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Alkaline earth metal
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Not to be confused with Alkali metal.
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
alkali metals ←    → group 3
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
Image: Lump of beryllium
Beryllium (Be)
4
3
Image: Magnesium crystals
Magnesium (Mg)
12
4
Image: Calcium stored under argon atmosphere
Calcium (Ca)
20
5
Image: Strontium floating in paraffin oil
Strontium (Sr)
38
6
Image: Barium stored under argon atmosphere
Barium (Ba)
56
7
Image: Radium electroplated on copper foil and covered with polyurethane to prevent reaction with air
Radium (Ra)
88

Legend

primordial element
element by radioactive decay

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]

Together with helium, these elements have in common an outer s orbital which is full[2][3][4]—that is, this orbital contains its full complement of two electrons, which the alkaline earth metals readily lose to form cations with charge +2, and an oxidation state of +2.[5] Helium is grouped with the noble gases and not with the alkaline earth metals, but it is theorized to have some similarities to beryllium when forced into bonding and has sometimes been suggested to belong to group 2.[6][7][8]

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.[9] There have been experiments, all unsuccessful, to try to synthesize element 120, the next potential member of the group.

Characteristics

[edit]

Chemical

[edit]

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 Electrons per 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, beryllium bromide and beryllium iodide, which are 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 unless at very high temperatures,[10] 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.[11] Even the compound beryllium fluoride, which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted.[12][13][14]

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

[edit]

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 (magnesium also reacts, but much more slowly), and also undergo transmetalation reactions to exchange ligands.

Solubility-related constants for alkaline-earth-metal fluorides
Metal M2+ hydration (-MJ/mol)[15] "MF2" unit hydration (-MJ/mol)[16] MF2 lattice (-MJ/mol)[17] Solubility (mol/kL)[18]
Be 2.455 3.371 3.526 soluble
Mg 1.922 2.838 2.978 1.2
Ca 1.577 2.493 2.651 0.2
Sr 1.415 2.331 2.513 0.8
Ba 1.361 2.277 2.373 6

Physical and atomic

[edit]
Key physical and atomic properties of the alkaline earth metals
Alkaline earth metal Standard atomic weight
(Da)[n 2][20][21] 		Melting point
(K) 		Melting point
(°C) 		Boiling point
(K)[4] 		Boiling point
(°C)[4] 		Density
(g/cm3)[22] 		Electronegativity
(Pauling) 		First ionization energy
(kJ·mol−1)		 Covalent radius
(pm)[23] 		Flame test color
Beryllium 9.012182(3) 1560 1287 2744 2471 1.845 1.57 899.5 105 White[24]
Magnesium 24.3050(6) 923 650 1363 1090 1.737 1.31 737.7 150 Brilliant-white[2]
Calcium 40.078(4) 1115 842 1757 1484 1.526 1.00 589.8 180 Brick-red[2]
Strontium 87.62(1) 1050 777 1655 1382 2.582 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 3] 969 696 2010 1737 5.502 0.9 509.3 221 Crimson red[n 4]

Nuclear stability

[edit]

Isotopes of all six alkaline earth metals are present in the Earth's crust and the Solar System at varying concentrations, dependent upon the nuclides' half-lives and, hence, their nuclear stabilities. The first five have one, three, five, four, and six stable (or observationally stable) isotopes respectively, for a total of 19 stable nuclides, as listed here: beryllium-9; magnesium-24, -25, -26; calcium-40, -42, -43, -44, -46; strontium-84, -86, -87, -88; barium-132, -134, -135, -136, -137, -138. The four underlined isotopes in the list are predicted by radionuclide decay energetics to be only observationally stable and to decay with extremely long half-lives through double-beta decay, though no decays attributed definitively to these isotopes have yet been observed as of 2024. Radium has no stable nor primordial isotopes.

In addition to the stable species, calcium and barium each have one extremely long-lived and primordial radionuclide: calcium-48 and barium-130, with half-lives of 5.6×1019 and 1.6×1021 years, respectively. Both are far longer than the current age of the universe (4.7× and 117× billion times longer, respectively) and less than one part per ten billion has decayed since the formation of the Earth. The two isotopes are stable for practical purposes.

Apart from the 21 stable or nearly-stable isotopes, the six alkaline earth elements each possess a large number of known radioisotopes. None of the isotopes other than the aforementioned 21 are primordial: all have half-lives too short for even a single atom to have survived since the Solar System's formation, after the seeding of heavy nuclei by nearby supernovae and collisions between neutron stars, and any present are derived from ongoing natural processes. Beryllium-7, beryllium-10, and calcium-41 are trace, as well as cosmogenic, nuclides, formed by the impact of cosmic rays with atmospheric or crustal atoms. The longest half-lives among them are 1.387 million years for beryllium-10, 99.4 thousand years for calcium-41, 1599 years for radium-226 (radium's longest-lived isotope), 28.90 years for strontium-90, 10.51 years for barium-133, and 5.75 years for radium-228. All others have half-lives of less than half a year, most significantly shorter.

Calcium-48 and barium-130, the two primordial and non-stable isotopes, decay only through double beta emission[n 5] and have extremely long half-lives, by virtue of the extremely low probability of both beta decays occurring at the same time. All isotopes of radium are highly radioactive and are primarily generated through the decay of heavier radionuclides. The longest-lived of them is radium-226, a member of the decay chain of uranium-238.[27] Strontium-90 and barium-140 are common fission products of uranium in nuclear reactors, accounting for 5.73% and 6.31% of uranium-235's fission products respectively when bombarded by thermal neutrons.[28] The two isotopes have half-lives each of 28.90 years and 12.7 days. Strontium-90 is produced in appreciable quantities in operating nuclear reactors running on uranium-235 or plutonium-239 fuel, and a minuscule secular equilibrium concentration is also present due to rare spontaneous fission decays in naturally occurring uranium.

Calcium-48 is the lightest nuclide known to undergo double beta decay.[29] Naturally occurring calcium and barium are very weakly radioactive: calcium contains about 0.1874% calcium-48,[30] and barium contains about 0.1062% barium-130.[31] On average, one double-beta decay of calcium-48 will occur per second for every 90 tons of natural calcium, or 230 tons of limestone (calcium carbonate).[32] Through the same decay mechanism, one decay of barium-130 will occur per second for every 16,000 tons of natural barium, or 27,000 tons of baryte (barium sulfate).[33]

The longest-lived isotope of radium is radium-226 with a half-life of 1600 years; it, along with radium-223, -224, and -228, occurs naturally in the decay chains of primordial thorium and uranium. Beryllium-8 is notable by its absence as it splits in half virtually instantaneously into two alpha particles whenever it is formed. The triple alpha process in stars can only occur at energies high enough for beryllium-8 to fuse with a third alpha particle before it can decay, forming carbon-12. This thermonuclear rate-limiting bottleneck is the reason most main sequence stars spend billions of years fusing hydrogen within their cores, and only rarely manage to fuse carbon before collapsing into a stellar remnant, and even then merely for a timescale of ~1000 years.[34] The radioisotopes of alkaline earth metals tend to be "bone seekers" as they behave chemically similar to calcium, an integral component of hydroxyapatite in compact bone, and gradually accumulate in the human skeleton. The incorporated radionuclides inflict significant damage to the bone marrow over time through the emission of ionizing radiation, primarily alpha particles. This property is made use of in a positive manner in the radiotherapy of certain bone cancers, since the radionuclides' chemical properties causes them to preferentially target cancerous growths in bone matter, leaving the rest of the body relatively unharmed.

Compared to their neighbors in the periodic table, alkaline earth metals tend to have a larger number of stable isotopes as they all possess an even number of protons, owing to their status as group 2 elements. Their isotopes are generally more stable due to nucleon pairing. This stability is further enhanced if the isotope also has an even number of neutrons, as both kinds of nucleons can then participate in pairing and contribute to nuclei stability.

History

[edit]

Etymology

[edit]

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia, and baria. These oxides are basic (alkaline) when combined with water. "Earth" was a 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,[35] thus supporting Lavoisier's hypothesis and causing the group to be named the alkaline earth metals.

Discovery

[edit]

The calcium compounds calcite and lime have been known and used since prehistoric times.[36] The same is true for the beryllium compounds beryl and emerald.[37] 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.[38][39][40]

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.[38][39][40]

Beryllium

[edit]
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.[37] Although it was originally thought that beryl was an aluminum silicate,[41] beryl was later found to contain a then-unknown element when, in 1797, Louis-Nicolas Vauquelin dissolved aluminum hydroxide from beryl in an alkali.[42] In 1828, Friedrich Wöhler[43] and Antoine Bussy[44] 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.[45] 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.[45]

Magnesium

[edit]

Magnesium was first produced by Humphry Davy in England in 1808 using electrolysis of a mixture of magnesia and mercuric oxide.[46] Antoine Bussy prepared it in coherent form in 1831. Davy's first suggestion for a name was magnium,[46] but the name magnesium is now used.

Calcium

[edit]

Lime has been used as a material for building since 7000 to 14,000 BCE,[36] and kilns used for lime have been dated to 2,500 BCE in Khafaja, Mesopotamia.[47][48] 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,[49] after hearing that Jöns Jakob Berzelius had prepared a calcium amalgam from the electrolysis of lime in mercury.

Strontium

[edit]

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,[50] who confirmed Crawford's discovery. Strontium was eventually isolated in 1808 by 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.[51]

Barium

[edit]
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 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.[52][53]

Radium

[edit]

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.[54] Radium was named in 1899 from the word radius, meaning ray, as radium emitted power in the form of rays.[55]

Occurrence

[edit]
Series of alkaline earth metals.

Beryllium occurs in the Earth's crust at a concentration of two to six parts per million (ppm),[56] 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.[57][58] However, in freshwater, beryllium is somewhat more common, with a concentration of 0.1 parts per billion.[59]

Magnesium and calcium are very common in the Earth's crust, being respectively the fifth and 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 15th most abundant element in the Earth's crust. The principal minerals are celestite and strontianite.[60] Barium is slightly less common, much of it in the mineral barite.[61]

Radium, being a decay product of uranium, is found in all uranium-bearing ores.[62] Due to its relatively short half-life,[63] radium from the Earth's early history has decayed, and present-day samples have all come from the much slower decay of uranium.[62]

Production

[edit]
Emerald, colored green with trace amounts of chromium, is a variety of the mineral beryl which is beryllium aluminum silicate.

Production of the six alkaline earth metals can occur in many ways.

Beryllium

[edit]
Main article: Beryllium § Extraction

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, aluminum 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.[11]

Strontium

[edit]
Main article: Strontium § Production

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.[64]

Barium

[edit]
Main article: Barium § Occurrence and production

To produce barium, barite (impure barium sulfate) is converted to barium sulfide by carbothermic reduction (such as with coke). The sulfide is water-soluble and easily reacted to form pure barium sulfate, used for commercial pigments, or other compounds, such as barium nitrate. These in turn are calcined into barium oxide, which eventually yields pure barium after reduction with aluminum.[61] The most important supplier of barium is China, which produces more than 50% of world supply.[65]

Magnesium

[edit]
Main article: Magnesium § Production

Magnesium is usually produced from magnesite ore, as well as dolomite. When dolomite is crushed, roasted and mixed with seawater in large tanks, magnesium hydroxide settles to the bottom. Heating, mixing in coke, and reacting with chlorine, then produces molten magnesium chloride. This can be electrolyzed, releasing magnesium, which floats to the surface.[66]

Calcium

[edit]
Main article: Calcium § Occurrence and production

Radium

[edit]
Main article: Radium § Production

Applications

[edit]

Beryllium is used mainly in military applications,[67] but non-military uses exist. In electronics, beryllium is used as a p-type dopant in some semiconductors,[68] and beryllium oxide is used as a high-strength electrical insulator and heat conductor.[69] Beryllium alloys are used for mechanical parts when stiffness, light weight, and dimensional stability are required over a wide temperature range.[70][71] Beryllium-9 is used in small-scale neutron sources that use the reaction 9Be + 4He (α) → 12C + 1n, the reaction used by James Chadwick when he discovered the neutron. Its low atomic weight and low neutron absorption cross-section would make beryllium suitable as a neutron moderator, but its high price and the readily available alternatives such as water, heavy water and nuclear graphite have limited this to niche applications. In the FLiBe eutectic used in molten salt reactors, beryllium's role as a moderator is more incidental than the desired property leading to its use.

Magnesium has many uses. It offers advantages over other structural materials such as aluminum, but magnesium's usage is hindered by its flammability.[72] Magnesium is often alloyed with aluminum, zinc and manganese to increase its strength and corrosion resistance.[73] Magnesium has many other industrial applications, such as its role in the production of iron and steel,[further explanation needed] and in the Kroll process for production of titanium.[74]

Calcium is used as a reducing agent in the separation of other metals such as uranium from ore. It is a major component of many alloys, especially aluminum and copper alloys, and is also used to deoxidize alloys. Calcium has roles in the making of cheese, mortars, and cement.[75]

Strontium and barium have fewer applications than the lighter alkaline earth metals. Strontium carbonate is used in the manufacturing of red fireworks.[76] Pure strontium is used in the study of neurotransmitter release in neurons.[77][78] Radioactive strontium-90 finds some use in RTGs,[79][80] which utilize its decay heat. Barium is used in vacuum tubes as a getter to remove gases.[61] Barium sulfate has many uses in the petroleum industry,[4][81] and other industries.[4][61][82]

Radium has many former applications based on its radioactivity, but its use is no longer common because of the adverse health effects and long half-life. Radium was frequently used in luminous paints,[83] although this use was stopped after it sickened workers.[84] The nuclear quackery that alleged health benefits of radium formerly led to its addition to drinking water, toothpaste, and many other products.[72] Radium is no longer used even when its radioactive properties are desired because its long half-life makes safe disposal challenging. For example, in brachytherapy, shorter-lived alternatives such as iridium-192 are usually used instead.[85][86]

Representative reactions of alkaline earth metals

[edit]

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 sulfur

Ca + 1/8S8 → CaS

Reaction with carbon

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

2Be + C → Be2C
CaO + 3C → CaC2 + CO (at 2500 °C 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 vapor.

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

[edit]

The flame test

The table below[87] presents the colors 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.[88]

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 sulfate 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

[edit]

Oxides

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

CaCO3 → CaO + CO2 (at approx. 900°C)

In laboratory, they are obtained from hydroxides:

Mg(OH)2 → MgO + H2O

or nitrates:

Ca(NO3)2 → CaO + 2NO2 + 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 2500°C)
CaC2 + N2 → CaCN2 + C
CaCN2 + H2SO4 → CaSO4 + H2N—CN
H2N—CN + H2O → (H2N)2CO (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 + 2 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

[edit]

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 (i.e. Behaves as a "bone seeker") where it bio-accumulates with a significant biological half life. While the bones themselves have higher radiation tolerance than other tissues, the rapidly dividing bone marrow does not and can thus be significantly harmed by Sr-90. The effect of ionizing radiation on bone marrow is also the reason why acute radiation syndrome can have anemia-like symptoms and why donation of red blood cells can increase survivability.

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.[11] Radium has a low availability and is highly radioactive, making it toxic to life.

Extensions

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The next alkaline earth metal after radium is thought to be element 120, although this may not be true due to relativistic effects.[89] 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.[90] 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.[91]

The chemistry of element 120 is predicted to be closer to that of calcium or strontium[92] instead of barium or radium. This noticeably contrasts with periodic trends, which 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.[92]

The next alkaline earth metal after element 120 has not been definitely predicted. Although a simple extrapolation using the Aufbau principle would suggest that element 170 is a congener of 120, relativistic effects may render such an extrapolation invalid. The next element with properties similar to the alkaline earth metals has been predicted to be element 166, though due to overlapping orbitals and lower energy gap below the 9s subshell, element 166 may instead be placed in group 12, below copernicium.[93][94]

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Explanatory notes

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Alkaline earth metal

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The alkaline earth metals are the six chemical elements in group 2 of the periodic table—beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra)—known for their similar chemical behaviors and role as s-block metals. These elements are all silvery-white, reactive metals that readily lose their two valence electrons to form stable +2 cations, exhibiting increasing reactivity down the group due to decreasing ionization energies. The term "alkaline earth" derives from the basic (alkaline) nature of their oxides and hydroxides when dissolved in water, combined with the historical classification of these insoluble oxides as "earths." Named for their position in the periodic table and shared electron configuration of [noble gas] ns², alkaline earth metals possess low densities, relatively low melting and boiling points compared to transition metals, and high melting points relative to alkali metals. Beryllium stands out as anomalous, showing more covalent character in its compounds due to its small size and high charge density, while radium is radioactive and rare. They react vigorously with water (except beryllium) to produce hydrogen gas and basic hydroxides, and their oxides are strongly basic, with basicity increasing down the group. Historically, compounds of these metals, such as lime (calcium oxide) and magnesia (), have been used since ancient times for and , but the pure metals of magnesium, calcium, , and were isolated in 1808 by , beryllium in 1828 by and Antoine Bussy, and radium in 1910 by . Beryllium was discovered in 1798 by Louis-Nicolas Vauquelin from beryl; magnesium, calcium, , and were isolated in 1808 by via ; and radium was discovered in 1898 by Marie and from pitchblende. In nature, these metals occur primarily as minerals like dolomite and rather than in elemental form due to their reactivity, with calcium and magnesium being among the most abundant elements in the . Alkaline earth metals have diverse applications leveraging their properties: magnesium in lightweight alloys for and automotive industries, calcium in production and as a dietary supplement for health, strontium in for red flames and in , barium in contrast agents and drilling fluids, beryllium in nuclear reactors and alloys for its stiffness, and radium historically in luminescent paints though now limited due to . Their compounds also play key roles in , (e.g., lime for adjustment), and .

Properties

Physical properties

The alkaline earth metals are silvery-white metals with high electrical and thermal conductivity and a lustrous appearance when freshly cut, though they rapidly in air due to formation. These properties arise from their , involving delocalized valence electrons, which becomes weaker down the group as atomic size increases. Densities of the alkaline earth metals vary, with at 1.85 g/cm³, magnesium at 1.74 g/cm³, calcium at 1.55 g/cm³, at 2.64 g/cm³, at 3.51 g/cm³, and at 5.5 g/cm³. While there is an initial decrease from to calcium due to a more rapid increase in atomic relative to , density generally increases down the group as atomic rises faster than , reflecting larger atomic radii.
ElementDensity (g/cm³)Melting Point (°C)Boiling Point (°C)
Be1.8512872470
Mg1.746501090
Ca1.558421484
Sr2.647771382
Ba3.517271897
Ra5.5~700~1737
The and boiling points are anomalously high for due to its small atomic size and partial covalent bonding. For the heavier members, melting points show a general decrease down the group (Mg: 650°C, Ca: 842°C, Sr: 777°C, Ba: 727°C, Ra: ~700°C), despite the increase from Mg to Ca, while boiling points initially increase from Mg to Ca then decrease slightly before rising again for Ba (Mg: 1090°C, Ca: 1484°C, Sr: 1382°C, Ba: 1897°C, Ra: ~1737°C). This irregularity arises from the increasing atomic size weakening metallic bonds, with variations due to changes in coordination and . Beryllium's anomalously high values stem from its small size and strong bonding. Atomic radii increase progressively from (112 pm) to (estimated 250 pm) due to the addition of shells, while ionic radii for the M²⁺ ions also increase, from ²⁺ (27 pm) to ²⁺ (135 pm), reflecting weaker per . In terms of , the crystal vary: and magnesium adopt a hexagonal close-packed (hcp) lattice, calcium and a face-centered cubic (fcc) lattice, and and a body-centered cubic (bcc) . These differences influence properties like and . Regarding mechanical properties, and magnesium are relatively hard and brittle owing to directional covalent character in bonding, whereas calcium, , and are softer and more malleable, allowing easier deformation.

Chemical properties

The alkaline earth metals, consisting of beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra), all exhibit a general electron configuration of [noble gas] ns² in their valence shell, where n is the principal quantum number for the outermost shell. This ns² configuration results in these elements predominantly forming ions with a +2 oxidation state, as they lose both valence electrons to achieve a stable noble gas configuration. Unlike some other metals, higher oxidation states are rare and unstable for these elements due to the high energy required to remove additional electrons. The chemical reactivity of alkaline earth metals increases down the group, primarily driven by a decreasing trend in ionization energies, which facilitates easier loss of the valence electrons. The first ionization energy decreases from 899 kJ/mol for Be to approximately 509 kJ/mol for Ra, reflecting the increasing atomic size and shielding effect that reduces the effective nuclear charge on the valence electrons. For instance, Mg has a first ionization energy of 738 kJ/mol and a second of 1451 kJ/mol, while the sum of the first two ionization energies becomes progressively lower down the group, stabilizing the M²⁺ ions and enhancing reactivity; Be is the least reactive, often forming covalent compounds due to its small size and high charge density, whereas Ra is the most reactive, approaching the behavior of alkali metals. This trend underscores the metals' tendency to form ionic compounds, with reactivity manifesting in vigorous reactions with water and oxygen, though specifics vary by element. Hydration energies of the M²⁺ ions decrease down the group from -2494 kJ/mol for Be²⁺ to -1305 kJ/mol for Ba²⁺, as the larger ionic radii reduce the electrostatic attraction between the ion and water molecules, influencing the of their salts in aqueous solutions. This decrease in contributes to anomalies in patterns, such as the increasing of sulfates from MgSO₄ to BaSO₄, where and hydration effects balance differently. Flame tests provide a distinctive method for identifying alkaline earth metals based on their characteristic emission colors, arising from electronic transitions in excited atoms. Calcium produces a brick-red flame corresponding to emissions around nm, strontium yields a crimson red at approximately 641 nm, and barium emits an apple-green color from lines near 524 nm, allowing qualitative detection in . Beryllium exhibits a notable with aluminum in group 13, attributed to their similar charge-to-radius ratios and atomic sizes, leading to comparable chemical behaviors such as the formation of amphoteric oxides and covalent halides like BeCl₂, which is structurally analogous to AlCl₃ in being dimeric and Lewis acidic. This relationship highlights deviations from typical group trends for the lighter alkaline earth metals.

Nuclear properties

The alkaline earth metals exhibit a range of nuclear properties characterized by their isotopic compositions, stability patterns, and interactions in nuclear processes. possesses only one , ^{9}Be, which constitutes 100% of naturally occurring beryllium. has three : ^{24}Mg (78.99%), ^{25}Mg (10.00%), and ^{26}Mg (11.01%). Calcium's include ^{40}Ca (96.948%), ^{42}Ca (0.647%), ^{43}Ca (0.135%), ^{44}Ca (2.086%), ^{46}Ca (0.004%), and ^{48}Ca (0.187%), with ^{40}Ca dominating. features four : ^{84}Sr (0.56%), ^{86}Sr (9.86%), ^{87}Sr (7.00%), and ^{88}Sr (82.58%). has seven : ^{130}Ba (0.11%), ^{132}Ba (0.10%), ^{134}Ba (2.42%), ^{135}Ba (6.59%), ^{136}Ba (7.85%), ^{137}Ba (11.23%), and ^{138}Ba (71.70%), led by ^{138}Ba. , in contrast, has no , with all known isotopes being radioactive.
ElementStable IsotopesNatural Abundance (%)
Beryllium^{9}Be100
Magnesium^{24}Mg, ^{25}Mg, ^{26}Mg78.99, 10.00, 11.01
Calcium^{40}Ca (dominant)96.948
^{88}Sr (dominant)82.58
^{138}Ba (dominant)71.70
NoneN/A
Radioactive isotopes of alkaline earth metals play roles in natural decay chains and cosmogenic processes. The most prominent is ^{226}Ra, with a of 1600 years, formed as part of the decay series and decaying via alpha emission to ^{222}Rn. , a cosmogenic isotope produced by , has a of 53.22 days and decays by to lithium-7. , another trace cosmogenic , persists with a of approximately 99,400 years, generated primarily through on calcium in the upper atmosphere. Nuclear stability among alkaline earth metals follows general trends observed in nuclides, influenced by the even-odd rule and magic numbers. These elements have even atomic numbers (Z = 4, 12, 20, 38, 56, 88), favoring isotopes with even neutron numbers (N) due to enhanced pairing stability, which contributes to the prevalence of even-N stable isotopes like ^{9}Be (N=5, odd but light nucleus exception), ^{24}Mg (N=12), ^{40}Ca (N=20), ^{88}Sr (N=50), and ^{138}Ba (N=82). A notable example is ^{40}Ca, a doubly magic nucleus with Z=20 and N=20—both magic numbers—resulting in exceptional stability from closed proton and neutron shells. This shell structure enhances binding and resistance to decay, explaining ^{40}Ca's dominance despite a neutron-to-proton ratio slightly below the typical stable band for mid-mass nuclei. Certain isotopes participate in nuclear reactions relevant to s and historical applications. Barium isotopes, particularly even-mass ones like ^{130}Ba, ^{132}Ba, ^{134}Ba, ^{136}Ba, and ^{138}Ba, undergo to produce radioisotopes used in and research, with cross-sections measured for reactor production of ^{131}Ba and ^{133}Ba. Radium-226 was employed in early 20th-century for treating skin conditions and cancers via , leveraging its alpha and gamma emissions before safer alternatives emerged. Binding energies per nucleon reflect increasing nuclear stability across the group, peaking near calcium and strontium before a gradual decline. For representative stable isotopes, the values are approximately 6.46 MeV for ^{9}Be, 8.26 MeV for ^{24}Mg, 8.55 MeV for ^{40}Ca, 8.73 MeV for ^{88}Sr, and 8.39 MeV for ^{138}Ba, illustrating the semi-empirical mass formula's trends where shell effects and pairing boost stability around mid-mass regions. These energies underscore why lighter alkaline earth nuclei like beryllium are less tightly bound, while calcium's magic configuration maximizes cohesion.

History

Etymology

The term "alkaline earth metals" refers to the group of elements in group 2 of the periodic table, whose oxides were historically termed "alkaline earths" in early chemistry because they are insoluble in but dissolve in acids to produce alkaline (basic) solutions, distinguishing them from the more soluble "alkalis" formed by metals. This nomenclature evolved in the 18th and 19th centuries as chemists like and isolated and characterized these substances, with further systematizing the classification in his 1828 publication on atomic weights, where he grouped the elements based on their chemical similarities. Initially, was excluded from the group due to the amphoteric (both acidic and basic) nature of its oxide, unlike the strongly basic oxides of the others; it was later included based on periodic table trends. The individual elements have names rooted in their discovery contexts, minerals, or properties: Beryllium derives its name from the mineral beryl, in which it is found, from the Greek word beryllos meaning a . Magnesium is named after Magnesia, an ancient district in , , where magnesium-rich minerals like were abundant. Calcium comes from the Latin calx, meaning lime, referring to (quicklime), one of the earliest known compounds of the element used in and . Strontium is named after , a village in near a lead mine where the strontianite () was first identified in 1790. Barium originates from the Greek barys, meaning heavy, alluding to the high density of its compounds like (). Radium was named by Marie and in 1898 from the Latin , meaning ray, due to the intense radiation emitted by the element.

Discovery and isolation of beryllium

In 1798, French chemist Louis-Nicolas Vauquelin identified a new earth, later known as beryllia (, BeO), during his chemical analysis of the beryl and emeralds. Beryl, with the composition Be₃Al₂Si₆O₁₈, had long been suspected to contain an unknown component distinct from alumina (Al₂O₃), and Vauquelin's work confirmed this by isolating the oxide from these gemstones. He named it "glucina" due to the sweet taste of its salts, a property that initially masked its distinct identity. Vauquelin's key experiments involved dissolving beryl in acids and separating the resulting earth through and , revealing properties that set it apart from alumina./Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group_2_Elements:The_Alkaline_Earth_Metals/Z004_Chemistry_of_Beryllium(Z4)) Unlike alumina, which exhibits amphoteric behavior by dissolving in both acids and bases, beryllia dissolved only in acids and resisted fusion with alkalis, indicating its unique chemical nature. This distinction was crucial amid late-18th-century advances in , where chemists like Vauquelin were systematically analyzing silicates to uncover new elements. The elemental metal was first isolated in 1828 through independent efforts by German chemist and French chemist Antoine-Alexandre-Brutus Bussy. Both used the reduction of (BeCl₂) with metal, heating the mixture to yield small quantities of impure beryllium, described as a gray, brittle substance. Wöhler's method mirrored his earlier success in isolating aluminum, applying the same potassium reduction technique to prepared from beryllia. Beryllium's recognition as an alkaline earth metal was delayed by its anomalous covalent character, which caused its compounds to behave more like those of aluminum than typical group 2 elements, complicating early classifications./Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group_2_Elements:The_Alkaline_Earth_Metals/Z004_Chemistry_of_Beryllium(Z4)) This period from 1798 to 1828 coincided with broader progress in , including the identification of other rare earths, though beryllium's scarcity and handling difficulties limited further immediate study.

Discovery and isolation of magnesium

The mineral known as , or heptahydrate (MgSO₄·7H₂O), was first identified in 1618 near , , where it was noted for its bitter taste and medicinal properties, including use as a and . Local cowherd Henry Wicker discovered the substance when his cattle refused to drink from a spring. In 1755, Scottish chemist conducted experiments that distinguished magnesia alba (, MgO) from lime (, CaO), recognizing MgO as a distinct earth with unique properties, such as lighter weight and resistance to certain acids, thereby establishing magnesium as a separate element. Black's work, detailed in his lectures at the , laid the foundation for by demonstrating that magnesia alba lost weight differently upon heating compared to lime and did not effervesce with acids in the same manner. The isolation of metallic magnesium was achieved in 1808 by English chemist , who electrolyzed (MgO) using a mercury in a large battery setup, producing a magnesium-mercury amalgam that was subsequently heated to expel the mercury and yield impure magnesium metal. Davy's method, performed at the Royal Institution in , marked the first production of the element in metallic form, though the yield was small and contaminated. Purer magnesium metal was obtained in 1831 by French chemist Antoine-Alexandre-Brutus Bussy, who reduced anhydrous (MgCl₂) with metal in a sealed , yielding several grams of the element in a more refined state suitable for further study. Bussy's technique, published in the Annales de Chimie et de Physique, improved upon Davy's amalgam process by avoiding mercury impurities and producing larger quantities, highlighting magnesium's potential as a lightweight, reactive metal.

Discovery and isolation of calcium

The use of calcium compounds dates back to prehistoric times, with lime (calcium oxide, CaO) being employed in construction as early as 7000 BCE in the . Archaeological evidence indicates that quicklime was produced by heating (, CaCO3) in rudimentary kilns, resulting in a material used for mortar and in early buildings and structures. This ancient process of , involving temperatures around 900–1000°C to decompose limestone into CaO and , laid the foundation for widespread applications in binding materials for architecture across civilizations. In the late , advances in chemical understanding elevated lime from a practical substance to a subject of elemental inquiry. In 1789, proposed that lime was the of an unknown metal, classifying it among the "salifiable earths" in his revolutionary and thereby recognizing its elemental nature rather than treating it as a simple compound. This insight bridged empirical observations with emerging atomic theory, setting the stage for isolation efforts. The element calcium was first isolated in metallic form in 1808 by British chemist at the Royal Institution in . Davy achieved this through of a molten mixture of lime (CaO) and mercuric (HgO), using a battery-powered apparatus to decompose the compound and deposit calcium amalgam, from which the pure metal was subsequently obtained. He named the new element "calcium," derived from the Latin "" meaning lime, honoring its historical compound form. Davy's work marked a pivotal shift from oxide-based uses to the recognition of calcium as a distinct alkaline earth metal. Further refinements in isolation techniques came in the late , with French achieving a significantly purer form of calcium in 1898. Building on electrolytic methods, Moissan electrolyzed fused (CaI2) in his electric furnace setup, yielding calcium metal with approximately 99% purity and enabling more accurate studies of its . This advancement, part of Moisson's broader contributions to high-temperature , overcame impurities in earlier samples and solidified calcium's place in metallurgical research.

Discovery and isolation of strontium

In 1790, Scottish chemist Adair Crawford and his colleague William Cruickshank identified a new mineral, (SrCO₃), in lead mines near the village of in , initially mistaking it for a barium compound similar to . In 1791, , a at the , conducted a detailed analysis of and distinguished it from lime () based on its greater solubility in water and its unique crimson-red flame coloration when heated, confirming as a distinct element. named the element after the Scottish locality where the mineral was found. The sulfate mineral celestite (SrSO₄), another key source of strontium, was formally identified and described in 1792 as part of the broader recognition of strontium compounds. This mineral, noted for its sky-blue crystals, further supported the element's distinction from calcium and barium through chemical tests. The isolation of metallic strontium occurred in 1808 when Sir Humphry Davy, using electrolysis on a mixture of strontium chloride (SrCl₂) and mercuric chloride (HgCl₂), obtained the pure metal for the first time at the Royal Institution in . Davy's method, part of his pioneering work on alkaline earth metals, involved passing an through molten salts to decompose them. Early recognition of strontium's red flame led to its use in fireworks and by the early , where strontium salts produced vibrant crimson displays.

Discovery and isolation of barium

The compound , known as baryta (BaO), was first isolated in 1774 by Swedish chemist from the heavy spar, or barite (BaSO₄), by dissolving it in and observing the low of the resulting precipitate, which distinguished it from calcium compounds. Scheele named this new "earth" terra ponderosa, or heavy earth, due to its high density compared to other alkaline earths. The nomenclature for baryta originated from the Greek word "barys," meaning heavy. In the late 18th century, French chemist Louis-Bernard Guyton de Morveau initially termed the oxidized form barote in his 1782 proposals for chemical nomenclature, emphasizing its weighty properties; this was later refined to baryta by to align with systematic naming conventions for earths./04:Group_2-_The_Alkaline_Earth_Metals/4.01:_The_Alkaline_Earth_Elements) Barite itself had earlier practical applications, notably as a white pigment in the 17th century under names like "Bologna stone," serving as a non-toxic substitute for in artists' paints due to its opacity and stability. The elemental metal was first isolated in 1808 by English chemist Sir Humphry Davy through of a mixture of molten baryta and mercuric oxide, building on earlier electrolytic attempts by and Magnus Pontin. Davy named the silvery-white metal , derived from baryta, and presented his findings to the Royal Society, completing the identification of the stable alkaline earth metals. To distinguish barium from the chemically similar strontium, early chemists relied on precipitation tests; for instance, barium ions form a bright yellow precipitate of (BaCrO₄) with chromate solutions, whereas strontium chromate (SrCrO₄) appears white, allowing reliable separation in mineral analyses. This method, leveraging the insolubility of and the color difference in chromates, was crucial for confirming barium's presence in heavy spar deposits amid overlapping alkaline earth properties.

Discovery and isolation of radium

The discovery of radium occurred in December 1898, when , , and their collaborator Gustave Bémont announced the isolation of a highly radioactive element from pitchblende residues, a processed at the Joachimsthal mines in . The Curies had observed that pitchblende exhibited far greater than pure , prompting them to fractionate the chemically and identify a barium-rich component with exceptional activity—over 300 times that of —which they named radium from the Latin word for ray. This breakthrough built on Henri Becquerel's 1896 observation of rays and was detailed in their publication in the Comptes rendus hebdomadaires des séances de l'Académie des sciences. The isolation process was arduous, requiring the processing of several tons of pitchblende residues to yield trace amounts of radium compounds, as the element occurs in minute concentrations—about 1 part per 3 million in the ore. employed fractional crystallization of and chlorides (or bromides) to separate the new element, exploiting subtle differences in despite their close , which initially led to confusion with salts. By April 1902, after thousands of recrystallizations, she obtained 0.1 grams of pure , determining its atomic weight as approximately 226—twice that of —thus identifying as the primary in the . Further refinement culminated in 1910, when , with André-Louis Debierne, produced metallic through of pure in a mercury , followed by to obtain about 0.1 grams of the shiny white metal, which rapidly tarnished in air. This achievement definitively proved 's existence as an element distinct from . For her contributions to the discovery and isolation of , received the in 1911, recognizing the profound impact on understanding .

Occurrence and abundance

Cosmic and terrestrial abundance

The alkaline earth metals display distinct abundance patterns in the , reflecting their nucleosynthetic origins and stability in stellar environments. Calcium is the most abundant among them in solar system materials, with a of approximately 6.5×10−56.5 \times 10^{-5} (or 0.0065% by mass), primarily produced through alpha-particle capture during and carbon burning in massive stars. Magnesium ranks sixth overall in cosmic abundance, at about 7.2×10−47.2 \times 10^{-4} (0.072% by mass), also resulting from efficient alpha-capture processes in stellar fusion. and are less common, with around 5×10−85 \times 10^{-8} and 1.7×10−81.7 \times 10^{-8} respectively, formed mainly via slow and rapid neutron-capture reactions in asymptotic giant branch stars and explosive events. stands out as the rarest, at roughly 2.3×10−102.3 \times 10^{-10} (or 2.3×10−8%2.3 \times 10^{-8} \%), since it is readily destroyed by fusion in stars and mainly arises from of heavier nuclei in the . has virtually no primordial cosmic abundance, estimated below 1 part per trillion by , as it forms exclusively through the of and . These values are based on solar photospheric abundances from 3D non-LTE models. Terrestrial abundances of these elements have been modified by planetary accretion, core formation, and crustal differentiation from a chondritic starting composition. In the continental crust, calcium ranks fifth in overall abundance at 3.6% by weight, concentrated in feldspars and minerals. Magnesium is eighth at 2.1% by weight, largely bound in ferromagnesian silicates like and . occurs at 0.037% by weight and at 0.0425%, both behaving geochemically like calcium and substituting into its lattice sites in minerals. is trace-level at 2.8 parts per million by weight, while remains exceedingly scarce, below 1 part per trillion by weight due to its short and dependence on parent radionuclides. In the mantle and core, fractionation further accentuates differences. Magnesium dominates the silicate portion, comprising about 22% by weight in upper mantle rocks such as peridotite, where it forms the backbone of olivine and orthopyroxene. Calcium is present at roughly 2.2% by weight, incorporated into clinopyroxene, garnet, and minor plagioclase. The core is depleted in these lithophile elements, with negligible amounts beyond trace levels in iron-nickel alloys. These distributions stem from siderophile-lithophile partitioning during Earth's differentiation, preserving magnesium and calcium in the silicate Earth while enriching the crust via partial melting.

Principal minerals and deposits

Beryllium primarily occurs in the minerals beryl (Be3Al2Si6O18\mathrm{Be_3Al_2Si_6O_{18}}) and bertrandite, with the latter being the main source for industrial production. The principal deposit in the United States is the Spor Mountain site in , which supplies bertrandite ore, while beryl is largely imported from deposits in and . Magnesium is found in (MgCO3\mathrm{MgCO_3}), dolomite (CaMg(CO3)2\mathrm{CaMg(CO_3)_2}), and extracted from and brines. Major global deposits of are concentrated in , which dominates production, and , where significant reserves support mining operations. Seawater brines provide an additional vast resource, particularly along coastal regions. Calcium is abundant in (CaCO3\mathrm{CaCO_3}), (CaSO4⋅2H2O\mathrm{CaSO_4 \cdot 2H_2O}), and (CaF2\mathrm{CaF_2}), making it one of the most widespread elements in the . These minerals form extensive deposits worldwide, with limestone particularly prominent in landscapes such as those in the Yunnan region of and the in , . Gypsum beds are common in sequences, like those in the . Strontium occurs mainly in celestite (SrSO4\mathrm{SrSO_4}) and (SrCO3\mathrm{SrCO_3}), with celestite being the dominant commercial source. Key deposits are located in , particularly in , and in , where large-scale supports global supply; other notable sites include those in and . Barium is chiefly derived from barite (BaSO4\mathrm{BaSO_4}), which forms in sedimentary, hydrothermal, and vein deposits. The largest reserves are in , followed by and the , where Mississippi Valley-type deposits in states like and yield significant quantities. Radium is present only in trace amounts within uranium-bearing minerals such as and carnotite, with no dedicated commercial deposits. It is recovered as a from the tailings of processing, historically from sites like those in the in the United States.

Production

Beryllium production

Beryllium is primarily extracted from two minerals: bertrandite and , which are processed through a multi-step industrial method to produce high-purity metal. The process begins with the roasting of crushed ore in the presence of or to form (), followed by conversion to beryllium fluoride (BeF₂) via reaction with . This fluoride is then reduced at approximately 1200°C using magnesium metal in a , yielding beryllium metal and magnesium fluoride as a according to the reaction:
\ceBeF2+Mg−>Be+MgF2\ce{BeF2 + Mg -> Be + MgF2}
The resulting beryllium ingots are further purified through or to achieve 99.9% purity or higher. Global annual production of beryllium metal is estimated at around 300 metric tons, with the United States accounting for the majority (over 70%) through operations at facilities like those operated by Materion Corporation in Utah, which process domestically mined bertrandite from the Spor Mountain deposit. The production process faces significant challenges due to beryllium's toxicity, necessitating inert atmospheres, specialized ventilation, and protective equipment to prevent inhalation of fine particles, which can cause chronic beryllium disease. Recycling from scrap alloys, such as those from aerospace components, contributes up to 20% of supply and involves similar purification steps to recover usable metal. Byproducts like magnesium fluoride and other fluorides are typically managed as hazardous waste, requiring neutralization and disposal in accordance with environmental regulations to minimize fluoride emissions.

Magnesium production

Magnesium is extracted on a large scale from magnesium-rich ores, such as dolomite (CaMg(CO₃)₂), and from natural brines and , which serve as primary sources for industrial production. Seawater contains approximately 0.13% magnesium by weight, making it a vast but dilute resource, while hypersaline brines like those in the Dead Sea have much higher concentrations, around 46 g/L of magnesium. Global primary magnesium production reached about 1,057,000 metric tons in 2024, with accounting for over 80% of output, primarily from domestic dolomite deposits. The dominant method worldwide is the Pidgeon process, a thermal reduction technique that accounts for the majority of production, especially in . In this process, calcined dolomite is mixed with (containing about 75% ) and formed into briquettes, which are then heated to 1100–1200°C under in horizontal . The reaction proceeds as follows: 2MgO+2CaO+Si→2Mg (g)+Ca2SiO42 \text{MgO} + 2 \text{CaO} + \text{Si} \rightarrow 2 \text{Mg (g)} + \text{Ca}_2\text{SiO}_4 Magnesium vapor is produced and condensed into liquid metal at the cooler end of the retort, yielding crude magnesium at 90–95% purity. This method is favored for its relatively low capital investment and flexibility with heat sources, though it requires high temperatures and generates significant waste slag. An alternative is the electrolytic Dow process, used primarily in regions with access to low-cost and resources, such as former operations in the United States and . Anhydrous (MgCl₂) is electrolyzed in cells at around 700°C, with a typical cell voltage of 4.5–5.0 V. Magnesium is deposited as at the , while gas is liberated at the for . The process operates with an mixture including NaCl, KCl, and CaCl₂ to lower the and improve conductivity. This method produces higher-purity magnesium directly but is more energy-intensive electrically. Electrolytic production consumes 12–14 kWh per kg of magnesium, reflecting the high energy demand for maintaining molten conditions and driving the decomposition (theoretical voltage ~3 V, but practical higher due to overpotentials). In contrast, thermal processes like Pidgeon rely on coal or other fuels for heat, with total energy input equivalent to about 35–40 GJ per ton, but they use minimal electricity and can be considered greener in terms of electrical grid burden when paired with renewable thermal sources, though they emit more CO₂ overall. Efforts to hybridize these methods, such as using solar thermal input for electrolysis, aim to reduce environmental impacts. Crude magnesium from either process undergoes purification to remove impurities like iron, silicon, and manganese, which affect its properties. Vacuum distillation is a common industrial technique, heating the metal to 600–700°C under low pressure (10–100 Pa) to selectively evaporate and recondense magnesium, achieving purities up to 99.99%. For ultra-high purity applications, zone refining involves passing a narrow molten zone along a magnesium ingot using induction heating, segregating impurities to the ends, which are then cropped off; this can yield 99.999% purity or higher. These methods ensure the metal meets standards for alloys and other uses.

Calcium production

Calcium metal is produced on a relatively small industrial scale compared to other alkaline earth metals, primarily through the aluminothermic reduction of (lime, CaO) with aluminum under high vacuum and temperature conditions. The process involves mixing calcined lime with aluminum powder and heating to approximately 1200°C in a , where the exothermic reaction produces calcium vapor and aluminum oxide slag: 3\ceCaO+2\ceAl→3\ceCa+\ceAl2O33 \ce{CaO} + 2 \ce{Al} \rightarrow 3 \ce{Ca} + \ce{Al2O3} The calcium vapor is then condensed and collected as crude metal, which is further purified by vacuum sublimation or at reduced pressures (around 0.1–1 mbar) and temperatures of 800–1000°C to achieve purities exceeding 99%. This method dominates production due to its efficiency in separating the volatile calcium from the . An alternative electrolytic method involves the decomposition of molten (CaCl₂) in a Downs cell-like setup at 800–900°C, with calcium deposited at the and gas at the . This process, historically used in the early , is less common today due to higher energy costs (about 10–12 kWh/kg) but offers direct high-purity output in regions with cheap . Preparation of anhydrous CaCl₂ typically starts from via to CaCO₃, calcination to CaO, and chlorination. Global production of calcium metal is estimated at around 10,000 metric tons annually as of 2023, with accounting for the majority (over 80%), followed by and smaller outputs in and the . The metal's production is niche, driven by demand in alloys and , and faces challenges from high energy requirements and reactivity, necessitating inert handling to prevent oxidation. Byproducts like aluminum are recovered for reuse in refractories.

Strontium production

Strontium is primarily extracted from its main ore, celestite (SrSO₄), through a multi-step chemical reduction to obtain the metal. The initial step involves roasting celestite ore at high temperatures to form strontium (SrSO₃), which is then reduced using coke in a controlled furnace environment. This reduction converts the to strontium (SrS) via the reaction: SrSO3+4C→SrS+4CO\text{SrSO}_3 + 4\text{C} \rightarrow \text{SrS} + 4\text{CO} The strontium is subsequently subjected to aluminothermic reduction with aluminum metal under conditions, yielding crude strontium metal and aluminum according to: SrS+2Al→Sr+Al2S3\text{SrS} + 2\text{Al} \rightarrow \text{Sr} + \text{Al}_2\text{S}_3 This exothermic reaction occurs at temperatures around 1,000–1,200°C, with the low pressure aiding in the separation of the volatile strontium vapor from the slag. An alternative production method involves the electrolytic decomposition of molten strontium chloride (SrCl₂), typically prepared from celestite via conversion to the carbonate and subsequent chlorination. In this process, SrCl₂ is electrolyzed in a Downs cell-like setup at approximately 800–900°C, producing strontium metal at the cathode and chlorine gas at the anode. While less common than the aluminothermic route due to higher energy requirements, electrolysis offers potential for higher purity output in specialized applications. Global production of strontium compounds, primarily carbonates and sulfates derived from celestite processing, reaches approximately 200,000 tons annually, with major contributors being (around 100,000 tons of celestite mined) and (about 80,000 tons). In contrast, strontium metal production remains limited to roughly 10 tons per year, concentrated in facilities in and , reflecting its niche industrial demand. The crude metal obtained from either method is purified via at reduced pressures (1–5 mbar) and temperatures of 700–800°C, removing impurities like aluminum and to achieve purities exceeding 97%.

Barium production

Barium production begins with the and processing of barite (, BaSO₄), the most abundant and commercially important of . Global barite production reached approximately 8.16 million metric tons in 2022, with major producers including , , and ; the majority is used in oil and gas drilling fluids, while a smaller fraction serves as feedstock for barium chemicals and metal. The primary industrial process for extracting barium involves the carbothermic reduction of crushed and ground barite ore mixed with coke or in a at around 1000°C. This yields (BaS) via the black ash process, represented by the reaction: \ceBaSO4+4C−>BaS+4CO\ce{BaSO4 + 4C -> BaS + 4CO} The resulting black ash—a of BaS, carbon residues, and impurities—is quenched and leached with hot water to dissolve the soluble BaS, producing a lye after to remove insoluble materials such as silica and iron compounds. Further purification of the lye occurs through to remove sulfides or precipitation of impurities, with high-purity intermediates obtained via steps, such as to form (BaO). To produce barium metal, the purified BaS is first converted to (BaCl₂), typically by reacting with or . The molten BaCl₂ is then electrolyzed at high temperatures (around 800–900°C) in a Downs cell-like setup, where deposits at the and gas is liberated at the : \ceBaCl2−>Ba+Cl2\ce{BaCl2 -> Ba + Cl2} An alternative thermal method involves converting BaS to (via carbonation to BaCO₃ followed by ) and then reducing BaO with aluminum in a at about 1100°C to form an Ba-Al compound, which is subsequently distilled to separate pure metal. Barium metal production remains minor compared to compounds, reflecting its niche applications in and alloys. A critical precursor in barium processing is (BaCO₃), produced by reacting the BaS with under controlled conditions to form a precipitate: \ceBaS+H2O+CO2−>BaCO3+H2S\ce{BaS + H2O + CO2 -> BaCO3 + H2S} The BaCO₃ is filtered, washed, and calcined for use in further syntheses, including those leading to metal production; global output of exceeds 700,000 tons annually, underscoring its role as an intermediate.

Radium production

Radium was first isolated in 1898 by Marie and Pierre Curie from pitchblende ore through a chemical that exploited its similarity to . The method entailed dissolving the ore in hydrochloric or to solubilize the components, precipitating radium alongside as insoluble sulfates or carbonates after removing , converting the precipitate to radium-, and achieving separation via repeated fractional crystallization, leveraging the marginally higher solubility of . This labor-intensive technique yielded approximately 0.14 grams of radium per ton of pitchblende ore. In their initial efforts, the Curies processed several tons of ore to obtain less than 1 milligram of . Industrial-scale production, primarily in the United States and , adopted and refined the Curie process in the early , with output peaking in the at about 18.5 grams annually from major facilities like the Standard Chemical Company in . Between 1913 and 1920, U.S. refineries alone produced around 70 grams, much of it destined for luminous paints and medical applications.30555-7/fulltext) Worldwide cumulative production until the mid- totaled slightly more than 3,000 grams, after which commercial extraction halted in the due to reduced demand and the rise of synthetic radioisotopes. In modern times, radium-226 is obtained in trace amounts as a from , primarily from mill of uranium processing operations, rather than dedicated . Extraction involves leaching the with or salt solutions, followed by purification using cation exchange resins to selectively bind and elute radium, or solvent extraction with organic phases to isolate it from interfering ions like and calcium. These methods yield highly pure radium for research but are not conducted commercially, with total contemporary output limited to milligrams annually for specialized applications such as isotope production. Owing to its emission of high-energy alpha particles, which pose severe internal hazards despite low penetration, is manipulated exclusively within shielded hot cells equipped with remote handling tools and ventilation systems to contain radioactive aerosols and prevent personnel exposure.

Compounds

Oxides, peroxides, and hydroxides

The oxides of the alkaline earth metals generally adopt the formula MO, where M is the metal, and exhibit increasing basicity down the group due to the decreasing of the metal cations. () is amphoteric with a high of approximately 2575 °C, making it suitable as a refractory material, though it shows limited in and reacts slowly with acids or bases. (MgO) is a basic, white solid used in high-temperature refractories owing to its stability up to 2800 °C; it reacts with to form but does so less vigorously than heavier analogs. (CaO), known as quicklime, is strongly basic and reacts exothermically with to produce (Ca(OH)₂), a process termed slaking. (SrO) and (BaO) are even more strongly basic, readily absorbing moisture from air to form the corresponding hydroxides and displaying high reactivity with acids./08:_Chemistry_of_the_Main_Group_Elements/8.05:_Group_2_The_Alkaline_Earth_Metals/8.5.02:_Alkaline_Earth_Metals'_Chemical_Properties) Peroxides of the alkaline earth metals have the formula MO₂ and are more stable for the heavier members (Ca, Sr, ) than for beryllium or magnesium, where normal s predominate. These peroxides decompose upon heating to yield the corresponding and oxygen gas via the reaction MO₂ → MO + ½O₂. (BaO₂), a pale yellow solid, is notably stable and serves as a bleaching agent in industrial applications by releasing oxygen to oxidize organic stains. Calcium and peroxides exhibit similar decomposition behavior but are less commonly utilized due to lower stability compared to ./08:_Chemistry_of_the_Main_Group_Elements/8.05:_Group_2_The_Alkaline_Earth_Metals/8.5.02:_Alkaline_Earth_Metals'_Chemical_Properties) The hydroxides of the alkaline earth metals follow the formula M(OH)₂ and show a solubility trend that increases from beryllium to barium, reflecting the decreasing and increasing ionic size down the group. Beryllium and magnesium hydroxides have low solubility, with Be(OH)₂ being amphoteric and dissolving in strong bases to form tetrahydroxoberyllates, while Mg(OH)₂ is sparingly soluble and used in antacids for its mild basicity. (Ca(OH)₂), often called milk of lime due to its suspension in , has moderate solubility (K_{sp} ≈ 5.5 × 10^{-6} at 25 °C) and is applied in for adjustment. Strontium and barium hydroxides are more soluble, with Ba(OH)₂ having K_{sp} ≈ 5.0 × 10^{-3} at 25 °C, allowing it to function as a strong base in ; all M(OH)₂ compounds are white solids that decompose thermally to the and ./08:_Chemistry_of_the_Main_Group_Elements/8.05:_Group_2_The_Alkaline_Earth_Metals/8.5.02:_Alkaline_Earth_Metals'_Chemical_Properties) These oxygen-containing compounds form through direct reaction of the metals with oxygen: and magnesium yield primarily normal oxides via 2M + O₂ → 2MO, while and produce mixtures of oxides and peroxides (M + O₂ → MO + MO₂). Additionally, the oxides are commonly prepared by of the corresponding : MCO₃ → MO + CO₂, with the decomposition temperature increasing down the group due to the decreasing and polarizing power of the larger cations, which results in less distortion of the ions. The hydroxides, in turn, arise from the hydration of oxides, underscoring the basic character that defines the alkaline earth metals./Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements:_The_Alkaline_Earth_Metals/1Group_2:_Chemical_Reactions_of_Alkali_Earth_Metals/The_Thermal_Stability_of_the_Nitrates_and_Carbonates)

Halides and oxyhalides

The halides of alkaline earth metals generally adopt the formula MX₂, where M is the metal and X is the halogen, and exhibit ionic character for magnesium through barium, with structures influenced by ion sizes and packing efficiency. For instance, calcium fluoride (CaF₂) crystallizes in the fluorite structure, featuring a face-centered cubic array of Ca²⁺ ions with F⁻ ions occupying all tetrahedral voids, resulting in a coordination number of 8 for calcium and 4 for fluoride. Beryllium halides, however, display more covalent and polymeric characteristics due to the small size and high charge density of Be²⁺; beryllium chloride (BeCl₂), for example, forms infinite chains in the solid state, with each beryllium atom tetrahedrally coordinated to four chlorine atoms via bridging chlorides. These halides are typically prepared by direct combination of the metal with the gas, as alkaline earth metals react vigorously with X₂ to yield MX₂; for example, magnesium burns in to form MgCl₂./20%3A_Periodic_Trends_and_the_s-Block_Elements/20.05%3A_The_Alkaline_Earth_Metals_(Group_2)) Alternatively, they can be synthesized from the corresponding oxides by treatment with halides, following the reaction MO + 2HX → MX₂ + H₂O, which proceeds under heating for products./20%3A_Periodic_Trends_and_the_s-Block_Elements/20.05%3A_The_Alkaline_Earth_Metals_(Group_2)) Solubility in water for these halides shows an anomalous trend compared to typical group trends: fluorides are generally insoluble, with decreasing from to due to increasing lattice energies outweighing hydration energies, while chlorides, bromides, and iodides exhibit increasing down the group owing to decreasing lattice energies relative to hydration energies. fluoride (BeF₂) is exceptional among fluorides, forming a highly soluble, -like solid with a network of corner-sharing BeF₄ tetrahedra similar to silica ./20%3A_Periodic_Trends_and_the_s-Block_Elements/20.05%3A_The_Alkaline_Earth_Metals_(Group_2)) Iodides, in contrast, are notably soluble, with iodide (BaI₂) dissolving readily due to its low ./20%3A_Periodic_Trends_and_the_s-Block_Elements/20.05%3A_The_Alkaline_Earth_Metals_(Group_2)) Oxyhalides of alkaline earth metals, such as MOX₂, arise from partial halogenation and find applications in bleaching; calcium oxychloride (CaOCl₂), known as bleaching powder, is prepared industrially by passing chlorine gas over slaked lime via the reaction Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O, yielding a compound that releases hypochlorite in water for disinfection and whitening. Hygroscopicity of the halides increases down the group, as larger metal ions form more stable hydrates with decreasing lattice energies; magnesium chloride, for example, readily forms the hexahydrate MgCl₂·6H₂O and is highly moisture-absorbent, a property that intensifies for calcium chloride, which deliquesces in air./Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/1Group_2%3A_Chemical_Reactions_of_Alkali_Earth_Metals/Alkaline_Earth_(Group_II)_Trends)

Sulfates, carbonates, and other salts

The sulfates of the alkaline earth metals exhibit a marked decrease in in as one proceeds down the group from to . Beryllium sulfate (\ceBeSO4\ce{BeSO4}) and (\ceMgSO4\ce{MgSO4}) are highly soluble, while (\ceCaSO4\ce{CaSO4}) has limited solubility with a solubility product constant (KspK_{sp}) of 2.4×10−52.4 \times 10^{-5} at 25°C, and (\ceBaSO4\ce{BaSO4}) is essentially insoluble with Ksp=1.1×10−10K_{sp} = 1.1 \times 10^{-10} at 25°C. This trend is utilized in qualitative analysis, where the of white, insoluble barite (\ceBaSO4\ce{BaSO4}) serves as a confirmatory test for either or ions in solution. The carbonates of the alkaline earth metals are all insoluble in water, with calcium carbonate (\ceCaCO3\ce{CaCO3}) forming the primary component of limestone, a widespread sedimentary rock. Their thermal stability increases down the group due to the decreasing polarizing power of the larger cations, which weakens the distortion of the carbonate anion; for instance, beryllium carbonate (\ceBeCO3\ce{BeCO3}) decomposes at relatively low temperatures around 100°C, whereas barium carbonate (\ceBaCO3\ce{BaCO3}) requires heating above approximately 1000 °C to yield the oxide and carbon dioxide. The nitrates of the alkaline earth metals are highly soluble in water across the group, reflecting the weak lattice energies relative to hydration energies for these compounds. Upon heating, the nitrates decompose to the corresponding oxide, nitrogen dioxide, and oxygen: \ce2M(NO3)2−>2MO+4NO2+O2\ce{2M(NO3)2 -> 2MO + 4NO2 + O2}, with the decomposition temperature increasing slightly down the group. Among other important salts, the phosphates include calcium-based (\ceCa5(PO4)3(F,Cl,OH)\ce{Ca5(PO4)3(F,Cl,OH)}), a key mineral in phosphate rock deposits used for fertilizers and a structural component in certain materials. Silicates such as (\ceMg2SiO4\ce{Mg2SiO4}), the magnesium end-member of the series, are prevalent in ultramafic rocks and mantle-derived materials. These solubility trends in alkaline earth salts arise from the interplay between , which decreases down the group due to increasing cation size, and , which also diminishes but to a lesser extent for polyatomic anions like and ; the resulting net energy favors solubility for smaller cations with higher .

Reactions and identification

Representative reactions

The alkaline earth metals exhibit increasing reactivity down the group, as illustrated by their reactions with . Beryllium is inert toward cold but reacts slowly with hot or steam to form and gas. Magnesium reacts very slowly with cold but more readily with hot or steam, producing and gas according to the equation: \ceMg+2H2O−>Mg(OH)2+H2\ce{Mg + 2H2O -> Mg(OH)2 + H2} Calcium, strontium, and barium react vigorously even with cold water, evolving hydrogen gas and forming the corresponding metal hydroxides, as exemplified by calcium: \ceCa+2H2O−>Ca(OH)2+H2\ce{Ca + 2H2O -> Ca(OH)2 + H2} This trend reflects the decreasing ionization energy and increasing atomic size down the group, facilitating easier electron donation to water molecules./Descriptive_Chemistry/Main_Group_Reactions/Reactions_of_Main_Group_Elements_with_Water) All alkaline earth metals react with acids to liberate hydrogen gas, forming soluble metal salts, due to their reducing nature. For instance, magnesium reacts with hydrochloric acid as follows: \ceMg+2HCl−>MgCl2+H2\ce{Mg + 2HCl -> MgCl2 + H2} However, beryllium is an exception; it forms a passive oxide layer that protects it from reaction with dilute acids, though it dissolves in concentrated acids or those that complex the oxide. This anomaly arises from beryllium's high charge density, which promotes strong bonding in its oxide layer./Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements:_The_Alkaline_Earth_Metals/1Group_2:_Chemical_Reactions_of_Alkali_Earth_Metals/Reactions_of_Group_2_Elements_with_Acids) Except for , the alkaline earth metals react with at high temperatures to form with the general formula M₃N₂. These hydrolyze upon contact with water to yield the and gas. A representative example is calcium nitride: \ceCa3N2+6H2O−>3Ca(OH)2+2NH3\ce{Ca3N2 + 6H2O -> 3Ca(OH)2 + 2NH3} This reaction highlights the metals' ability to reduce , with reactivity increasing down the group due to more favorable for heavier members./20:_Periodic_Trends_and_the_s-Block_Elements/20.05:The_Alkaline_Earth_Metals(Group_2)) The combustion of alkaline earth metals in air produces their oxides, often with intense light emission. Magnesium exemplifies this, burning brightly to form : \ce2Mg+O2−>2MgO\ce{2Mg + O2 -> 2MgO} The brilliant white light results from the high of formation of MgO and the metal's volatility in flame, making it useful in . Heavier metals like calcium also combust but produce less intense light. In aqueous solutions, the coordination chemistry of alkaline earth metal ions varies with ionic size. The small Be²⁺ ion forms a tetrahedral aquo complex, [Be(H₂O)₄]²⁺, which is acidic due to : \ce[Be(H2O)4]2++H2O⇌[Be(H2O)3(OH)]++H3O+\ce{[Be(H2O)4]^2+ + H2O ⇌ [Be(H2O)3(OH)]+ + H3O+} In contrast, the larger ions of , , , and adopt octahedral coordination, as in [Mg(H₂O)₆]²⁺ or [Ca(H₂O)₆]²⁺, with weaker ligand fields and less pronounced acidity. This difference stems from beryllium's high charge-to-radius ratio, favoring four-coordinate geometry over six./20:_Periodic_Trends_and_the_s-Block_Elements/20.05:The_Alkaline_Earth_Metals(Group_2))

Qualitative identification of cations

Qualitative identification of alkaline earth metal cations (Be²⁺, Mg²⁺, Ca²⁺, Sr²⁺, Ba²⁺, and Ra²⁺) relies on a combination of classical wet chemical tests and instrumental methods, primarily exploiting differences in , color of precipitates, and spectral characteristics. In traditional qualitative analysis schemes, these cations are grouped as Group IV, separated from other metal ions through selective . The process begins in ammoniacal medium, where alkaline earth cations are precipitated as carbonates by adding ((NH₄)₂CO₃). , , and calcium form insoluble carbonates (BaCO₃, SrCO₃, CaCO₃), while magnesium and carbonates remain soluble due to higher solubility products, allowing their separation in the supernatant. The precipitated carbonates are then dissolved in acetic acid, and the individual cations are separated sequentially based on solubility differences of their sulfates. is first precipitated as insoluble (BaSO₄, white precipitate) by adding (Na₂SO₄). The filtrate is treated to precipitate as strontium sulfate (SrSO₄, white precipitate, less soluble than CaSO₄ but more than BaSO₄), and calcium remains in solution until confirmed separately. and magnesium, if present, do not precipitate in this sequence and are tested in the initial supernatant using other reagents like for Mg²⁺ (yellow precipitate). , due to its radioactivity and rarity, is not typically included in standard qualitative schemes and requires radiochemical separation followed by alpha spectrometry for identification./6%3A_Group_IV_and_Group_V_cations/6.2%3A_Separation_and_confirmation_of_individual_ions_in_group_IV_precipitates_and_group_V_mixture) Flame tests provide a rapid preliminary identification based on characteristic emission colors from excited atoms:
CationFlame Color
Be²⁺None
Mg²⁺None
Ca²⁺Brick-red (or orange-red)
Sr²⁺ (deep red)
Ba²⁺Apple green
Ra²⁺Not performed (radioactive)
These colors arise from electronic transitions in the valence electrons, with calcium, , and showing distinct hues due to their s² ./6%3A_Group_IV_and_Group_V_cations/6.2%3A_Separation_and_confirmation_of_individual_ions_in_group_IV_precipitates_and_group_V_mixture) Confirmatory tests involve specific reagents forming unique precipitates. For Ba²⁺, addition of chromate ions (CrO₄²⁻) produces a yellow precipitate of (BaCrO₄), insoluble in acetic acid, while strontium chromate (pale yellow, SrCrO₄) and calcium chromate (CaCrO₄) are sufficiently soluble and do not precipitate under these conditions. is confirmed by its sulfate precipitate or, more selectively, by a crimson after separation, as SrSO₄ is sparingly soluble but requires careful control to differentiate from calcium. Calcium is identified by precipitating (CaC₂O₄, white crystalline precipitate) with in neutral or slightly acidic solution; this precipitate is insoluble in acetic acid but dissolves in strong acids like HCl. Magnesium can be confirmed by forming a blue lake with p-nitrobenzeneazoresorcinol in alkaline medium, while yields a white gelatinous precipitate with NH₄OH due to Be(OH)₂./6%3A_Group_IV_and_Group_V_cations/6.2%3A_Separation_and_confirmation_of_individual_ions_in_group_IV_precipitates_and_group_V_mixture) Spectroscopic methods enhance specificity through (AES), where each cation exhibits characteristic line spectra. For example, calcium shows prominent emission lines at 422.7 nm (violet-blue) and 616.2 nm (red), at 460.7 nm (blue) and 421.5 nm (violet), and at 455.4 nm (blue) and 493.4 nm (green). These lines allow unambiguous identification even in mixtures, with detection limits in the ppm range using or AES. For separation and confirmation in complex samples, (IC) separates alkaline earth cations on cation-exchange columns using eluents like or EDTA-based systems, with conductometric detection providing baseline resolution (e.g., retention times increasing from Mg²⁺ to Ba²⁺ due to decreasing hydration). serves for semi-quantitative confirmation, as each cation forms stable 1:1 complexes with EDTA at 10–12, indicated by color change with indicator (red to blue for Mg²⁺/Ca²⁺). These methods are particularly useful for trace-level analysis in environmental or biological samples.

Applications

Metallurgical and structural uses

Alkaline earth metals play crucial roles in and structural applications, primarily through alloying to enhance mechanical properties, reduce weight, or modify microstructures in various materials. Magnesium and are valued for their lightweight characteristics in high-performance structures, while calcium, , and serve as effective modifiers in and non-ferrous alloys to improve castability, strength, and inclusion control. These uses leverage the metals' reactivity and atomic properties to tailor material performance without compromising integrity. Magnesium alloys, such as AZ91 (a composition of magnesium with approximately 9% aluminum and 1% ), are widely employed in automotive components like engine blocks and transmission cases, where they enable significant weight reductions—up to 75% lighter than equivalents—contributing to improved and . In aircraft applications, these alloys support structural elements requiring high strength-to-weight ratios, and their suitability for die-casting processes allows for complex, thin-walled parts with enhanced rigidity. Beryllium-copper alloys, typically containing about 2% , exhibit exceptional strength, conductivity, and fatigue resistance, making them ideal for precision springs and non-sparking tools used in hazardous environments like oil refineries and explosives handling. In nuclear reactors, these alloys serve as structural components due to beryllium's low absorption cross-section, which minimizes interference in and supports efficient . Calcium is primarily utilized as a desulfurizing agent in steel production, where it reacts with sulfur impurities to form calcium sulfide (CaS) inclusions that can be readily removed via slag formation, thereby improving steel cleanliness and ductility. Additionally, aluminum-calcium alloys are developed for overhead electrical conductors, offering a balance of high electrical conductivity and mechanical strength through deformation-processed nanocomposites that enhance tensile properties while maintaining low density. Strontium is incorporated into aluminum-strontium (Al-Sr) master alloys to refine grain structures in aluminum-silicon castings, promoting finer eutectic morphology and reducing for better mechanical uniformity and . , similarly, acts as an inoculant in production, nucleating formation to prevent fade during solidification and yielding with improved nodularity and tensile strength. A key trend in these applications distinguishes the lighter alkaline earth metals like magnesium and , which drive lightweighting in transportation and for energy efficiency, from the heavier ones—calcium, , and —that function as metallurgical modifiers to control inclusions, refine grains, and enhance cast quality in bulk alloys.

Chemical and industrial applications

Alkaline earth metals and their compounds play crucial roles in various chemical processes and industrial applications, leveraging their reactivity, properties, and unique physical characteristics. Calcium compounds, in particular, are foundational in and due to their abundance and chemical versatility. Calcium oxide (lime, CaO) is a primary ingredient in production, where it reacts with clay and other materials during high-temperature clinkering, followed by the addition of (3CaSO₄·2H₂O) to control setting time and form the final cement product. In , calcium hydroxide (Ca(OH)₂) is employed in processes to precipitate hardness-causing ions like calcium and magnesium carbonates, thereby reducing water hardness and improving quality for municipal and industrial use. Calcium oxide and hydroxide (lime) are also used in to neutralize acidic soils, raising and improving nutrient availability for crops. Magnesium's reactivity enables its use in organic synthesis through Grignard reagents (RMgX), which are organomagnesium halides formed by reacting magnesium with alkyl halides; these serve as powerful nucleophiles for carbon-carbon bond formation in the production of alcohols, hydrocarbons, and pharmaceuticals. Additionally, magnesium powder is utilized in for , where its high combustion energy produces intense white light and sparks upon ignition, enhancing visual effects in displays. Strontium compounds, such as and carbonate, are used in to produce crimson red colors in . Barium sulfate (BaSO₄) is a key additive in oil and gas drilling muds, functioning as a weighting agent to increase fluid density—typically up to around 80% of the required hydrostatic pressure—to prevent blowouts by countering formation pressures. (BaCO₃) has been historically applied as a , exploiting its toxicity to barium ions that disrupt channels and cause in pests like rats. Strontium ferrite (SrFe₁₂O₁₉) is widely used in permanent magnets for industrial applications such as electric motors and speakers, owing to its high , cost-effectiveness, and resistance to demagnetization in hexagonal structures. (SrO) is incorporated into glass formulations for cathode-ray tube (CRT) television screens, where it constitutes about 8% by weight in the faceplate to absorb X-rays and enhance shielding without compromising transparency. Radium bromide (RaBr₂) was historically mixed with (ZnS) to create luminous paints for watch dials and instrument panels in the early 20th century, relying on radium's to excite the for sustained glow-in-the-dark illumination.

Medical and specialized uses

plays a vital role in medical applications, particularly through supplements like (CaCO₃), which is widely used to address calcium deficiencies and support by increasing bone mineral density and reducing the risk of . is employed in specialized diagnostic procedures, such as selective intra-arterial calcium stimulation to localize insulinomas prior to , enhancing accuracy in endocrine evaluations. Magnesium compounds are integral to gastrointestinal and obstetric therapies; magnesium hydroxide (Mg(OH)₂) serves as an to neutralize stomach acid and alleviate , acid indigestion, and upset stomach. Magnesium sulfate (MgSO₄), commonly known as Epsom salt, acts as a to treat by drawing water into the intestines, and it is administered intravenously to prevent seizures in patients with by stabilizing neuronal membranes. Strontium-89, a radioactive of , is utilized in for bone metastases, particularly in and cancers, where it targets osteoblastic lesions as a beta emitter to provide significant and reduce the need for analgesics. Barium sulfate (BaSO₄) is a non-absorbable radiographic routinely used in upper and lower gastrointestinal imaging to visualize the , , and intestines during or fluoroscopic examinations, aiding in the diagnosis of conditions like ulcers and obstructions. Radium-226 (Ra-226) was historically applied in for treating cervical and other cancers using needles or tubes placed directly into tumors to deliver localized radiation, but it has largely been replaced by safer alternatives like due to radium's long and associated health risks. Beryllium exposure in medical and specialized contexts, such as in certain alloys for dental or applications, carries a significant precaution due to the risk of , a chronic characterized by granulomatous , necessitating strict exposure controls and medical surveillance for sensitized individuals.

Biological and environmental roles

Biological functions

Calcium is vital for numerous biological processes in living organisms, serving both structural and regulatory roles. Approximately 99% of the calcium in the is stored in bones and teeth, where it forms , \ceCa10(PO4)6(OH)2\ce{Ca10(PO4)6(OH)2}, providing rigidity and support to the skeletal system. Beyond its structural function, calcium ions act as second messengers in cellular signaling pathways, facilitating processes such as , nerve transmission, and hormone secretion; cytosolic calcium concentrations are tightly regulated at around 10−710^{-7} M to enable these transient signaling events. Magnesium is another essential alkaline earth metal, functioning primarily as a cofactor in enzymatic reactions and supporting metabolic processes. It participates in over 300 systems, notably stabilizing the ATP-Mg complex required for energy transfer in , , and synthesis. In photosynthetic organisms, magnesium occupies the central position in the ring, enabling light absorption and electron transport during . The typically contains about 25 g of magnesium, with roughly half residing in and the remainder distributed in soft tissues and fluids. Among the other alkaline earth metals, can partially substitute for calcium in processes, incorporating into structures like shells where it replaces calcium in lattices. In contrast, and lack essential biological roles and instead act as non-functional mimics of calcium, potentially disrupting normal physiological processes. Homeostatic mechanisms ensure adequate levels of these metals, particularly for calcium, which is regulated by to maintain serum concentrations through mobilization, intestinal absorption, and renal handling. daily intake recommendations are around 1000 mg for calcium and 400 mg for magnesium to support these functions. Deficiencies in these metals can impair health; insufficient calcium intake contributes to , a condition involving reduced and increased fracture risk. Similarly, , known as hypomagnesemia, is linked to cardiac arrhythmias due to disrupted balance and neuromuscular function.

Health precautions and environmental impact

exposure primarily affects the , leading to chronic disease (CBD), a condition characterized by due to an immunological response to inhaled particles. To prevent CBD and related acute effects, occupational safety standards set a of 0.2 μg/m³ as an 8-hour time-weighted average. Additionally, is classified as a probable carcinogen, with long-term exposure increasing risks of . For magnesium, —elevated serum magnesium levels—is uncommon and generally arises in cases of renal dysfunction or excessive supplementation, with symptoms including , , and emerging at levels above approximately 5 mEq/L and becoming severe beyond 10 mEq/L. Environmentally, magnesium extraction through generates runoff that can elevate metal concentrations in nearby water bodies, potentially disrupting aquatic ecosystems. However, magnesium-based compounds like show promise in mitigating by enhancing seawater and reducing dissolved CO₂. Excess calcium absorption can result in hypercalcemia, a condition linked to the formation of stones through increased urinary calcium excretion. In contrast, moderate calcium levels from offer health benefits by contributing to daily mineral intake and supporting health without significant risk. , a radioactive isotope from , chemically mimics calcium and preferentially accumulates in tissue, where its beta emissions elevate the risk of bone and marrow cancers. Barium ions (Ba²⁺) pose acute cardiac risks by blocking inward rectifier channels, which disrupts membrane potentials and induces hypokalemia-like arrhythmias. Radium exposure, mainly via emission, targets bone tissue and induces sarcomas, as dramatically illustrated by the 1920s incident, where dial painters ingested -laced paint and suffered fatal osteonecrosis and malignancies. Ecologically, contaminates environments through uranium mine tailings, which release radionuclides into and if containment fails. (BaSO₄), widely used in , remains largely inert and immobile in soils due to low , posing minimal direct , though barite mining operations contribute to and via open-pit excavation.

Extensions

Isotopic variations and stability

The alkaline earth metals exhibit a range of isotopic compositions, with the number of stable isotopes increasing from lighter to heavier elements in the group. Beryllium has only one isotope, ^{9}Be, while magnesium possesses three isotopes: ^{24}Mg (abundance 78.99%), ^{25}Mg (10.00%), and ^{26}Mg (11.01%). Calcium features six isotopes, dominated by ^{40}Ca at approximately 96.94% abundance, alongside ^{42}Ca (0.65%), ^{43}Ca (0.14%), ^{44}Ca (2.09%), ^{46}Ca (0.004%), and ^{48}Ca (0.19%); it also has more than 20 known radioactive isotopes. has four isotopes: ^{84}Sr (0.56%), ^{86}Sr (9.86%), ^{87}Sr (7.00%), and ^{88}Sr (82.58%). Barium includes seven isotopes, ranging from ^{130}Ba to ^{138}Ba, with ^{138}Ba being the most abundant at 71.66%. , the heaviest in the group, has no isotopes, with all known isotopes being radioactive.
ElementAtomic Number (Z)Number of Stable IsotopesExample Cosmogenic/Radioactive Isotope
Beryllium41 (^{9}Be)^{10}Be (t_{1/2} = 1.39 \times 10^{6} years)
Magnesium123N/A (all stable are non-cosmogenic)
Calcium206^{41}Ca (t_{1/2} = 99,400 years)
Strontium384^{90}Sr (t_{1/2} = 28.8 years, fission product)
Barium567^{133}Ba (t_{1/2} = 10.55 years)
Radium880^{223}Ra (t_{1/2} = 11.4 days)
Isotopic stability in alkaline earth metals follows trends influenced by the odd-even nuclear pairing effect, where nuclei with even numbers of both protons () and neutrons () are generally more stable. Although all group 2 elements have even , lighter members like and magnesium exhibit fewer stable isotopes due to low and limited neutron-proton pairing opportunities compared to heavier even- counterparts calcium, strontium, and , which benefit from greater neutron-proton pairing and thus support more stable configurations. Certain isotopes of alkaline earth metals serve as valuable tracers in and environmental studies. For instance, the ^{26}Mg isotope, produced via the decay of short-lived ^{26}Al, is used in isochron dating of silicates to determine the age of early solar system materials, providing insights into planetary formation timelines. The long-lived radioactive ^{41}Ca is employed in (AMS) for tracing calcium dynamics in human diet and studies, allowing measurement of absorption and turnover rates over decades due to its 99,400-year and trace natural abundance. The ^{87}Sr/^{86}Sr isotopic ratio, where ^{87}Sr arises from decay, is widely applied in analysis for and migration studies, as it reflects geological sources in biological tissues like without significant . Artificial isotopes of alkaline earth metals have practical applications in and instrumentation. , with a of 11.4 days, is used in targeted alpha therapy for treating metastases in , where its alpha emissions deliver high localized doses to tumor sites while sparing surrounding tissue. Barium-133, half-life 10.55 years, serves as a calibration source for and equipment due to its well-characterized emissions at 356 keV and 81 keV. Additionally, , a fission byproduct with a 28.8-year , is notable for its beta emissions and role in historical studies, though its applications are limited by radiotoxicity.

Relativistic effects in heavier homologues

Radium serves as the heaviest confirmed homologue in the alkaline earth metal group, with theoretical extensions predicting properties for element 120 () within the superheavy . Relativistic effects become prominent in these heavier elements, primarily through the contraction of s-orbitals, which stabilizes the valence electrons and alters expected . In and its lighter congener , relativistic s-orbital contraction increases the first (IE) of radium to 5.28 eV, higher than barium's 5.21 eV, reversing the typical decrease down the group. This stabilization, akin to an , reduces the metallic character and reactivity of radium compared to barium. Radium exhibits a brilliant white luster when pure, contrasting with barium's gray-white appearance upon tarnishing, though the precise influence of relativistic effects on remains under study. For element 120, predictions indicate an electron configuration of [Og] 8s², with intensified relativistic effects further contracting the 8s orbital and enhancing the inert pair, potentially rendering it less metallic than radium. Isotopes of unbinilium are expected to be highly unstable, with alpha decay half-lives on the order of microseconds (less than 1 second), limiting opportunities for chemical characterization. Some theoretical models suggest relativistic influences could lower the melting point, possibly making it liquid near room temperature, though this remains speculative. As of 2025, a breakthrough in synthesizing livermorium-116 using a titanium-50 beam has improved prospects for element 120 production, with experiments planned to begin that year. Ongoing at facilities like GSI Helmholtz Centre and the () aims to synthesize element 120 through fusion reactions, such as ²⁴⁸Cm + ⁵⁴Cr at GSI, which targets the compound nucleus ³⁰²₁₂₀; no confirmed detections have been reported to date. Alternative approaches, including ²⁴⁹Cf + ⁵⁰Ti at , explore similar neutron-rich isotopes predicted to lie near the superheavy island. These efforts highlight the challenges posed by relativistic stabilization, which may deviate unbinilium's chemistry from classical alkaline earth behavior.

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