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{{distinguish|Alkali metal}}
{{Periodic table (alkaline earth metals)}}⏎
The '''alkaline earth metals''' are six [[chemical element]]s in column (group) 2 of the [[Periodic table]]. They are [[beryllium]] (Be), [[magnesium]] (Mg), [[calcium]] (Ca), [[strontium]] (Sr), [[barium]] (Ba), and [[radium]] (Ra).<ref name="redbook">{{RedBook2005|pages=51}}.</ref> The elements have very similar properties: they are all shiny, silvery-white, somewhat [[reactivity (chemistry)|reactive]] [[metal]]s at [[standard temperature and pressure]].<ref name="rsc">{{cite web|url=http://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupii_data.html |title=Visual Elements: Group 2–The Alkaline Earth Metals |author =[[Royal Society of Chemistry]] |work=Visual Elements |publisher=Royal Society of Chemistry|accessdate=13 January 2012}}</ref>

Structurally, they have in common an outer s- [[electron shell]] which is full;<ref name="rsc"/><ref>
{{cite web
 | title = Periodic Table: Atomic Properties of the Elements
 | publisher = [[National Institute of Standards and Technology]]
 | date = September 2010
 | work = nist.gov
 | url = http://www.nist.gov/pml/data/upload/periodic_table_composite_2010_nobleed.pdf
 | accessdate = 17 February 2012
}}
</ref><ref name="RubberBible84th">
{{cite book
 | editor = Lide, D. R.
 | title = CRC Handbook of Chemistry and Physics
 | edition = 84th
 | location = Boca Raton, FL
 | publisher = CRC Press
 | year = 2003
}}
</ref>
that is, this orbital contains its full complement of two electrons, which these elements readily lose to form [[cation]]s with [[electric charge|charge]] +2, and an [[oxidation state]] (oxidation number) of +2.<ref name="Greenwood&Earnshaw">{{Greenwood&Earnshaw2nd}}</ref>

All the discovered alkaline earth metals occur in nature.<ref name="webelements-occurrence">{{cite web|url = http://www.webelements.com/webelements/properties/text/image-flash/abund-crust.html|title = Abundance in Earth's Crust|publisher = WebElements.com|accessdate =14 April 2007}}</ref> Experiments have been conducted to attempt the synthesis of [[Unbinilium|element 120]], the next potential member of the group, but they have all met with failure.

== Characteristics ==

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

{| class="wikitable" border="1" cellpadding="3" cellspacing="0"
|-
!''[[Atomic number|Z]]'' !! [[Chemical element|Element]] !! [[Electron shell|No. of electrons/shell]] !! [[Electron configuration]]<ref group="n">[[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.</ref>
|-
| 4 || [[beryllium]] || 2, 2 || &#91;[[Helium|He]]&#93; 2s<sup>2</sup>
|-
| 12 || [[magnesium]] || 2, 8, 2 || &#91;[[Neon|Ne]]&#93; 3s<sup>2</sup>
|-
| 20 || [[calcium]] || 2, 8, 8, 2 || &#91;[[Argon|Ar]]&#93; 4s<sup>2</sup>
|-
| 38 || [[strontium]] || 2, 8, 18, 8, 2 || &#91;[[Krypton|Kr]]&#93; 5s<sup>2</sup>
|-
| 56 || [[barium]] || 2, 8, 18, 18, 8, 2 || &#91;[[Xenon|Xe]]&#93; 6s<sup>2</sup>
|-
| 88 || [[radium]] || 2, 8, 18, 32, 18, 8, 2 || &#91;[[Radon|Rn]]&#93; 7s<sup>2</sup>
|}

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 [[radioactive decay|radioactivity]];<ref name="rsc" /> thus, the presentation of its properties here is limited.

The alkaline earth metals are all silver-colored and soft, and have relatively low [[density|densities]], [[melting point]]s, and [[boiling point]]s. [[chemistry|In chemical terms]], all of the alkaline metals react with the [[halogen]]s to form the alkaline earth metal [[halide]]s, all of which are [[ionic crystal|ionic crystalline compounds]] (except for [[beryllium chloride]], which is [[covalent bond|covalent]]). All the alkaline earth metals except beryllium also react with water to form strongly [[alkali]]ne [[hydroxide]]s and, thus, should be handled with great care. The heavier alkaline earth metals react more vigorously than the lighter ones.<ref name="rsc"/> The alkaline metals have the second-lowest first [[ionization energy|ionization energies]] in their respective periods of the [[periodic table]]<ref name="RubberBible84th"/> because of their somewhat low [[effective nuclear charge]]s and the ability to attain a [[full outer shell]] configuration by losing just two [[electron]]s. The second ionization energy of all of the alkaline metals is also somewhat low.<ref name="rsc"/><ref name="RubberBible84th" />

[[Beryllium]] is an exception: It does not react with water or steam, and its halides are covalent. If beryllium did form compounds with an ionization state of +2, it would polarize electron clouds that are near it very strongly and would cause extensive [[orbital overlap]], since beryllium has a high charge density. All compounds that include beryllium have a covalent bond.<ref name=deGruyter>{{cite book
 | others = trans. rev. Eagleson, Mary
 | editor1-first = Hans-Dieter | editor1-last=Jakubke
 | editor2-first = Hans | editor2-last = Jeschkeit
 | title = Concise Encyclopedia Chemistry
 | publisher = Walter de Gruyter
 | location = Berlin
 | year = 1994}}</ref> Even the compound [[beryllium fluoride]], which is the most ionic beryllium compound, has a low melting point and a low electrical conductivity when melted.<ref name="Beryllium halide and pseudohalides">{{cite book|title=Advances in inorganic chemistry and radiochemistry, Volume 14|year=1972|publisher=Academic Press|location=New York|isbn=978-0-12-023614-5|pages=256–277|url=https://books.google.com/?id=VupzlLU9NB0C&pg=PA257&dq=beryllium+fluoride+covalent#v=onepage&q=beryllium%20fluoride%20covalent&f=false|author =Bell, N. A.|editor1=Emeléus, Harry Julius |editor2=Sharpe, A. G. |chapter=Beryllium halide and pseudohalides}}</ref><ref name="Beryllium chemistry">{{cite book|last=Walsh|first=Kenneth A.|title=Beryllium chemistry and processing|publisher=ASM International|isbn=978-0-87170-721-5|pages=99–102, 118–119|url=https://books.google.com/?id=3-GbhmSfyeYC&pg=PA119&dq=beryllium+fluoride+covalent#v=onepage&q=beryllium%20fluoride%20covalent&f=false|date=2009-08-01}}</ref><ref name="General analytical chemistry of beryllium">{{cite book|title=Chemical analysis of metals: a symposium|year=1987|publisher=ASTM|isbn=978-0-8031-0942-1|pages=74–75|url=https://books.google.com/?id=uaWTfwrG644C&pg=PA74&dq=beryllium+fluoride+covalent#v=onepage&q=beryllium%20fluoride&f=false|author =Hertz, Raymond K.|editor=Coyle, Francis T.|chapter=General analytical chemistry of beryllium}}</ref>

All the alkaline earth metals have two [[electron]]s in their valence shell, so the energetically preferred state of achieving a filled [[electron shell]] is to lose two electrons to form doubly [[electric charge|charged]] [[Positive (electricity)|positive]] [[ion]]s.

==== Compounds and reactions ====
The alkaline earth metals all react with the [[halogens]] to form ionic halides, such as [[calcium chloride]] ({{chem|Ca|Cl|2}}), as well as reacting with [[oxygen]] to form oxides such as [[strontium oxide]] ({{chem|Sr|O}}). Calcium, strontium, and barium react with water to produce [[hydrogen gas]] and their respective [[hydroxide]]s, and also undergo [[transmetalation]] reactions to exchange [[ligand]]s.

<div style="float: left; margin: 1px; font-size:85%;">
:{| class="wikitable sortable"
|+ Alkaline earth metals fluorides solubility-related constants{{refn|Energies are given in −kJ/mol, solubilities in mol/L; HE means "[[hydration energy]]".|group="n"}}
! Metal<br />
! M<sup>2+</sup><br />HE<br />{{sfn|Wiberg|Wiberg|Holleman|2001|pp=XXXVI–XXXVII}}{{clarify |reason=this sfn missing matching citation |date=April 2013}}
! F<sup>−</sup><br />HE<br />{{sfn|Wiberg|Wiberg|Holleman|2001|p=XXXVI}}{{clarify |reason=this sfn missing matching citation |date=April 2013}}
! "MF<sub>2</sub>"<br />unit<br />HE
! MF<sub>2</sub><br />lattice<br />energies<br />{{sfn|Lide|2004|p=12-23}} 
! Solubility<br />{{sfn|Wiberg|Wiberg|Holleman|2001|p=1073}}{{clarify |reason=this sfn missing matching citation |date=April 2013}}<!--|gives g/L values, did an easy conversation to make it more about chem than industry-->
|-
| Be
| 2,455
| 458
| 3,371
| 3,526
| soluble
|- 
| Mg
| 1,922
| 458
| 2,838
| 2,978
| 0.0012
|-
| Ca
| 1,577
| 458
| 2,493
| 2,651
| 0.0002
|-
| Sr
| 1,415
| 458
| 2,331
| 2,513
| 0.0008
|-
| Ba
| 1,361
| 458
| 2,277
| 2,373
| 0.006
|}
</div>
{{Clear}}

=== Physical and atomic ===

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

{| class="wikitable" style="text-align:center"
!Alkaline earth metal
!Standard [[atomic weight]]<br>([[unified atomic mass unit|u]]){{refn|The number given in [[bracket|parentheses]] refers to the [[Uncertainty#Measurements|measurement uncertainty]]. This uncertainty applies to the [[significant figure|least significant figure]](s) of the number prior to the parenthesized value (i.e., counting from rightmost digit to left). For instance, {{val|1.00794|(7)}} stands for {{val|1.00794|0.00007}}, whereas {{val|1.00794|(72)}} stands for {{val|1.00794|0.00072}}.<ref>{{cite web|url=http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html|title=Standard Uncertainty and Relative Standard Uncertainty|work=[[CODATA]] reference|publisher=[[National Institute of Standards and Technology]]|accessdate=26 September 2011}}</ref>|group=n}}<ref name="atomicweights2007">{{cite journal|last1=Wieser |first1=Michael E. |last2=Berglund |first2=Michael |year=2009 |title=Atomic weights of the elements 2007 (IUPAC Technical Report) |journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=81|issue=11|pages= 2131–2156 |publisher=[[International Union of Pure and Applied Chemistry|IUPAC]]|doi=10.1351/PAC-REP-09-08-03|url=http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf|accessdate=7 February 2012 }}</ref><ref name="atomicweights2009">{{cite journal |last1=Wieser |first1=Michael E.|last2=Coplen|first2=Tyler B. |year=2011|title=Atomic weights of the elements 2009 (IUPAC Technical Report)|journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=83 |issue=2 |pages=359–396|publisher=[[International Union of Pure and Applied Chemistry|IUPAC]]|doi=10.1351/PAC-REP-10-09-14|url=http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf|accessdate=11 February 2012 }}</ref>
![[Melting point]]<br>([[kelvin|K]])
![[Melting point]]<br>([[Celsius|°C]])
![[Boiling point]]<br>([[kelvin|K]])<ref name="RubberBible84th"/>
![[Boiling point]]<br>([[Celsius|°C]])<ref name="RubberBible84th"/>
![[Density]]<br>(g/cm<sup>3</sup>)
![[Electronegativity]]<br>([[Pauling scale|Pauling]])
!First [[ionization energy]]<br>([[Kilojoule per mole|kJ·mol<sup>−1</sup>]]) !! [[Covalent radius]]<br>([[picometre|pm]])<ref name="slater64">{{cite journal|last=Slater|first=J. C.|year=1964|title=Atomic Radii in Crystals|journal=[[Journal of Chemical Physics]]|volume=41|issue=10|pages=3199–3205|bibcode=1964JChPh..41.3199S|doi=10.1063/1.1725697}}</ref>
! colspan="2" | [[Flame test]] color
|-
| [[Beryllium]] || 9.012182(3) || 1560 || 1287 || 2742 || 2469 || 1.85 || 1.57 || 899.5 || 105 || White<ref name="Jensen">{{cite journal|last1=Jensen |first1=William B. |year=2003 |title=The Place of Zinc, Cadmium, and Mercury in the Periodic Table |journal=Journal of Chemical Education |volume=80 |issue=8 |pages=952–961 |publisher=[[American Chemical Society]] |doi=10.1021/ed080p952 |bibcode=2003JChEd..80..952J |url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |accessdate=2012-05-06 |deadurl=yes |archiveurl=https://web.archive.org/web/20100611152417/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |archivedate=2010-06-11 |df= }}</ref> ||
|-
| [[Magnesium]] || 24.3050(6) || 923 || 650 || 1363 || 1090 || 1.738 || 1.31 || 737.7 || 150 || Brilliant-white<ref name="rsc" /> ||
|-
| [[Calcium]] || 40.078(4) || 1115 || 842 || 1757 || 1484 || 1.54 || 1.00 || 589.8 || 180 || Brick-red<ref name="rsc" /> || [[File:FlammenfärbungCa.png|40px]]
|-
| [[Strontium]] || 87.62(1) || 1050 || 777 || 1655 || 1382 || 2.64 || 0.95 || 549.5 || 200 || Crimson<ref name="rsc" /> || [[File:FlammenfärbungSr.png|40px]]
|-
| [[Barium]] || 137.327(7) || 1000 || 727 || 2170 || 1897 || 3.594 || 0.89 || 502.9 || 215 || Apple-green<ref name="rsc" /> ||
|-
| [[Radium]] || [226]{{refn|The element does not have any stable [[nuclide]]s, and a value in brackets indicates the [[mass number]] of the longest-lived [[isotope]] of the element.<ref name="atomicweights2007"/><ref name="atomicweights2009"/>|group=n}} || 973 || 700 || 2010 || 1737 || 5.5 || 0.9 || 509.3 || 221 || Crimson red{{refn|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.<ref name="RaFlameTest">{{cite book | url =https://books.google.com/books/about/The_Radiochemistry_of_Radium.html?id=3cgQLgEACAAJ | title = The Radiochemistry of Radium | author1 = Kirby | first1 = H. W | last2 = Salutsky | first2 = Murrell L | year = 1964 |publisher=National Academies Press}}</ref>|group=n}}
|
|}

==== Nuclear stability ====
{{expand section|date=November 2012}}

All of the alkaline earth metals except magnesium and strontium have at least one naturally occurring [[radioisotope]]: [[beryllium-7]], [[beryllium-10]], and [[calcium-41]] are [[trace radioisotope]]s, [[calcium-48]] and [[barium-130]] have very long [[half-life|half-lives]] and, thus, occur naturally, and all [[isotopes of radium]] are [[radioactivity|radioactive]]. Calcium-48 is the lightest nuclide to undergo [[double beta decay]].<ref>{{cite journal|author1=G. Audi |author2=A. H. Wapstra |author3=C. Thibault |author4=J. Blachot |author5=O. Bersillon |year=2003 |title=The NUBASE evaluation of nuclear and decay properties |url=http://www.nndc.bnl.gov/amdc/nubase/Nubase2003.pdf |journal=[[Nuclear Physics A]] |volume=729 |issue= |pages=3–128 |doi=10.1016/j.nuclphysa.2003.11.001 |bibcode=2003NuPhA.729....3A |deadurl=yes }}</ref>

The natural radioisotope of calcium, calcium-48, makes up about 0.1874% of natural calcium,<ref>{{cite web |url=http://ie.lbl.gov/education/parent/Ca_iso.htm |title=Isotopes of Calcium (Z=20) |author =Richard B. Firestone |date=15 March 2010 |publisher=Lawrence Berkeley National Laboratory |accessdate=12 June 2012}}</ref> and, thus, natural calcium is weakly radioactive. Barium-130 makes up approximately 0.1062% of natural barium, and, thus, barium is weakly radioactive, as well.<ref>{{cite web|url=http://ie.lbl.gov/education/parent/Ba_iso.htm |title=Isotopes of Barium (Z=56) |author =Richard B. Firestone |date=15 March 2010 |publisher=Lawrence Berkeley National Laboratory |accessdate=12 June 2012}}</ref>

== History ==

=== Etymology ===

The alkaline earth metals are named after their [[oxide]]s, the ''alkaline earths'', whose old-fashioned names were [[beryllia]], [[Magnesium oxide|magnesia]], [[Calcium oxide|lime]], [[strontia]], and [[baryta]]. These oxides are basic (alkaline) when combined with water. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but [[chemical compound|compound]]s 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,<ref name="Krebs">{{cite book| url = https://books.google.com/?id=yb9xTj72vNAC| title = The history and use of our earth's chemical elements: a reference guide| author = Robert E. Krebs| publisher = Greenwood Publishing Group| year = 2006| isbn = 0-313-33438-2| pages=65–81}}</ref> thus supporting Lavoisier's hypothesis and causing the group to be named the ''alkaline earth metals''.

=== Discovery ===

The calcium compounds [[calcite]] and [[Lime (material)|lime]] have been known and used since prehistoric times.<ref name="minerals.usgs"/> The same is true for the beryllium compounds [[beryl]] and [[emerald]].{{sfn|Weeks|1968|p=535}} 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.<ref name="Weeks1">{{cite journal | doi = 10.1021/ed009p1046 | title = The discovery of the elements. X. The alkaline earth metals and magnesium and cadmium|year = 1932 | last1 = Weeks | first1 = Mary Elvira |authorlink1=Mary Elvira Weeks| journal = Journal of Chemical Education | volume = 9 |issue = 6 | pages = 1046|bibcode = 1932JChEd...9.1046W }}</ref><ref name="Weeks2">{{cite journal | doi = 10.1021/ed009p1386 | title = The discovery of the elements. XII. Other elements isolated with the aid of potassium and sodium: Beryllium, boron, silicon, and aluminum| year = 1932 |last1 = Weeks | first1 = Mary Elvira |authorlink1=Mary Elvira Weeks| journal = Journal of Chemical Education | volume = 9|issue = 8 | pages = 1386|bibcode = 1932JChEd...9.1386W }}</ref><ref name="Weeks3">{{cite journal | doi = 10.1021/ed010p79 | title = The discovery of the elements. XIX. The radioactive elements | year = 1933 | last1 = Weeks | first1 = Mary Elvira |authorlink1=Mary Elvira Weeks| journal = Journal of Chemical Education | volume = 10 | issue = 2 | pages = 79|bibcode = 1933JChEd..10...79W }}</ref>

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 [[Marie Curie|Curie]] and [[André-Louis Debierne]] also by electrolysis.<ref name="Weeks1"/><ref name="Weeks2"/><ref name="Weeks3"/>

==== Beryllium ====
[[File:Béryl var. émeraude sur gangue (Muzo Mine Boyaca - Colombie) 15.jpg|thumb|right|200px|[[Emerald]], a variety of beryl, the mineral that beryllium was first known to be in.]]
[[Beryl]], a mineral that contains beryllium, has been known since the time of the [[Ptolemaic Kingdom]] in Egypt.{{sfn|Weeks|1968|p=535}} Although it was originally thought that beryl was an [[aluminium silicate]],{{sfn|Weeks|1968|p=537}} beryl was later found to contain a then-unknown element when, in 1797, [[Louis-Nicolas Vauquelin]] dissolved [[aluminium hydroxide]] from beryl in an alkali.<ref>{{cite journal|journal = Annales de Chimie|url = https://books.google.com/books?id=dB8AAAAAMAAJ&pg=RA1-PA155|pages = 155–169| first = Louis-Nicolas|last = Vauquelin|title = De l'Aiguemarine, ou Béril; et découverie d'une terre nouvelle dans cette pierre| year = 1798| issue = 26}}</ref> In 1828, [[Friedrich Wöhler]]<ref>{{Cite journal|journal = Annalen der Physik|volume = 89|issue = 8|pages = 577–582|title = Ueber das Beryllium und Yttrium|first = Friedrich|last = Wöhler|authorlink = Friedrich Wöhler|doi = 10.1002/andp.18280890805|year = 1828|bibcode = 1828AnP....89..577W}}</ref> and [[Antoine Bussy]]<ref>{{cite journal|journal = Journal de Chimie Medicale| url = https://books.google.com/books?id=pwUFAAAAQAAJ&pg=PA456|pages=456–457| first = Antoine |last = Bussy| title = D'une travail qu'il a entrepris sur le glucinium| year = 1828| issue = 4}}</ref> 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.{{sfn|Weeks|1968|p=539}} 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.{{sfn|Weeks|1968|p=539}}

==== Magnesium ====

Magnesium was first produced by [[Humphry Davy|Sir Humphry Davy]] in England in 1808 using electrolysis of a mixture of magnesia and [[Mercury(II) oxide|mercuric oxide]].<ref name="Davy1808">{{cite journal| last = Davy |first = H. | year= 1808 | title = Electro-chemical researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia | journal = Philosophical Transactions of the Royal Society of London | volume = 98 | pages = 333–370|bibcode = 1808RSPT...98..333D | jstor=107302 |doi=10.1098/rstl.1808.0023| url = https://books.google.com/books?id=gpwEAAAAYAAJ&pg=102#v=onepage&q&f=false}}</ref> [[Antoine Bussy]] prepared it in coherent form in 1831. Davy’s first suggestion for a name was magnium,<ref name="Davy1808"/> but the name magnesium is now used.

==== Calcium ====

[[Lime (material)|Lime]] has been used as a material for building since 7000 to 14,000&nbsp;BCE,<ref name="minerals.usgs">{{cite web |title = Commodity report:Lime |publisher = United States Geological Survey |first = M. Michael | last = Miller | url = http://minerals.usgs.gov/minerals/pubs/commodity/lime/390498.pdf |accessdate= 2012-03-06}}</ref> and [[Lime kiln|kilns used for lime]] have been dated to 2,500&nbsp;BCE in [[Khafaja]], [[Mesopotamia]].<ref>{{cite book | url =https://books.google.com/?id=ryap1yyEGAgC&pg=PA4 | page = 4 |title =Lime Kilns and Lime Burning | isbn =978-0-7478-0596-0 | author1 =Williams | first1 =Richard | year =2004}}</ref><ref>{{cite book | url = https://books.google.com/books/about/Lime_and_limestone.html?id=vHQsGAKAdYoC| title = Lime and Limestone: Chemistry and Technology, Production and Uses | isbn = 978-3-527-61201-7 | author1 = Oates | first1 = J. A. H | date = 2008-07-01}}</ref> Calcium as a material has been known since at least the first century, as the [[Ancient Rome|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]],<ref>{{cite journal | author = Davy H | year = 1808 | title = Electro-chemical researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia | url = https://books.google.com/books?id=gpwEAAAAYAAJ&pg=102#v=onepage&q&f=false |journal = Philosophical Transactions of the Royal Society of London | volume = 98 | issue = | pages = 333–370|bibcode = 1808RSPT...98..333D | doi = 10.1098/rstl.1808.0023 }}</ref> after hearing that [[Jöns Jakob Berzelius]] had prepared a calcium amalgam from the electrolysis of lime in mercury.

==== Strontium ====

In 1790, physician [[Adair Crawford]], who had been working with barium, realized that Strontian ores showed different properties than other supposed ores of barium.<ref>{{cite journal |last=Crawford |first=Adair |year=1790 |title=On the medicinal properties of the muriated barytes |journal=Medical Communications |location=London |volume=2 |pages=301–359}}</ref> Therefore, he concluded that these ores contained new minerals, which were named ''strontites'' in 1793 by [[Thomas Charles Hope]], a chemistry professor at the [[University of Glasgow]],<ref>{{cite journal |author =Murray, T.|year=1993| title= Elemementary Scots: The Discovery of Strontium |journal = Scottish Medical Journal| volume = 38 |pages = 188–189 |pmid=8146640 |issue=6}}</ref> who confirmed Crawford's discovery. Strontium was eventually isolated in 1808 by Sir [[Humphry Davy]] by electrolysis of a mixture of [[strontium chloride]] and [[mercuric oxide]]. The discovery was announced by Davy on 30 June 1808 at a lecture to the Royal Society.<ref>{{cite book|last=Davy |first=Humphry |year=1808 |url=https://books.google.com/books?id=gpwEAAAAYAAJ&pg=102#v=onepage&q&f=falseElectro-chemical |title=researches on the decomposition of the earths; with observations on the metals obtained from the alkaline earths, and on the amalgam procured from ammonia |publisher=Philosophical Transactions of the Royal Society of London |volume=98|pages=333–370}}</ref>

==== Barium ====

[[File:6158M-barite2.jpg|thumb|right|150px|Barite, the material that was first found to contain barium.]]

[[Barite]], a mineral containing barium, was first recognized as containing a new element in 1774 by [[Carl Scheele]], although he was able to isolate only [[barium oxide]]. Barium oxide was isolated again two years later by [[Johan Gottlieb Gahn]]. Later in the 18th century, [[William Withering]] noticed a heavy mineral in the [[Cumberland]] lead mines, which are now known to contain barium. Barium itself was finally isolated in 1808 when [[Sir Humphry Davy]] used electrolysis with molten salts, and Davy named the element ''barium'', after [[baryta]]. Later, [[Robert Bunsen]] and [[Augustus Matthiessen]] isolated pure barium by electrolysis of a mixture of barium chloride and ammonium chloride.<ref>{{cite journal|doi = 10.1002/jlac.18550930301|title = Masthead|year = 1855|journal = Annalen der Chemie und Pharmacie|volume = 93|issue = 3|pages = fmi–fmi}}</ref><ref>{{cite journal|doi =10.1002/prac.18560670194|title =Notizen|year =1856|last1 =Wagner|first1 =Rud.|last2 =Neubauer|first2 =C.|last3 =Deville|first3 =H. Sainte-Claire|last4 =Sorel|last5 =Wagenmann|first5 =L.|last6 =Techniker|last7 =Girard|first7 =Aimé|journal =Journal für Praktische Chemie|volume =67|pages =490–508}}</ref>

==== Radium ====

While studying [[uraninite]], on 21 December 1898, [[Marie Curie|Marie]] and [[Pierre Curie]] discovered that, even after uranium had decayed, the material created was still radioactive. The material behaved somewhat similarly to [[Compounds of barium|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]].<ref>{{cite journal |year=1898|title=Sur une nouvelle substance fortement radio-active, contenue dans la pechblende (On a new, strongly radioactive substance contained in pitchblende)|journal=Comptes Rendus|volume= 127|pages= 1215–1217|url=http://www.aip.org/history/curie/discover.htm |accessdate=2009-08-01 |author1=Curie, Pierre |author2=Curie, Marie |author3=Bémont, Gustave }}</ref> Radium was named in 1899 from the word ''radius'', meaning ''ray'', as radium emitted power in the form of rays.<ref>{{cite web|url=http://www.etymonline.com/index.php?term=radium |title=radium |publisher=Online Etymology Dictionary |accessdate=20 August 2011}}</ref>

== Occurrence ==

[[File:Erdalkali.jpg|thumb|Series of alkaline earth metals.]]

Beryllium occurs in the earth's crust at a concentration of two to six [[parts per million]] (ppm),<ref name=Merck>{{cite book
 | author= Merck contributors
| editor1-last = O'Neil | editor1-first = Marydale J.
 | editor2-last = Heckelman | editor2-first = Patricia E.
 | editor3-last = Roman | editor3-first = Cherie B.
 | title = The Merck Index: An Encyclopedia of Chemicals, Drugs, and Biologicals
 | edition = 14th
 | publisher = Merck Research Laboratories, Merck & Co., Inc.
 | location = Whitehouse Station, NJ, USA
 | year = 2006
 | isbn = 0-911910-00-X}}</ref> 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.<ref name="emsley">{{cite book
 | title = Nature's Building Blocks: An A–Z Guide to the Elements
 | last = Emsley
 | first = John
 | publisher = Oxford University Press
 | year = 2001
 | location = Oxford, England, UK
 | isbn = 0-19-850340-7
 | chapter =
 | ref = CITEREFEmsley2001}}</ref><ref>{{cite web|url=http://www.webelements.com/periodicity/abundance_seawater/ |title=Abundance in oceans |work=Mark Winter, [[The University of Sheffield]] and WebElements Ltd, UK |publisher=WebElements |accessdate=6 August 2011}}</ref> However, in freshwater, beryllium is somewhat more common, with a concentration of 0.1 parts per billion.<ref>{{cite web|url=http://www.webelements.com/periodicity/abundance_stream/ |title=Abundance in stream water |work=Mark Winter, [[The University of Sheffield]] and WebElements Ltd, UK |publisher=WebElements|accessdate=6 August 2011}}</ref>

Magnesium and calcium are very common in the earth's crust, with calcium the fifth-most-abundant element, and magnesium the eighth. None of the alkaline earth metals are found in their elemental state, but magnesium and calcium are found in many rocks and minerals: magnesium in [[carnallite]], [[magnesite]], and [[dolomite]]; and calcium in [[chalk]], [[limestone]], [[gypsum]], and [[anhydrite]].<ref name="rsc"/>

Strontium is the fifteenth-most-abundant element in the Earth's crust. Most strontium is found in the minerals [[celestite]] and [[strontianite]].<ref name="usgs10">{{cite web|publisher = United States Geological Survey |accessdate = 2010-05-14|title = Mineral Commodity Summaries 2010: Strontium |first = Joyce A. |last = Ober |url =http://minerals.usgs.gov/minerals/pubs/commodity/strontium/mcs-2010-stron.pdf |format=PDF}}</ref> Barium is slightly less common, much of it in the mineral [[barite]].<ref name="ullman">{{cite book|author1=Kresse, Robert |author2=Baudis, Ulrich |author3=Jäger, Paul |author4=Riechers, H. Hermann |author5=Wagner, Heinz |author6=Winkler, Jocher |author7=Wolf, Hans Uwe |chapter=Barium and Barium Compounds|editor=Ullman, Franz|title=Ullmann's Encyclopedia of Industrial Chemistry|year=2007|publisher=Wiley-VCH|doi=10.1002/14356007.a03_325.pub2}}</ref>

Radium, being a [[decay product]] of [[uranium]], is found in all uranium-bearing [[ore]]s.<ref name="lanl">[http://periodic.lanl.gov/88.shtml "Radium"], Los Alamos National Laboratory. Retrieved on 2009-08-05.</ref> Due to its relatively short half-life,<ref>{{cite book | url =https://books.google.com/books?id=t-fpKQ54f44C&pg=PT115| pages = 115– | title = Radioactivity | isbn = 978-0-19-983178-4 | author1 = Malley |first1 = Marjorie C | date = 2011-08-25}}</ref> radium from the Earth's early history has decayed, and present-day samples have all come from the much slower decay of uranium.<ref name="lanl"/>

== Production ==
{{expand section|date=November 2012}}

[[File:Beryl-130023.jpg|thumb|upright=0.56|Emerald, a variety of [[beryl]], is a naturally occurring compound of beryllium.]]

Most beryllium is extracted from beryllium hydroxide. One production method is [[sintering]], done by mixing [[beryl]], [[sodium fluorosilicate]], and soda at high temperatures to form [[sodium fluoroberyllate]], [[aluminium oxide]], and [[silicon dioxide]]. A solution of sodium fluoroberyllate and [[sodium hydroxide]] in [[properties of water|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.<ref name=deGruyter/>

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]].<ref>{{cite book | url = https://books.google.com/books?id=5smDPzkw0wEC&pg=PA401 | title = Production of SrCO, by black ash process: Determination of reductive roasting parameters| pages = 401 | isbn = 9789054108290 | author1 = Kemal | first1 = Mevlüt | last2 = Arslan | first2 = V| last3 = Akar | first3 = A | last4 = Canbazoglu | first4 = M | year = 1996}}</ref>

To produce barium, barite ore is separated from quartz, sometimes by [[froth flotation]] methods, resulting in relatively pure barite. [[Carbon]] is then used to reduce the baryte into [[barium sulfide]], which is dissolved with other elements to form other compounds, such as [[barium nitrate]]. These in turn are thermally decompressed into [[barium oxide]], which eventually yields pure barium after a reaction with [[aluminium]].<ref name="ullman"/> The most important supplier of barium is [[China]], which produces more than 50% of world supply.<ref>{{cite web|author =Miller, M. M. |url=http://minerals.usgs.gov/minerals/pubs/commodity/barite/mcs-2012-barit.pdf |title=Barite |publisher=USGS.gov}}</ref>

== Applications ==

Beryllium is used mostly for military applications,<ref>{{Cite book | last1 = Petzow | first1 = G. N. | last2 = Aldinger | first2 = F. | last3 = Jönsson | first3 = S. | last4 = Welge | first4 = P. | last5 = Van Kampen | first5 = V. | last6 = Mensing | first6 = T. | last7 = Brüning | first7 = T. | chapter = Beryllium and Beryllium Compounds | doi = 10.1002/14356007.a04_011.pub2 | title = Ullmann's Encyclopedia of Industrial Chemistry | year = 2005 | isbn = 3527306730 | pmid =  | pmc = }}</ref> but there are other uses of beryllium, as well. In electronics, beryllium is used as a [[p-type semiconductor|p-type]] [[dopant]] in some semiconductors,<ref>{{Cite book|url=https://books.google.com/?id=oJs6nK3TZrwC&pg=PA104|page=104|title=High-power diode lasers|author =Diehl, Roland|publisher=Springer|year=2000|isbn=3-540-66693-1}}</ref> and [[beryllium oxide]] is used as a high-strength [[electrical insulator]] and [[heat conductor]].<ref>{{Cite web|url = http://www.purdue.edu/uns/html4ever/2005/050927.Solomon.nuclear.html|date = 27 September 2005|title = Purdue engineers create safer, more efficient nuclear fuel, model its performance |publisher = Purdue University|accessdate =18 September 2008}}</ref> Due to its light weight and other properties, beryllium is also used in mechanics when stiffness, light weight, and dimensional stability are required at wide temperature ranges.<ref>{{Cite book|url =https://books.google.com/?id=IpEnvBtSfPQC&pg=PA690| title = Metals handbook|chapter = Beryllium|first = Joseph R.|last = Davis|publisher = ASM International|year = 1998|isbn = 978-0-87170-654-6|pages = 690–691}}</ref><ref>{{Cite book|url=https://books.google.com/?id=6fdmMuj0rNEC&pg=PA62|page=62|title=Encyclopedia of materials, parts, and finishes|author =Schwartz, Mel M. |publisher=CRC Press|year=2002|isbn=1-56676-661-3}}</ref>

Magnesium has many different uses. One of its most common uses was in industry, where it has many structural advantages over other materials such as [[aluminium]], although this usage has fallen out of favor recently due to magnesium's flammability.<ref name="Gray">{{cite book|last=Gray|first=Theodore|authorlink=Theodore Gray|title=The Elements: A Visual Exploration of Every Known Atom in the Universe|year=2009|publisher=Black Dog & Leventhal Publishers|location=New York|isbn=978-1-57912-814-2}}</ref> Magnesium is also often alloyed with aluminium or [[zinc]] to form materials with more desirable properties than any pure metal.<ref name="BakerM. M. Avedesian1999">{{cite book|last1=Baker|first1=Hugh D. R.|last2=Avedesian|first2=Michael|title=Magnesium and magnesium alloys|year=1999|publisher=Materials Information Society|location=Materials Park, OH|isbn=0-87170-657-1|page=4}}</ref> Magnesium has many other uses in industrial applications, such as having a role in the production of [[iron]] and [[steel]], and the production of [[titanium]].<ref>{{Cite book | last1 = Amundsen | first1 = K. | last2 = Aune | first2 = T. K. | last3 = Bakke | first3 = P. | last4 = Eklund | first4 = H. R. | last5 = Haagensen | first5 = J. Ö. | last6 = Nicolas | first6 = C. | last7 = Rosenkilde | first7 = C. | last8 = Van Den Bremt | first8 = S. | last9 = Wallevik | first9 = O. 
| doi = 10.1002/14356007.a15_559 | chapter = Magnesium | title = Ullmann's Encyclopedia of Industrial Chemistry | year = 2003 | isbn = 3527306730 | pmid =  | pmc = }}</ref>

Calcium also has many uses. One of its uses is as a [[reducing agent]] in the separation of other metals from ore, such as [[uranium]]. It is also used in the production of the alloys of many metals, such as [[aluminium]] and [[copper]] alloys, and is also used to deoxidize alloys as well. Calcium also has a role in the making of [[cheese]], [[Mortar (masonry)|mortar]]s, and [[cement]].<ref>{{RubberBible86th}}</ref>

Strontium and barium do not have as many applications as the lighter alkaline earth metals, but still have uses. [[Strontium carbonate]] is often used in the manufacturing of red [[firework]]s,<ref>{{cite journal | doi =10.1016/j.atmosenv.2006.09.019 | title =Recreational atmospheric pollution episodes: Inhalable metalliferous particles from firework displays | year =2007 | last1 =Moreno | first1 =Teresa | last2 =Querol | first2 =Xavier | last3 =Alastuey | first3 =Andrés | last4 =Cruz Minguillón | first4 =Mari | last5 =Pey | first5 =Jorge | last6 =Rodriguez | first6 =Sergio | last7 =Vicente Miró | first7 =José |last8 =Felis | first8 =Carles | last9 =Gibbons | first9 =Wes | journal =Atmospheric Environment | volume =41 |issue =5 | page =913|bibcode = 2007AtmEn..41..913M }}</ref> and pure strontium is used in the study of [[neurotransmitter]] release in neurons.<ref>{{cite journal |doi=10.1038/2121233a0 |pmid=21090447|title=Strontium as a Substitute for Calcium in the Process of Transmitter Release at the Neuromuscular Junction |year=1966 |last=Miledi |first= R. |journal=Nature |volume=212 |issue=5067 |pages=1233–4 |bibcode = 1966Natur.212.1233M }}</ref><ref>{{cite journal |author1=Hagler D.J., Jr |author2=Goda Y. |title= Properties of synchronous and asynchronous release during pulse train depression in cultured hippocampal neurons |journal= J. Neurophysiol. |year =2001 |volume =85|pmid=11387379 |issue=6 |pages=2324–34}}</ref> Barium has some use in [[vacuum tube]]s to remove gases,<ref name="ullman"/> and [[barium sulfate]] has many uses in the [[petroleum]] industry,<ref name="RubberBible84th"/> as well as other industries.<ref name="RubberBible84th"/><ref name="ullman"/><ref>{{cite book| page = 102| url = https://books.google.com/?id=uEJHsZWyO-EC| title= Medicinal applications of coordination chemistry|author1=Jones, Chris J.  |author2=Thornback, John | publisher =Royal Society of Chemistry| year = 2007| isbn =0-85404-596-1}}</ref>

Due to its radioactivity, radium no longer has many applications, but it used to have many. Radium used to be used often in [[luminous paint]]s,<ref name="PMC2024184">{{cite journal|title=Radium in the healing arts and in industry: Radiation exposure in the United States|pmc=2024184|year=1954|volume=69|issue=3|pmid=13134440|last1=Terrill Jr|first1=JG|last2=Ingraham Sc|first2=2nd|last3=Moeller|first3=DW|pages=255–62|journal=Public Health Reports|doi=10.2307/4588736}}</ref> although this use was stopped after workers got sick.<ref>{{cite web|url=http://www.radford.edu/~wkovarik/envhist/radium.html |title=Mass Media & Environmental Conflict&nbsp;– Radium Girls|accessdate=2009-08-01}}</ref> As people used to think that radioactivity was a good thing, radium used to be added to [[drinking water]], [[toothpaste]], and many other products, although they are also not used anymore due to their health effects.<ref name="Gray"/> Radium is no longer even used for its radioactive properties, as there are more powerful and safer emitters than radium.<ref>{{cite book | url = https://books.google.com/books?id=3cT2REdXJ98C&pg=PA24| page =24 | title = Radiation source use and replacement: Abbreviated version | isbn = 978-0-309-11014-3 | author1 = Committee On Radiation Source Use And Replacement | first1 = National Research Council (U.S.) | last2 = Nuclear And Radiation Studies Board | first2 = National Research Council (U.S.) | date = January 2008}}</ref><ref>{{cite book | url = https://books.google.com/books?id=bk0go_-FO5QC&pg=PA8|pages =8 | title = Radiation therapy planning | isbn = 978-0-07-005115-7 | author1 = Bentel | first1 = Gunilla Carleson | year = 1996}}</ref>

== Biological role and precautions ==
{{expand section|date=January 2013}}

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 transporter|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 [[exoskeleton]]s. It and barium have some uses in medicine, for example "[[barium meal]]s" in radiographic imaging, whilst strontium compounds are employed in some [[toothpaste]]s. Excessive amounts of strontium-90 are toxic due to its radioactivity and strontium-90 mimics calcium and then can kill.

Beryllium and radium, however, are toxic. Beryllium's low aqueous solubility means it is rarely available to biological systems; it has no known role in living organisms and, when encountered by them, is usually highly toxic.<ref name=deGruyter/> Radium has a low availability and is highly radioactive, making it toxic to life.

== Extensions ==

The next alkaline earth metal after radium is thought to be [[Unbinilium|element 120]], although this may not be true due to [[relativistic quantum chemistry|relativistic effects]].<ref name="tanm">{{cite web |url=http://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf|title=Gas Phase Chemistry of Superheavy Elements |author =Gäggeler, Heinz W. |date=5–7 November 2007 |work=Lecture Course Texas A&M |accessdate=26 February 2012}}</ref> 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 [[barn (unit)|fb]] for the cross-section at the energy studied.<ref>{{cite journal|journal=Phys. Rev. C|volume=79|page=024603|year=2009|title=Attempt to produce element 120 in the <sup>244</sup>Pu+<sup>58</sup>Fe reaction|doi=10.1103/PhysRevC.79.024603|last1=Oganessian|first1=Yu. Ts.|last2=Utyonkov|first2=V.|last3=Lobanov|first3=Yu.|last4=Abdullin|first4=F.|last5=Polyakov|first5=A.|last6=Sagaidak|first6=R.|last7=Shirokovsky|first7=I.|last8=Tsyganov|first8=Yu.|last9=Voinov|first9=A.|issue=2|bibcode = 2009PhRvC..79b4603O }}</ref> In April 2007, a team at the [[Gesellschaft für Schwerionenforschung|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.<ref name=Duellmann>http://fias.uni-frankfurt.de/kollo/Duellmann_FIAS-Kolloquium.pdf</ref>

The chemistry of element 120 is predicted to be closer to that of [[calcium]] or [[strontium]]<ref name=EB/> instead of [[barium]] or [[radium]]. This is unusual as [[periodic trends]] would predict element 120 to be more reactive than barium and radium. This lowered [[reactivity (chemistry)|reactivity]] is due to the expected energies of element 120's valence electrons, increasing element 120's [[ionization energy]] and decreasing the [[metallic radius|metallic]] and [[ionic radius|ionic radii]].<ref name=EB>{{cite web|author =Seaborg, G. T.|url=http://www.britannica.com/EBchecked/topic/603220/transuranium-element|title=transuranium element (chemical element)|publisher=Encyclop&aelig;dia Britannica|date=c. 2006|accessdate=16 March 2010}}</ref>

Although simple extrapolation would put element 170 as the next member of this series, calculations suggest that the next element of this series might actually be element 166.

== Notes ==
{{reflist|group="n"}}

== References ==
{{reflist|30em}}

== Bibliography ==
* {{cite book
 | last = Weeks
 | first = Mary Elvira
 | authorlink=Mary Elvira Weeks
 | last2= Leichester|first2=Henry M.
 | year = 1968
 | title = Discovery of the Elements
 | publisher = Journal of Chemical Education
 | location = Easton, PA
 | chapter =
 | id = LCCCN 68-15217
 | ref = CITEREFWeeks1968
 }}

== Further reading ==
* [http://www.rsc.org/chemsoc/visualelements/pages/data/intro_groupii_data.html Group 2 – Alkaline Earth Metals], Royal Chemistry Society.
* Hogan, C.Michael. 2010. [https://web.archive.org/web/20120612123626/http://www.eoearth.org/article/Calcium?topic=49557 ''Calcium''. eds. A.Jorgensen, C. Cleveland. Encyclopedia of Earth]. National Council for Science and the Environment.
* Maguire, Michael E. "Alkaline Earth Metals." ''Chemistry: Foundations and Applications''. Ed. [[J. J. Lagowski]]. Vol. 1. New York: Macmillan Reference USA, 2004. 33–34. 4 vols. Gale Virtual Reference Library. Thomson Gale.
* Silberberg, M.S., Chemistry: The molecular nature of Matter and Change (3e édition, McGraw-Hill 2009)
* Petrucci R.H., Harwood W.S. et Herring F.G., General Chemistry (8e édition, Prentice-Hall 2002)

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