Together withhelium, these elements have in common an outers 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 formcations withcharge +2, and anoxidation state of +2.[5] Helium is grouped with thenoble 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 thedecay chain ofuranium andthorium and not as a primordial element.[9] There have been experiments, all unsuccessful, to try to synthesizeelement 120, the next potential member of the group.
As with other groups, the members of this family show patterns in theirelectronic configuration, especially the outermost shells, resulting in trends in chemical behavior:
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 itsradioactivity;[2] thus, the presentation of its properties here is limited.
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 extensiveorbital overlap, since beryllium has a high charge density. All compounds that include beryllium have a covalent bond.[11] Even the compoundberyllium 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 twoelectrons in their valence shell, so the energetically preferred state of achieving a filledelectron shell is to lose two electrons to form doublychargedpositiveions.
The alkaline earth metals all react with thehalogens to form ionic halides, such ascalcium chloride (CaCl 2), as well as reacting withoxygen to form oxides such asstrontium oxide (SrO). Calcium, strontium, and barium react with water to producehydrogen gas and their respectivehydroxides (magnesium also reacts, but much more slowly), and also undergotransmetalation reactions to exchangeligands.
Solubility-related constants for alkaline-earth-metal fluorides
Isotopes of all six alkaline earth metals are present in theEarth's crust and theSolar System at varying concentrations, dependent upon the nuclides' half-lives and, hence, their nuclear stabilities. The first five haveone,three,five,four, andsix 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 throughdouble-beta decay, though no decays attributed definitively to these isotopes have yet been observed as of 2024. Radium has no stable norprimordial isotopes.
In addition to the stable species, calcium and barium each have one extremely long-lived andprimordial radionuclide:calcium-48 and barium-130, with half-lives of5.6×1019 and1.6×1021 years, respectively. Both are far longer than the currentage of the universe (4.7× and 117× billion times longer, respectively) and less than one part per ten billion has decayed since theformation 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 knownradioisotopes. None of the isotopes other than the aforementioned 21 areprimordial: 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 nearbysupernovae andcollisions between neutron stars, and any present are derived from ongoing natural processes.Beryllium-7,beryllium-10, andcalcium-41 aretrace, as well ascosmogenic, nuclides, formed by the impact ofcosmic 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 forradium-226 (radium's longest-lived isotope), 28.90 years forstrontium-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 throughdouble beta emission[n 5] and have extremely longhalf-lives, by virtue of the extremely low probability of both beta decays occurring at the same time. Allisotopes of radium are highlyradioactive and are primarily generated through the decay of heavier radionuclides. The longest-lived of them is radium-226, a member of thedecay chain ofuranium-238.[27] Strontium-90 and barium-140 are commonfission 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 onuranium-235 orplutonium-239 fuel, and a minusculesecular equilibrium concentration is also present due to rarespontaneous fission decays in naturally occurring uranium.
Calcium-48 is the lightest nuclide known to undergodouble 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 ofbaryte (barium sulfate).[33]
The longest-lived isotope of radium isradium-226 with a half-life of 1600 years; it, along withradium-223, -224, and -228, occurs naturally in thedecay chains of primordialthorium anduranium.Beryllium-8 is notable by its absence as it splits in half virtually instantaneously into twoalpha particles whenever it is formed. Thetriple 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, formingcarbon-12. This thermonuclear rate-limiting bottleneck is the reason mostmain sequence stars spend billions of yearsfusing 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 ofhydroxyapatite in compactbone, and gradually accumulate in the human skeleton. The incorporated radionuclides inflict significant damage to thebone marrow over time through the emission of ionizing radiation, primarilyalpha particles. This property is made use of in a positive manner in theradiotherapy of certainbone 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 aneven number of protons, owing to their status as group 2 elements. Their isotopes are generally more stable due tonucleon 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.
The alkaline earth metals are named after theiroxides, thealkaline earths, whose old-fashioned names wereberyllia,magnesia,lime,strontia, andbaria. 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 butcompounds is attributed to the chemistAntoine Lavoisier. In hisTraité 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 byelectrolysis of their molten earths,[35] thus supporting Lavoisier's hypothesis and causing the group to be named thealkaline earth metals.
The calcium compoundscalcite andlime have been known and used since prehistoric times.[36] The same is true for the beryllium compoundsberyl andemerald.[37] The other compounds of the alkaline earth metals were discovered starting in the early 15th century. The magnesium compoundmagnesium sulfate was first discovered in 1618 by a farmer atEpsom in England. Strontium carbonate was discovered inminerals in the Scottish village ofStrontian in 1790. The last element is the least abundant: radioactiveradium, which was extracted fromuraninite in 1898.[38][39][40]
All elements except beryllium were isolated by electrolysis of molten compounds. Magnesium, calcium, and strontium were first produced byHumphry Davy in 1808, whereas beryllium was independently isolated byFriedrich Wöhler andAntoine Bussy in 1828 by reacting beryllium compounds with potassium. In 1910, radium was isolated as a pure metal byCurie andAndré-Louis Debierne also by electrolysis.[38][39][40]
Magnesium was first produced byHumphry Davy in England in 1808 using electrolysis of a mixture of magnesia andmercuric 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.
Lime has been used as a material for building since 7000 to 14,000 BCE,[36] andkilns used for lime have been dated to 2,500 BCE inKhafaja,Mesopotamia.[47][48] Calcium as a material has been known since at least the first century, as theancient Romans were known to have usedcalcium 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, whenHumphry Davy, inEngland, usedelectrolysis on a mixture of lime andmercuric oxide,[49] after hearing thatJöns Jakob Berzelius had prepared a calcium amalgam from the electrolysis of lime in mercury.
In 1790, physicianAdair Crawford discovered ores with distinctive properties, which were namedstrontites in 1793 byThomas Charles Hope, a chemistry professor at theUniversity of Glasgow,[50] who confirmed Crawford's discovery. Strontium was eventually isolated in 1808 byHumphry Davy by electrolysis of a mixture ofstrontium chloride andmercuric oxide. The discovery was announced by Davy on 30 June 1808 at a lecture to the Royal Society.[51]
While studyinguraninite, on 21 December 1898,Marie andPierre Curie discovered that, even after uranium had decayed, the material created was still radioactive. The material behaved somewhat similarly tobarium 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 theFrench Academy of Sciences.[54] Radium was named in 1899 from the wordradius, meaningray, as radium emitted power in the form of rays.[55]
Beryllium occurs in the Earth's crust at a concentration of two to sixparts 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 asscandium, 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 arecarnallite,magnesite, anddolomite. Common calcium-containing minerals arechalk,limestone,gypsum, andanhydrite.[2]
Strontium is the 15th most abundant element in the Earth's crust. The principal minerals arecelestite andstrontianite.[60] Barium is slightly less common, much of it in the mineralbarite.[61]
Radium, being adecay product ofuranium, is found in all uranium-bearingores.[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]
Most beryllium is extracted from beryllium hydroxide. One production method issintering, done by mixingberyl,sodium fluorosilicate, and soda at high temperatures to form sodiumfluoroberyllate,aluminum oxide, andsilicon dioxide. A solution of sodium fluoroberyllate andsodium hydroxide inwater is then used to formberyllium hydroxide by precipitation. Alternatively, in the melt method, powdered beryl is heated to high temperature, cooled with water, then heated again slightly insulfuric acid, eventually yielding beryllium hydroxide. The beryllium hydroxide from either method then producesberyllium fluoride andberyllium chloride through a somewhat long process. Electrolysis or heating of these compounds can then produce beryllium.[11]
In general, strontium carbonate is extracted from the mineralcelestite through two methods: by leaching the celestite withsodium carbonate, or in a more complicated way involvingcoal.[64]
To produce barium,barite (impure barium sulfate) is converted tobarium sulfide bycarbothermic reduction (such as withcoke). The sulfide is water-soluble and easily reacted to form pure barium sulfate, used for commercial pigments, or other compounds, such asbarium nitrate. These in turn arecalcined intobarium oxide, which eventually yields pure barium after reduction withaluminum.[61] The most important supplier of barium isChina, which produces more than 50% of world supply.[65]
Magnesium is usually produced frommagnesite ore, as well asdolomite. 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 is the fifth most abundant element element in the Earth's crust, existing mostly as calcium carbonate. Other important minerals of calcium include gypsum, anhydrite, fluorite, and apatite.[5] The largest producers of calcium in the world are China and Russia, where Davy's method of electrolysis is used on calcium chloride to produce the metal. Other large producers of calcium are the U.S. and Canada, where lime is reduced with aluminium at high temperatures.[67]
Marie Curie developed the first methods of radium extraction using residues of pitchblende, following her extraction of uranium. These residues contained mostly barium sulfate. A lengthy purification process involving boiling with sodium hydroxide, treating with hydrochloric acid, then adding sodium carbonate afforded a mixture of barium and radium carbonates, which could be converted to chlorides, and from which the radium salt could be separated via fractional crystallization.[68] Curie's method was used until 1940, after which mixed bromides became more popular.[69] As of 2011, radium is extracted only from spent reactor fuel as oxides.[70] The pure metal can be isolated by reduction with aluminium in a vacuum at 1,200 °C.[71]
Beryllium is used mainly in military applications,[72] but non-military uses exist. In electronics, beryllium is used as ap-typedopant in some semiconductors,[73] andberyllium oxide is used as a high-strengthelectrical insulator andheat conductor.[74] Beryllium alloys are used for mechanical parts when stiffness, light weight, and dimensional stability are required over a wide temperature range.[75][76] Beryllium-9 is used in small-scaleneutron sources that use the reaction9Be +4He (α) →12C +1n, the reaction used byJames Chadwick when hediscovered the neutron. Its low atomic weight and low neutron absorption cross-section would make beryllium suitable as aneutron moderator, but its high price and the readily available alternatives such as water,heavy water andnuclear graphite have limited this to niche applications. In theFLiBe eutectic used inmolten 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 asaluminum, but magnesium's usage is hindered by its flammability.[77] Magnesium is oftenalloyed with aluminum, zinc and manganese to increase its strength and corrosion resistance.[78] Magnesium has many other industrial applications, such as its role in the production ofiron andsteel,[further explanation needed] and in theKroll process for production oftitanium.[79]
Calcium is used as areducing agent in the separation of other metals such asuranium from ore. It is a major component of many alloys, especiallyaluminum andcopper alloys, and is also used to deoxidize alloys. Calcium has roles in the making ofcheese,mortars, andcement.[80]
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 inluminous paints,[88] although this use was stopped after it sickened workers.[89] Thenuclear quackery that alleged health benefits of radium formerly led to its addition todrinking water,toothpaste, and many other products.[77] Radium is no longer used even when its radioactive properties are desired because its long half-life makes safe disposal challenging. For example, inbrachytherapy, shorter-lived alternatives such asiridium-192 are usually used instead.[90][91]
Anhydrous calcium chloride is ahygroscopic 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 asdeliquescence.
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 formhydroxide andhydrogen gas. Be and Mg arepassivated 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.
The table below[92] presents the colors observed when the flame of aBunsen burner is exposed to salts of alkaline earth metals. Be and Mg do not impart colour to the flame due to their small size.[93]
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.
They are generated from the corresponding oxides on reaction with water. They exhibit basic character: they turnphenolphthalein pink andlitmus, blue. Beryllium hydroxide is an exception as it exhibitsamphoteric 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 asdolomite,aragonite,magnesite (carbonate rocks). Calcium and magnesium ions are found inhard 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 ascalcium hydroxide,sodium carbonate orsodium phosphate. A more common method is to use ion-exchangealuminosilicates orion-exchange resins that trap Ca2+ and Mg2+ and liberate Na+ instead:
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 calciumion pumps playing a role in some cellular processes, magnesium functioning as the active center in someenzymes, 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 theirexoskeletons. It and barium have some uses in medicine, for example "barium meals" in radiographic imaging, whilst strontium compounds are employed in sometoothpastes. Excessive amounts ofstrontium-90 are toxic due to its radioactivity and strontium-90 mimics calcium (i.e. Behaves as a "bone seeker") where itbio-accumulates with a significantbiological half life. While the bones themselves have higher radiation tolerance than other tissues, the rapidly dividingbone marrow does not and can thus be significantly harmed by Sr-90. The effect ofionizing radiation on bone marrow is also the reason whyacute radiation syndrome can haveanemia-like symptoms and why donation ofred 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.
The next alkaline earth metal after radium is thought to beelement 120, although this may not be true due torelativistic effects.[94] The synthesis of element 120 was first attempted in March 2007, when a team at theFlerov Laboratory of Nuclear Reactions inDubna bombardedplutonium-244 withiron-58 ions; however, no atoms were produced, leading to a limit of 400fb for the cross-section at the energy studied.[95] In April 2007, a team at theGSI attempted to create element 120 by bombardinguranium-238 withnickel-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.[96]
The chemistry of element 120 is predicted to be closer to that ofcalcium orstrontium[97] instead ofbarium orradium. This noticeably contrasts withperiodic trends, which would predict element 120 to be more reactive than barium and radium. This loweredreactivity is due to the expected energies of element 120's valence electrons, increasing element 120'sionization energy and decreasing themetallic andionic radii.[97]
The next alkaline earth metal after element 120 has not been definitely predicted. Although a simple extrapolation using theAufbau 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 ingroup 12, belowcopernicium.[98][99]
^Noble gas notation is used for conciseness; the nearest noble gas that precedes the element in question is written first, and then the electron configuration is continued from that point forward.
^The number given inparentheses refers to themeasurement uncertainty. This uncertainty applies to theleast significant figure(s) of the number prior to the parenthesized value (i.e., counting from rightmost digit to left). For instance,1.00794(7) stands for1.00794±0.00007, whereas1.00794(72) stands for1.00794±0.00072.[19]
^The element does not have any stablenuclides, and a value in brackets indicates themass number of the longest-livedisotope of the element.[20][21]
^The color of the flame test of pure radium has never been observed; the crimson-red color is an extrapolation from the flame test color of its compounds.[25]
^Calcium-48 is theoretically capable of singlebeta decay, but such process has never been observed.[26]
^Bent Weberg, Libby (18 January 2019).""The" periodic table".Chemical & Engineering News.97 (3).Archived from the original on 1 February 2020. Retrieved27 March 2020.
^Clark, Jim (December 2021)."Reactions of the Group 2 elements with water". Retrieved2012-08-14.'Beryllium has no reaction with water or steam even at red heat.' This was commonly quoted in textbooks....However, a researcher...sent me a photo showing the result of exposing beryllium to steam at 800°C. It definitely reacts. I think the problem is that beryllium is both expensive and carries major health risks....Textbook (or these days, web) statements about it never get checked.
^abcJakubke, Hans-Dieter; Jeschkeit, Hans, eds. (1994).Concise Encyclopedia Chemistry. trans. rev. Eagleson, Mary. Berlin: Walter de Gruyter.
^Arblaster, John W. (2018).Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International.ISBN978-1-62708-155-9.
^Richard B. Firestone (15 March 2010)."Isotopes of Calcium (Z=20)". Lawrence Berkeley National Laboratory. Archived fromthe original on 6 May 2012. Retrieved12 June 2012.
^Richard B. Firestone (15 March 2010)."Isotopes of Barium (Z=56)". Lawrence Berkeley National Laboratory. Archived fromthe original on 6 May 2012. Retrieved12 June 2012.
^abMiller, M. Michael."Commodity report:Lime"(PDF). United States Geological Survey.Archived(PDF) from the original on 2011-11-12. Retrieved2012-03-06.
^abcdKresse, Robert; Baudis, Ulrich; Jäger, Paul; Riechers, H. Hermann; Wagner, Heinz; Winkler, Jocher; Wolf, Hans Uwe (2007). "Barium and Barium Compounds". In Ullman, Franz (ed.).Ullmann's Encyclopedia of Industrial Chemistry. Wiley-VCH.doi:10.1002/14356007.a03_325.pub2.ISBN978-3-527-30673-2.
^Baker, Hugh D. R.; Avedesian, Michael (1999).Magnesium and magnesium alloys. Materials Park, OH: Materials Information Society. p. 4.ISBN0-87170-657-1.
^Amundsen, K.; Aune, T. K.; Bakke, P.; Eklund, H. R.; Haagensen, J. Ö.; Nicolas, C.; Rosenkilde, C.; Van Den Bremt, S.; Wallevik, O. (2003). "Magnesium".Ullmann's Encyclopedia of Industrial Chemistry.doi:10.1002/14356007.a15_559.ISBN3-527-30673-0.
^Lide, D. R., ed. (2005).CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton, Florida: CRC Press.ISBN0-8493-0486-5.
^Hagler D.J., Jr; Goda Y. (2001). "Properties of synchronous and asynchronous release during pulse train depression in cultured hippocampal neurons".J. Neurophysiol.85 (6):2324–34.doi:10.1152/jn.2001.85.6.2324.PMID11387379.S2CID2907823.
^Fricke, B.; Greiner, W.; Waber, J. T. (1971). "The continuation of the periodic table up to Z = 172. The chemistry of superheavy elements".Theoretica Chimica Acta.21 (3):235–260.doi:10.1007/BF01172015.S2CID117157377.
^Hoffman, Darleane C.; Lee, Diana M.; Pershina, Valeria (2006). "Transactinides and the future elements". In Morss; Edelstein, Norman M.; Fuger, Jean (eds.).The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands:Springer Science+Business Media.ISBN978-1-4020-3555-5.
Hogan, C. Michael. 2010."Calcium". A. Jorgensen, C. Cleveland, eds.Encyclopedia of Earth. National Council for Science and the Environment.
Maguire, Michael E. "Alkaline Earth Metals".Chemistry: Foundations and Applications. Ed.J. J. Lagowski. Vol. 1. New York: Macmillan Reference USA, 2004. 33–34. 4 vols. Gale Virtual Reference Library. Thomson Gale.
Petrucci R.H., Harwood W.S., and Herring F.G., General Chemistry (8th edition, Prentice-Hall, 2002)
Silberberg, M.S.,Chemistry: The Molecular Nature of Matter and Change (3rd edition, McGraw-Hill, 2009)
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