Alkaline earth metal

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Group → 2
↓ Period
2 Lump of beryllium
4
Be
3 Magnesium crystals
12
Mg
4 Calcium stored under argon atmosphere
20
Ca
5 Strontium floating in paraffin oil
38
Sr
6 Barium stored under argon atmosphere
56
Ba
7 Radium electroplated on copper foil and covered with polyurethane to prevent reaction with air
88
Ra
Legend
Alkaline earth metal
Primordial
From decay

The alkaline earth metals are a group of chemical elements in the periodic table with very similar properties: they are all shiny, silvery-white, somewhat reactive metals at standard temperature and pressure[1] and readily lose their two outermost electrons to form cations with charge +2.[2] In the modern IUPAC nomenclature, the alkaline earth metals comprise the group 2 elements.[note 1]

The alkaline earth metals are beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra).[4] This group lies in the s-block of the periodic table as all alkaline earth metals have their outermost electron in an s-orbital.[1][5][6]

All the discovered alkaline earth metals occur in nature.[7] Experiments have been conducted to attempt the synthesis of unbinilium (Ubn), which is likely to be the next member of the group, but they have all met with failure. However, unbinilium may not be an alkaline earth metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements.[8]

Contents

[edit] Characteristics

[edit] Chemical

Series of alkaline earth metals.

Like other groups, the members of this family show patterns in its electronic configuration, especially the outermost shells, resulting in trends in chemical behavior:

Z Element No. of electrons/shell Electron configuration[note 2]
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
120 unbinilium 2, 8, 18, 32, 32, 18, 8, 2 (predicted)[9]:1722 [Uuo] 8s2 (predicted)[9]:1722

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,[1] and unbinilium has not yet been discovered; thus, the presentation of their properties here is limited.

H   He
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba * Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra ** Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Uuq Uup Uuh Uus Uuo
Uue Ubn  
  * La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
  ** Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr
Alkaline earth metals in the periodic table

The alkaline earth metals are all silver-colored,[1] soft, and have relatively low densities,[1] melting points,[1] and boiling points.[1] In chemical terms, all of the alkaline metals react with the halogens to form the alkaline earth metal halides, all of which ionic crystalline compounds (except for beryllium chloride, which is covalent).[1] 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.[1] The alkaline metals have the second-lowest first ionization energies in their respective periods of the periodic table[6] because of their somewhat low effective nuclear charges[1] 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.[1][6]

The chemistry of unbinilium, the undiscovered seventh alkali metal, is predicted to be closer to that of calcium or strontium or rubidium[10] instead of barium or radium. This is unusual as periodic trends would predict unbinilium to be more reactive than barium and radium. This lowered reactivity is due to the energetic properties[clarification needed]of unbinilium's valence electron, increasing unbinilium's ionisation energy and decreasing the metallic and ionic radii.[10]

Beryllium is an exception: It does not react with water or steam, and its halides are covalent. If the Be2+ ion did exist, it would polarize electron clouds that are near it very strongly and would cause extensive orbital overlap, since Be has a high charge density. All compounds that include Be have a covalent bond. Even the compound BeF2, which is the most ionic Be compound, has a low melting point and a low electrical conductivity when melted.

All the alkaline earth metals have two electrons in their valence shell, so the energetically preferred state of achieving a filled electron shell is to lose two electrons to form doubly charged positive ions.

[edit] Compounds and reactions

Reactions:
Note: E = elements that act as reducing agents

  1. The metals reduce halogens to form ionic halides: E(s) + X2 → EX2(s)where X = F, Cl, Br or I
  2. The metals reduce O2 to form the oxides: 2E(s) + O2 → 2EO(s)
  3. The larger metals react with water to produce hydrogen gas: E(s) + 2H2O(l) → E2+(aq) + 2OH-aq + H2 (g) where E = Ca, Sr or Ba
  4. The metals undergo transmetalation reactions to exchange ligands: Ae + Hg{N(SiMe3)2}2 → [Ae{N(SiMe3)2}2(THF)2] where Ae = Ca, Sr, or Ba.

Compounds

  1. Alkylmagnesium halides (RMgX where R = hydrocarbon group and X = halogen). They are used to synthetise organic compounds.
    Here’s an example: 3RMgCl + SnCl4 → 3MgCl2 + R3SnCl
  2. Magnesium oxide (MgO). It is used as a material to refract furnace brick and wire insulation (melting point of 2852°C).
  3. Calcium carbonate (CaCO3). It is mainly used in the construction industry and for making limestone, marble, chalk, and coral.

[edit] Physical and atomic

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

Alkaline earth metal Standard atomic weight
(u)[note 3][12][13]		 Melting point
(K)		 Melting point
(°C)		 Boiling point
(K)[6]		 Boiling point
(°C)[6]		 Density
(g/cm3)		 Electronegativity
(Pauling)		 First ionisation energy
(kJ·mol−1)		 Covalent radius
(pm)[14]		 Flame test color
Beryllium 9.012182(3) 1560.15 1287.00 2742 2469 1.85 1.57 899.5 105 Unknown 40px
Magnesium 24.3050(6) 923.15 650.00 1363 1090 1.738 1.31 737.7 150 Brilliant white[1] 40px
Calcium 40.078(4) 1112.15 839.00 1757 1484 1.54 1.00 589.8 180 Brick-red[1] FlammenfärbungCa.png
Strontium 87.62(1) 1042.15 769.00 1655 1382 2.64 0.95 549.5 200 Crimson[1] FlammenfärbungSr.png
Barium 137.327(7) 1002.15 729.00 2170 1897 3.594 0.89 502.9 215 Apple green[1] 40px
Radium [226][note 4] 973.15 700.00 2010 1737 5.5 0.9 509.3 221 Crimson red[citation needed] 40px

[edit] Radioactivity

All of the alkaline earth metals except magnesium and strontium have at least one naturally occurring radioisotope: beryllium-7, beryllium-10, and calcium-41 are trace radioisotopes, calcium-48 and barium-130 have very long half-lives and thus occur naturally, and all isotopes of radium are radioactive.

[edit] Other substances similar to the alkaline earth metals

Just like their names, the alkaline earth metals and the alkali metals do not differ completely. The main difference is the electron configuration, which is ns2 for alkaline earth metals and ns1 for alkali metals. For the alkaline earth metals, there are two electrons that are available to form a metallic bond, and the nucleus contains an additional positive charge. Also, the elements of group 2A (alkaline earth) have much higher melting points and boiling points compared to those of group 1A (alkali metals). The alkali also have a softer and more lighweight figure whereas the alkaline earth metals are much harder and denser.

The second valence electron is very important when it comes to comparing chemical properties of the alkaline earth and the alkali metals. The second valence electron is in the same “sublevel” as the first valence electron. Therefore, the Zeff is much greater. This means that the elements of the group 2A contain a smaller atomic radius and much higher ionization energy than the group 1A. Even though the group 2A contains much higher ionization energy, they still form an ionic compound with 2+ cations. Beryllium, however, behaves differently. This is because in order to remove two electrons from this particular atom, it requires significantly more energy. It never forms Be2+ and its bonds are polar covalent.

[edit] History

[edit] Etymology

The alkaline earth metals are named after their oxides, the alkaline earths, whose old-fashioned names were beryllia, magnesia, lime, strontia and baryta. These oxides are basic (alkaline) when combined with water. "Earth" is an old term applied by early chemists to nonmetallic substances that are insoluble in water and resistant to heating—properties shared by these oxides. The realization that these earths were not elements but compounds is attributed to the chemist Antoine Lavoisier. In his Traité Élémentaire de Chimie (Elements of Chemistry) of 1789 he called them salt-forming earth elements. Later, he suggested that the alkaline earths might be metal oxides, but admitted that this was mere conjecture. In 1808, acting on Lavoisier's idea, Humphry Davy became the first to obtain samples of the metals by electrolysis of their molten earths.

[edit] Discovery

The calcium compounds calcite and lime are known and used since prehistoric times the same is true for the beryllium compounds beryl and emerald. The mineral baryte is also known for a very long time because it is associated with hydrothermal deposits of sulfide ores.

The 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 farmer at Epsom in England. Strontium carbonate was discovered in minerals in the Scottish village of Strontian in 1790. The last element was the least abundant radioactive radium which was extracted from uraninite in 1898.[15][16][17]

All elements except beryllium were isolated by electrolysis of molten compounds. Magnesium, calcium and strontium were first produced by Humphry Davy in 1808, while beryllium was independently isolated Friedrich Wöhler and Antoine Bussy in 1828 by reacting berylium compounds with potassium. In 1910, radium was isolated as a pure metal by Curie and André-Louis Debierne also by electrolysis.[15][16][17]


[edit] Occurrence

Emerald is a naturally occurring compound of beryllium.

Calcium and magnesium are abundant in earths crust making up several important rock forming minerals like dolomite (dolostone) and calcite (limestone). The other non-radioactive members of the group only are present in smaller quantities but also form minerals like beryl, strontianite and baryte. Deposits of those minerals are mined to extract the elements for further use. Radium with a maximum half-life of 1601 years is only present in nature when it is resupplied by a decay chain from the radioactive decay from heavier elements.


[edit] Biological role and precautions

  • 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.
  • Magnesium and calcium are ubiquitous and essential to all known living organisms. They are involved in more than one role, with, for example, Mg/Ca 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 (e.g. bones).
  • Strontium and barium have a lower availability in the biosphere. Strontium plays an important role in marine aquatic life, especially hard corals. They use strontium to build their exoskeleton. These elements have some uses in medicine, for example "barium meals" in radiographic imaging, whilst strontium compounds are employed in some toothpastes.
  • Radium has a low availability and is highly radioactive, making it toxic to life.


[edit] Notes

  1. ^ In both the old IUPAC and the CAS systems for group numbering, this group is known as group IIA (pronounced as "group two A", as the "II" is a Roman numeral).[3]
  2. ^ 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.
  3. ^ The number given in parenthesesrefers to the measurement uncertainty. This uncertainty applies to the least significant figure(s) of the number prior to the parenthesized value (ie. counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, while 1.00794(72) stands for 1.00794±0.00072.[11]
  4. ^ The element does not have any stable nuclides, and a value in brackets indicates the mass number of the longest-lived isotope of the element.[12][13]

[edit] References

  1. ^ a b c d e f g h i j k l m n o Royal Society of Chemistry. "Visual Elements: Group 2–The Alkaline Earth Metals". Visual Elements. Royal Society of Chemistry. http://www.rsc.org/chemsoc/visualelements/PAGES/data/intro_groupii_data.html. Retrieved 13 January 2012. 
  2. ^ Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements (2nd ed.). Butterworth–Heinemann. ISBN 0080379419. 
  3. ^ Fluck, E. (1988). "New Notations in the Periodic Table". Pure Appl. Chem. (IUPAC) 60 (3): 431–436. DOI:10.1351/pac198860030431. http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf. Retrieved 24 March 2012. 
  4. ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. pp. 51. Electronic version..
  5. ^ "Periodic Table: Atomic Properties of the Elements". nist.gov. National Institute of Standards and Technology. September 2010. http://www.nist.gov/pml/data/upload/periodic_table_composite_2010_nobleed.pdf. Retrieved 17 February 2012. 
  6. ^ a b c d e Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press. 
  7. ^ "Abundance in Earth's Crust". WebElements.com. http://www.webelements.com/webelements/properties/text/image-flash/abund-crust.html. Retrieved 14 April 2007. 
  8. ^ Gäggeler, Heinz W. (5–7 November 2007). "Gas Phase Chemistry of Superheavy Elements". Lecture Course Texas A&M. http://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf. Retrieved 26 February 2012. 
  9. ^ a b Hoffman, Darleane C.; Lee, Diana M.; Pershina, Valeria (2006). "Transactinides and the future elements". In Morss; Edelstein, Norman M.; Fuger, Jean. The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer Science+Business Media. ISBN 1-4020-3555-1. 
  10. ^ a b Seaborg, G. T. (ca. 2006). "transuranium element (chemical element)". Encyclopædia Britannica. http://www.britannica.com/EBchecked/topic/603220/transuranium-element. Retrieved 16 March 2010. 
  11. ^ "Standard Uncertainty and Relative Standard Uncertainty". CODATA reference. National Institute of Standards and Technology. http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html. Retrieved 26 September 2011. 
  12. ^ a b Wieser, Michael E.; Berglund, Michael (2009). "Atomic weights of the elements 2007 (IUPAC Technical Report)". Pure Appl. Chem. (IUPAC) 81 (11): 2131–2156. DOI:10.1351/PAC-REP-09-08-03. http://iupac.org/publications/pac/pdf/2009/pdf/8111x2131.pdf. Retrieved 7 February 2012. 
  13. ^ a b Wieser, Michael E.; Coplen, Tyler B. (2011). "Atomic weights of the elements 2009 (IUPAC Technical Report)". Pure Appl. Chem. (IUPAC) 83 (2): 359–396. DOI:10.1351/PAC-REP-10-09-14. http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf. Retrieved 11 February 2012. 
  14. ^ Slater, J. C. (1964). "Atomic Radii in Crystals". Journal of Chemical Physics 41 (10): 3199–3205. Bibcode 1964JChPh..41.3199S. DOI:10.1063/1.1725697. 
  15. ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. X. The alkaline earth metals and magnesium and cadmium". Journal of Chemical Education 9 (6): 1046. Bibcode 1932JChEd...9.1046W. DOI:10.1021/ed009p1046. 
  16. ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. XII. Other elements isolated with the aid of potassium and sodium: Beryllium, boron, silicon, and aluminum". Journal of Chemical Education 9 (8): 1386. Bibcode 1932JChEd...9.1386W. DOI:10.1021/ed009p1386. 
  17. ^ a b Weeks, Mary Elvira (1933). "The discovery of the elements. XIX. The radioactive elements". Journal of Chemical Education 10 (2): 79. Bibcode 1933JChEd..10...79W. DOI:10.1021/ed010p79. 
Layouts
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Periodic table
H   He
Li Be   B C N O F Ne
Na Mg   Al Si P S Cl Ar
K Ca   Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
Rb Sr   Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe
Cs Ba La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn
Fr Ra Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo
Alkali metals Alkaline earth metals Lanthanides Actinides Transition metals Post-​transition metals Metalloids Other nonmetals Halogens Noble gases Unknown chemical properties
Large version
Alkaline earth metals
   

Beryllium
Be
Atomic Number: 4
Atomic Weight: 9.012182
Melting Point: 1560.15 K
Boiling Point: 2742 K
Specific mass: 1.85 g/cm3
Electronegativity: 1.57

Magnesium
Mg
Atomic Number: 12
Atomic Weight: 24.3050
Melting Point: 923.15 K
Boiling Point: 1363 K
Specific mass: 1.738 g/cm3
Electronegativity: 1.31

Calcium
Ca
Atomic Number: 20
Atomic Weight: 40.078
Melting Point: 1112.15 K
Boiling Point: 1757 K
Specific mass: 1.54 g/cm3
Electronegativity: 1

Strontium
Sr
Atomic Number: 38
Atomic Weight: 87.62
Melting Point: 1042.15 K
Boiling Point: 1655 K
Specific mass: 2.64 g/cm3
Electronegativity: 0.95

Barium
Ba
Atomic Number: 56
Atomic Weight: 137.327
Melting Point: 1002.15 K
Boiling Point: 2170 K
Specific mass: 3.594 g/cm3
Electronegativity: 0.89

Radium
Ra
Atomic Number: 88
Atomic Weight: [226]
Melting Point: 973.15 K
Boiling Point: 2010 K
Specific mass: 5.5 g/cm3
Electronegativity: 0.9
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