World Library  
Flag as Inappropriate
Email this Article

Alkali metal

Article Id: WHEBN0000000666
Reproduction Date:

Title: Alkali metal  
Author: World Heritage Encyclopedia
Language: English
Subject: Periodic table, Caesium, Lithium, Metal, Potassium
Collection: Alkali Metals, Articles Containing Video Clips, Groups in the Periodic Table, Periodic Table
Publisher: World Heritage Encyclopedia

Alkali metal

A sample of petalite
Petalite, the lithium mineral from which lithium was first isolated

Sodium compounds have been known since ancient times; salt (

  • "Group 1: The Alkali Metals". Visual Elements.  
  • Atomic and Physical Properties of the Group 1 Elements An in-depth look at alkali metals
  • Alkali Metal Bangs Filmed reactions of five-gram samples of the alkali metals with water

External links

  • Bauer, Brent A., Robert Houlihan, Michael J. Ackerman, Katya Johnson, and Himeshkumar Vyas (2006). "Acquired Long QT Syndrome Secondary to Cesium Chloride Supplement".  
  • Campbell, Linda M., Aaron T. Fisk, Xianowa Wang, Gunter Kock, and Derek C. Muir (2005). "Evidence for Biomagnification of Rubidium in Freshwater and Marine Food Webs". Canadian Journal of Fisheries and Aquatic Sciences 62 (5): 1161–1167.  
  • Chang, Cheng-Hung, and Tian Y. Tsong (2005). "AIP Conference Proceedings: 18th International Conference on Noise and Fluctuations" 780.  
  • Erermis, Serpil, Muge Tamar, Hatice Karasoy, Tezan Bildik, Eyup S. Ercan, and Ahmet Gockay (1997). "Double-Blind Randomised Trial of Modest Salt Restriction in Older People".  
  • Joffe, Russell T., Stephen T. Sokolov and Anthony J. Levitt (2006). "Lithium and Triiodothyronine Augmentation of Antidepressants". Canadian Journal of Psychiatry 51 (12): 791–3.  
  • Krachler, M, and E Rossipal (1999). "Trace Elements Transfer From Mother to the Newborn – Investigations on Triplets of Colostrum, Maternal and Umbilical Sera". European Journal of Clinical Nutrition 53 (6): 486–494.  
  • Stein, Benjamin P., Stephen G. Benka, Phillip F. Schewe, and Bertram Schwarzhild (1996). "Physics Update".  

Further reading

  1. ^ RSC–IUPAC. ISBN 0-85404-438-8. pp. 248–49. Electronic version..
  2. ^ Coghill, Anne M.; Garson, Lorrin R., eds. (2006). The ACS Style Guide: Effective Communication of Scientific Information (3rd ed.). Washington, D.C.: American Chemical Society. p. 127.  .
  3. ^ Coplen, T. B.; Peiser, H. S. (1998). "History of the recommended atomic-weight values from 1882 to 1997: a comparison of differences from current values to the estimated uncertainties of earlier values". Pure Appl. Chem. 70 (1): 237–257.  .
  4. ^ a b c RSC–IUPAC. ISBN 0-85404-438-8. pp. 51. Electronic version..
  5. ^ Leach, Mark R. (1999–2012). "The Internet Database of Periodic Tables". Retrieved 20 May 2012. 
  6. ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad  
  7. ^ "Periodic Table: Atomic Properties of the Elements".  
  8. ^ a b c Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press. 
  9. ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am an ao ap aq ar as at au av aw ax  
  10. ^ a b c The OpenLearn team (2012). "Alkali metals". OpenLearn. The Open University. Retrieved 9 July 2012. 
  11. ^ Fluck, E. (1988). "New Notations in the Periodic Table".  
  12. ^ "IUPAC Periodic Table of the Elements".  
  13. ^ a b c "International Union of Pure and Applied Chemistry > Periodic Table of the Elements". IUPAC. Retrieved 1 May 2011. 
  14. ^ a b Folden, Cody (31 January 2009). "The Heaviest Elements in the Universe". Saturday Morning Physics at Texas A&M. Retrieved 9 March 2012. 
  15. ^ a b c d e f g Vinson, Greg (2008). "Hydrogen is a Halogen". Retrieved 14 January 2012. 
  16. ^ a b c Gray, Theodore. "Alkali Metal Bangs".  
  17. ^ a b "Abundance in Earth's Crust". Retrieved 14 April 2007. 
  18. ^ a b Adloff, Jean-Pierre; Kaufman, George B. (2005-09-25). Francium (Atomic Number 87), the Last Discovered Natural Element. The Chemical Educator 10 (5). Retrieved on 26 March 2007.
  19. ^ a b c Gäggeler, Heinz W. (5–7 November 2007). "Gas Phase Chemistry of Superheavy Elements". Lecture Course Texas A&M. Retrieved 26 February 2012. 
  20. ^ a b c d e f g h i j k l m n o p q r s t u 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:  
  21. ^ a b "Cesium Atoms at Work". Time Service Department—U.S. Naval Observatory—Department of the Navy. Retrieved 20 December 2009. 
  22. ^ a b c d e Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Archived from the original on 22 November 2009. Retrieved 27 December 2009. 
  23. ^ a b
  24. ^ a b Lindsey, Jack L (1997). Applied illumination engineering. p. 112.  
  25. ^ a b Kane, Raymond; Sell, Heinz (2001). Revolution in lamps: A chronicle of 50 years of progress. p. 241.  
  26. ^ a b c d Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | biological information". WebElements. Retrieved 13 January 2012. 
  27. ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Sodium | biological information". WebElements. Retrieved 13 January 2012. 
  28. ^ a b c d e f Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | biological information". Webelements. Retrieved 15 February 2011. 
  29. ^ a b c d Winter, Mark. "WebElements Periodic Table of the Elements | Rubidium | biological information". Webelements. Retrieved 15 February 2011. 
  30. ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Caesium | biological information". WebElements. Retrieved 13 January 2012. 
  31. ^ a b "Francium – Element information, properties and uses | Periodic Table". Visual Elements Periodic Table.  
  32. ^ a b c d e f g h i j k l m n o p q Averill, Bruce A.; Eldredge, Patricia (2007). "21.3: The Alkali Metals". Chemistry: Principles, Patterns, and Applications with Student Access Kit for Mastering General Chemistry (1st ed.). Prentice Hall.  
  33. ^ "Standard Uncertainty and Relative Standard Uncertainty".  
  34. ^ a b c d Wieser, Michael E.; Berglund, Michael (2009). "Atomic weights of the elements 2007 (IUPAC Technical Report)".  
  35. ^ a b c d e Wieser, Michael E.; Coplen, Tyler B. (2011). "Atomic weights of the elements 2009 (IUPAC Technical Report)".  
  36. ^ Klehr, Wolfram (21 May 2007). "Francium". Archived from the original on 9 May 2008. Retrieved 25 April 2012. 
  37. ^ a b c d Clark, Jim (2005). "Atomic and Physical Properties of the Group 1 Elements". chemguide. Retrieved 30 January 2012. 
  38. ^  
  39. ^ a b Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements: A Reference Guide. Westport, Conn.: Greenwood Press.  
  40. ^ J. L. Dye, J. M. Ceraso, Mei Lok Tak, B. L. Barnett, F. J. Tehan (1974). "Crystalline salt of the sodium anion (Na)".  
  41. ^ F. J. Tehan, B. L. Barnett, J. L. Dye (1974). "Alkali anions. Preparation and crystal structure of a compound which contains the cryptated sodium cation and the sodium anion".  
  42. ^ J. L. Dye (1979). "Compounds of Alkali Metal Anions".  
  43. ^ M. Y. Redko, R. H. Huang, J. E. Jackson, J. F. Harrison, J. L. Dye (2003). "Barium azacryptand sodide, the first alkalide with an alkaline Earth cation, also contains a novel dimer, (Na2)2−".  
  44. ^ a b c M. Y. Redko, M. Vlassa, J. E. Jackson, A. W. Misiolek, R. H. Huang RH, J. L. Dye (2002). ""Inverse sodium hydride": a crystalline salt that contains H+ and Na".  
  45. ^ a b Agnieszka Sawicka, Piotr Skurski, and Jack Simons (2003). "Inverse Sodium Hydride: A Theoretical Study". J. Am. Chem. Soc. 125 (13): 3954–3958.  
  46. ^ Burgess, John (1978). Metal Ions in Solution. Chichester: Ellis Horwood. p. 20.  
  47. ^ a b Richens, David. T. (1997). The Chemistry of Aqua Ions. Wiley.  
  48. ^ Persson, Ingmar (2010). "Hydrated metal ions in aqueous solution: How regular are their structures?". Pure Appl. Chem. 82 (10): 1901–1917.  
  49. ^ a b c d e Clark, Jim (2005). "Reaction of the Group 1 Elements with Oxygen and Chlorine". chemguide. Retrieved 27 June 2012. 
  50. ^ Shriver, Duward; Atkins, Peter (2006). Inorganic Chemistry. W. H. Freeman. p. 259.  
  51. ^ Andreev, S.V.; Letokhov, V.S.; Mishin, V.I., (1987). "Laser resonance photoionization spectroscopy of Rydberg levels in Fr".  
  52. ^ a b c d e Thayer, John S. (2010). Chemistry of heavier main group elements. pp. 81, 84.  
  53. ^ a b Various authors (2002). Lide, David R., ed. Handbook of Chemistry & Physics (88th ed.). CRC.  
  54. ^ "Universal Nuclide Chart". Nucleonica. Institute for Transuranium Elements. 2007–2012. Retrieved 2011-04-17. 
  55. ^ a b c Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 4 October 2012. 
  56. ^ Patton, I. Jocelyn; Waldbauer, L. J. (1926). "The Radioactivity of the Alkali Metals". Chemical Reviews 3: 81.  
  57. ^ McLennan, J. C.; Kennedy, W. T. (1908). "On the radioactivity of potassium and other alkali metals". Philosophical Magazine. 6 16 (93): 377–395.  
  58. ^ "Potassium-40". Human Health Fact Sheet.  
  59. ^ a b Fontani, Marco (10 September 2005). "International Conference on the History of Chemistry". Lisbon. pp. 1–8. Archived from the original on 24 February 2006. Retrieved 8 April 2007. 
  60. ^ a b Van der Krogt, Peter (10 January 2006). "Francium". Elementymology & Elements Multidict. Retrieved 8 April 2007. 
  61. ^ National Institute of Standards and Technology. "Radionuclide Half-Life Measurements". Retrieved 2011-11-07. 
  62. ^
  63. ^ a b The Radiological Accident in Goiânia.  
  64. ^ Clark, Jim (2005). "Reaction of the Group 1 Elements with Water". chemguide. Retrieved 18 June 2012. 
  65. ^ Catherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 14: The group 14 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 386.  
  66. ^ NIST Ionizing Radiation Division 2001 - Technical Highlights
  67. ^ N. Emery et al. (2008). "Review: Synthesis and superconducting properties of CaC6". Sci. Technol. Adv. Mater. 9 (4): 044102.  
  68. ^ S.M. Kauzlarich,(1994), Zintl Compounds, Encyclopedia of Inorganic Chemistry, John Wiley & sons, ISBN 0-471-93620-0
  69. ^ Hoch, Constantin; Wendorff, Marco; Röhr, Caroline (2002). "Tetrapotassium nonastannide, K4Sn9". Acta Crystallographica Section C Crystal Structure Communications 58 (4): i45.  
  70. ^ Duncan H. Gregory, Paul M. O'Meara, Alexandra G. Gordon, Jason P. Hodges, Simine Short, and James D. Jorgensen (2002). "Structure of Lithium Nitride and Transition-Metal-Doped Derivatives, Li3−xyMxN (M = Ni, Cu): A Powder Neutron Diffraction Study". Chem. Mater. 14 (5): 2063–2070.  
  71. ^ Fischer, D., Jansen, M. (2002). "Synthesis and structure of Na3N". Angew Chem 41 (10): 1755.  
  72. ^ Fischer, D.; Cancarevic, Z.; Schön, J. C.; Jansen, M. Z. (2004). "Synthesis and structure of K3N". Z. anorg allgem Chemie 630 (1): 156.  . 'Elusive Binary Compound Prepared' Chemical & Engineering News 80 No. 20 (20 May 2002)
  73. ^ H.G. Von Schnering, W. Hönle Phosphides - Solid-state Chemistry Encyclopedia of Inorganic Chemistry Ed. R. Bruce King (1994) John Wiley & Sons ISBN 0-471-93620-0
  74. ^ Kahlenberg, Louis (2008). Outlines of Chemistry – A Textbook for College Students. READ BOOKS. pp. 324–325.  
  75. ^ "Welcome to Arthur Mar's Research Group". University of Alberta. University of Alberta. 1999–2013. Retrieved 24 June 2013. 
  76. ^ Lindsay, D. M.; Garland, D. A. (1987). "ESR spectra of matrix-isolated lithium superoxide". The Journal of Physical Chemistry 91 (24): 6158–61.  
  77. ^ Vol'nov, I. I.; Matveev, V. V. (1963). "Synthesis of cesium ozonide through cesium superoxide". Bulletin of the Academy of Sciences, USSR Division of Chemical Science 12 (6): 1040–1043.  
  78. ^ Tokareva, S. A. (1971). "Alkali and Alkaline Earth Metal Ozonides". Russian Chemical Reviews 40 (2): 165–174.  
  79. ^ Simon, A. (1997). "Group 1 and 2 Suboxides and Subnitrides — Metals with Atomic Size Holes and Tunnels". Coordination Chemistry Reviews 163: 253–270.  
  80. ^ Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). "The Crystal Structure of Tricesium Monoxide". Journal of Physical Chemistry 60 (3): 345–347.  
  81. ^ Okamoto, H. (2009). "Cs-O (Cesium-Oxygen)". Journal of Phase Equilibria and Diffusion 31: 86.  
  82. ^ Band, A.; Albu-Yaron, A.; Livneh, T.; Cohen, H.; Feldman, Y.; Shimon, L.; Popovitz-Biro, R.; Lyahovitskaya, V.; Tenne, R. (2004). "Characterization of Oxides of Cesium". The Journal of Physical Chemistry B 108 (33): 12360–12367.  
  83. ^ Brauer, G. (1947). "Untersuchungen ber das System Cäsium-Sauerstoff". Zeitschrift für anorganische Chemie 255: 101.  
  84. ^ Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Retrieved 27 December 2009. 
  85. ^
  86. ^ Moyer, Harvey V. (1956). "Chemical Properties of Polonium". In Moyer, Harvey V. Polonium. Oak Ridge, Tenn.: United States Atomic Energy Commission. pp. 33–96.  .
  87. ^ Bagnall, K. W. (1962). "The Chemistry of Polonium". Adv. Inorg. Chem. Radiochem. Advances in Inorganic Chemistry and Radiochemistry 4: 197–229.  .
  88. ^ Alberto, R.; Ortner, K.; Wheatley, N.; Schibli, R.; Schubiger, A. P. (2001). "Synthesis and properties of boranocarbonate: a convenient in situ CO source for the aqueous preparation of [99mTc(OH2)3(CO)3]+".  
  89. ^ Cotton, F.A.; Wilkinson, G. (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc.  
  90. ^ Brown, T. L.; Rogers, M. T. (1957). "The Preparation and Properties of Crystalline Lithium Alkyls". Journal of the American Chemical Society 79 (8): 1859–1861.  
  91. ^ Manfred Schlosser (1988). "Superbases for organic synthesis". Pure and Appl. Chem. 60 (11): 1627–1634.  
  92. ^ Clegg, William; Conway, Ben; Kennedy, Alan R.; Klett, Jan; Mulvey, Robert E.; Russo, Luca (2011). "Synthesis and Structures of \(Trimethylsilyl)methyl]sodium and -potassium with Bi- and Tridentate N-Donor Ligands". European Journal of Inorganic Chemistry 2011 (5): 721.  
  93. ^ a b c d Pyykkö, Pekka (2011). "A suggested periodic table up to Z ≤ 172, based on Dirac–Fock calculations on atoms and ions". Physical Chemistry Chemical Physics 13 (1): 161–8.  
  94. ^ a b c Seaborg, G. T. (c. 2006). "transuranium element (chemical element)". Encyclopædia Britannica. Retrieved 16 March 2010. 
  95. ^ Fricke, Burkhard (1975). "Superheavy elements: a prediction of their chemical and physical properties". Recent Impact of Physics on Inorganic Chemistry 21: 89–144.  
  96. ^ Kratz, J. V. (5 September 2011). "The Impact of Superheavy Elements on the Chemical and Physical Sciences". 4th International Conference on the Chemistry and Physics of the Transactinide Elements. Retrieved 27 August 2013. 
  97. ^
  98. ^
  99. ^ Emsley, J. (1989). The Elements. Oxford: Clarendon Press. pp. 22–23. 
  100. ^ Chemical Bonding, Mark J. Winter, Oxford University Press, 1994, ISBN 0-19-855694-2
  101. ^ a b c d e f Cronyn, Marshall W. (August 2003). "The Proper Place for Hydrogen in the Periodic Table". Journal of Chemical Education 80 (8): 947–951.  
  102. ^ J.E. Huheey, E.A. Keiter, and R.L. Keiter in Inorganic Chemistry : Principles of Structure and Reactivity, 4th edition, HarperCollins, New York, USA, 1993.
  103. ^ A.M. James and M.P. Lord in Macmillan's Chemical and Physical Data, Macmillan, London, UK, 1992.
  104. ^ Wigner, E.; Huntington, H.B. (1935). "On the possibility of a metallic modification of hydrogen".  
  105. ^ Nellis, W. J.; Weir, S. T.; Mitchell, A. C. (1999). "Metallization of fluid hydrogen at 140 GPa (1.4 Mbar) by shock compression". Shock Waves 9: 301–305.  
  106. ^ Cousins, David M.; Davidson, Matthew G.; García-Vivó, Daniel (2013). "Unprecedented participation of a four-coordinate hydrogen atom in the cubane core of lithium and sodium phenolates". Chem. Commun. 49: 11809–11811.  
  107. ^ Mark R. Leach. "2002 Inorganic Chemist's Periodic Table". Retrieved 16 October 2012. 
  108. ^ Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press,  
  109. ^ a b c Stevenson, D. J. (20 November 1975). "Does metallic ammonium exist?".  
  110. ^ a b Bernal, M. J. M.; Massey, H. S. W. (3 February 1954). "Metallic Ammonium".  
  111. ^ "Solubility Rules!". Retrieved 4 January 2014. 
  112. ^ R. D. Shannon (1976). "Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides". Acta Cryst A32 (5): 751–767.  
  113. ^ a b  
  114. ^ a b c d Leach, Mark R. (1999–2012). "The Internet Database of Periodic Tables". Retrieved 6 April 2012. 
  115. ^ Anja-Verena Mudring (2007). "Thallium Halides - New Aspects of the Stereochemical Activity of Electron Lone Pairs of Heavier Main-Group Elements".  
  116. ^ Russell AM & Lee KL 2005, Structure-property relations in nonferrous metals, p. 302. Wiley-Interscience, New York, ISBN 047164952X
  117. ^ Deming HG 1940, Fundamental Chemistry, John Wiley & Sons, New York, pp. 705–7
  118. ^ a b c d e Jensen, William B. (2003). "The Place of Zinc, Cadmium, and Mercury in the Periodic Table". Journal of Chemical Education ( 
  119. ^ Marggraf, Andreas Siegmund (1761). Chymische Schriften. p. 167. 
  120. ^ du Monceau, H. L. D. "Sur la Base de Sel Marine". Memoires de l'Academie royale des Sciences (in French): 65–68. 
  121. ^ a b  
  122. ^ a b Siegfried, R. (1963). "The Discovery of Potassium and Sodium, and the Problem of the Chemical Elements". Isis 54 (2): 247–258.  
  123. ^ Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim.  
  124. ^ a b Davy, Humphry (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society of London 98: 1–44.  
  125. ^ Shaposhnik, V. A. (2007). "History of the discovery of potassium and sodium (on the 200th anniversary of the discovery of potassium and sodium)". Journal of Analytical Chemistry 62 (11): 1100–1102.  
  126. ^ Ralph, Jolyon; Chau, Ida (24 August 2011). "Petalite: Petalite mineral information and data". Retrieved 27 November 2011. 
  127. ^ a b Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | historical information". Retrieved 27 November 2011. 
  128. ^ Weeks, Mary (2003). Discovery of the Elements. Whitefish, Montana, United States: Kessinger Publishing. p. 124.  
  129. ^ "Johan Arfwedson". Archived from the original on 5 June 2008. Retrieved 10 August 2009. 
  130. ^ a b van der Krogt, Peter. "Lithium". Elementymology & Elements Multidict. Retrieved 5 October 2010. 
  131. ^ Clark, Jim (2005). "Compounds of the Group 1 Elements". chemguide. Retrieved 10 August 2009. 
  132. ^ Kaner, Richard (2003). "C&EN: It's Elemental: The Periodic Table – Cesium". American Chemical Society. Retrieved 25 February 2010. 
  133. ^  
  134. ^  
  135. ^ Oxford English Dictionary, 2nd Edition
  136. ^ Newlands, John A. R. (20 August 1864). "On Relations Among the Equivalents". Chemical News 10: 94–95. Archived from the original on January 1, 2011. Retrieved November 25, 2013. 
  137. ^ Newlands, John A. R. (18 August 1865). "On the Law of Octaves". Chemical News 12: 83. Archived from the original on January 1, 2011. Retrieved November 25, 2013. 
  138. ^ Mendelejew, Dimitri (1869). "Über die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente". Zeitschrift für Chemie (in German): 405–406. 
  139. ^ Fluck, E. (1988). "New Notations in the Periodic Table".  
  140. ^ "Alabamine & Virginium". TIME. 15 February 1932. Retrieved 1 April 2007. 
  141. ^ MacPherson, H. G. (1934). "An Investigation of the Magneto-Optic Method of Chemical Analysis". Physical Review (American Physical Society) 47 (4): 310–315.  
  142. ^ "Francium".  
  143. ^ van der Krogt, Peter. "Ununennium". Elementymology & Elements Multidict. Retrieved 14 February 2011. 
  144. ^ Schadel, M.; Brüchle, W.; Brügger, M.; Gäggeler, H.; Moody, K.; Schardt, D.; Sümmerer, K.; Hulet, E.; Dougan, A. et al. (1986). "Heavy isotope production by multinucleon transfer reactions with 254Es". Journal of the Less Common Metals 122: 411.  
  145. ^ "Modern alchemy: Turning a line".  
  146. ^ Emsley, John (2011). Nature's Building Blocks: An A-Z Guide to the Elements (New ed.). New York, NY: Oxford University Press. p. 593.  
  147. ^ a b Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591 (2): 1220–1247.  
  148. ^ Oddo, Giuseppe (1914). "Die Molekularstruktur der radioaktiven Atome". Zeitschrift für anorganische Chemie 87: 253.  
  149. ^ Harkins, William D. (1917). "The Evolution of the Elements and the Stability of Complex Atoms. I. A New Periodic System Which Shows a Relation Between the Abundance of the Elements and the Structure of the Nuclei of Atoms". Journal of the American Chemical Society 39 (5): 856.  
  150. ^ North, John (2008). Cosmos an illustrated history of astronomy and cosmology (Rev. and updated ed.). Univ. of Chicago Press. p. 602.  
  151. ^ Morgan, J. W.; Anders, E. (1980). "Chemical composition of Earth, Venus, and Mercury". Proceedings of the National Academy of Sciences 77 (12): 6973–6977.  
  152. ^ Albarède, Francis (2003). Geochemistry: an introduction. Cambridge University Press.  
  153. ^ "List of Periodic Table Elements Sorted by Abundance in Earth's crust". Israel Science and Technology Homepage. Retrieved 15 April 2007. 
  154. ^ a b c Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press.  
  155. ^ "Lithium Occurrence". Institute of Ocean Energy, Saga University, Japan. Retrieved 13 March 2009. 
  156. ^ "Some Facts about Lithium". ENC Labs. Retrieved 15 October 2010. 
  157. ^ Schwochau, Klaus (1984). "Extraction of metals from sea water". Topics in Current Chemistry. Topics in Current Chemistry. 124/1984: 91–133.  
  158. ^ Wise, M. A. (1995). "Trace element chemistry of lithium-rich micas from rare-element granitic pegmatites". Mineralogy and Petrology 55 (13): 203–215.  
  159. ^ a b CRC Handbook of Chemistry and Physics 4. CRC. 2006. p. 12.  
  160. ^ a b c Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 151–153.  
  161. ^ a b Gagnon, Steve. "Francium". Jefferson Science Associates, LLC. Archived from the original on 31 March 2007. Retrieved 1 April 2007. 
  162. ^ a b Winter, Mark. "Geological information". Francium. The University of Sheffield. Retrieved 26 March 2007. 
  163. ^ "It's Elemental — The Periodic Table of Elements". Jefferson Lab. Archived from the original on 29 April 2007. Retrieved 14 April 2007. 
  164. ^ Ober, Joyce A. "Lithium" (PDF).  
  165. ^ Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | Essential information". Webelements. Retrieved 27 November 2011. 
  166. ^ Lemke, Charles H.; Markant, Vernon H. (2001). "Kirk-Othmer Encyclopedia of Chemical Technology".  
  167. ^ Pauling, Linus. General Chemistry (1970 ed.). Dover Publications. 
  168. ^ "Los Alamos National Laboratory – Sodium". Retrieved 8 June 2007. 
  169. ^ Merck Index, 9th ed., monograph 8325
  170. ^ "Cesium and Rubidium Hit Market". Chemical & Engineering News 37 (22): 50–56. 1959.  
  171. ^ a b c d Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2003). "Mineral Commodity Profile: Rubidium" (PDF). United States Geological Survey. Retrieved 4 December 2010. 
  172. ^ bulletin 585. United States. Bureau of Mines. 1995. 
  173. ^ Burt, R. O. (1993). "Caesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology 5 (4th ed.). New York: John Wiley & Sons, Inc. pp. 749–764.  
  174. ^ Stancari, G.; Veronesi, S.; Corradi, L.; Atutov, S. N.; Calabrese, R.; Dainelli, A.; Mariotti, E.; Moi, L.; Sanguinetti, S.; Tomassetti, L. (2006). "Production of Radioactive Beams of Francium". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment 557 (2): 390–396.  
  175. ^ a b c Luis A. Orozco (2003). "Francium". Chemical and Engineering News. 
  176. ^ USGS (2011). "Lithium" (PDF). Retrieved 4 December 2011. 
  177. ^ Stampers, National Association of Drop Forgers and (1957). Metal treatment and drop forging. 
  178. ^ Harris, Jay C (1949). Metal cleaning bibliographical abstracts. p. 76. 
  179. ^ Cordel, Oskar (1868). Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung ... (in German). L. Schnock. Retrieved 29 May 2011. 
  180. ^ Toedt, John; Koza, Darrell; Cleef-Toedt, Kathleen Van (2005). "Personal Cleansing Products: Bar Soap". Chemical composition of everyday products. Greenwood Publishing Group.  
  181. ^ Schultz, H. et al. (2006). "Potassium compounds". Ullmann's Encyclopedia of Industrial Chemistry A22. p. 95.  
  182. ^ Koch, E.-C. (2002). "Special Materials in Pyrotechnics, Part II: Application of Caesium and Rubidium Compounds in Pyrotechnics". Journal Pyrotechnics 15: 9–24. 
  183. ^ Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. McGraw-Hill. pp. 201–203.  
  184. ^ Winter, Mark. "Uses". Francium. The University of Sheffield. Archived from the original on 31 March 2007. Retrieved 25 March 2007. 
  185. ^ Gomez, E; Orozco, L A; Sprouse, G D (7 November 2005). "Spectroscopy with trapped francium: advances and perspectives for weak interaction studies". Rep. Prog. Phys. 69 (1): 79–118.  
  186. ^ Peterson, I (11 May 1996). "Creating, cooling, trapping francium atoms". Science News 149 (19): 294.  
  187. ^  
  188. ^ Howland, Robert H. (September 2007). "Lithium: Underappreciated and Underused?". Psychiatric Annals 37 (9). Retrieved 6 November 2012. 
  189. ^ Zarse, Kim; Terao, Takeshi; Tian, Jing; Iwata, Noboru; Ishii, Nobuyoshi; Ristow, Michael (August 2011). "Low-dose lithium uptake promotes longevity in humans and metazoans". European Journal of Nutrition (Springer) 50 (5): 387–9.  
  190. ^ "Sodium". Northewestern University. Retrieved 21 November 2011. 
  191. ^ "Sodium and Potassium Quick Health Facts". Retrieved 7 November 2011. 
  192. ^ "Dietary Reference Intakes: Water, Potassium, Sodium, Chloride, and Sulfate". Food and Nutrition Board,  
  193. ^ Geleijnse, J. M.; Kok, F. J.; Grobbee, D. E. (2004). "Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations". European Journal of Public Health 14 (3): 235–239.  
  194. ^ Lawes, C. M.; Vander Hoorn, S.; Rodgers, A.; International Society of Hypertension (2008). "Global burden of blood-pressure-related disease, 2001". Lancet 371 (9623): 1513–1518.  
  195. ^ a b Mikko Hellgren, Lars Sandberg, Olle Edholm (2006). "A comparison between two prokaryotic potassium channels (KirBac1.1 and KcsA) in a molecular dynamics (MD) simulation study". Biophys. Chem. 120 (1): 1–9.  
  196. ^ Relman, AS (1956). "The Physiological Behavior of Rubidium and Cesium in Relation to That of Potassium". The Yale journal of biology and medicine 29 (3): 248–62.  
  197. ^ a b Meltzer, HL (1991). "A pharmacokinetic analysis of long-term administration of rubidium chloride". Journal of clinical pharmacology 31 (2): 179–84.  
  198. ^ Pinsky, Carl; Bose, Ranjan; Taylor, J. R.; McKee, Jasper; Lapointe, Claude; Birchall, James (1981). "Cesium in mammals: Acute toxicity, organ changes and tissue accumulation". Journal of Environmental Science and Health, Part A 16 (5): 549– 567.  
  199. ^ Johnson, Garland T.; Lewis, Trent R.; Wagner, D. Wagner (1975). "Acute toxicity of cesium and rubidium compounds".  
  200. ^ Sartori H. E. (1984). "Cesium therapy in cancer patients". Pharmacol Biochem Behav 21 (Suppl 1): 11–13.  
  201. ^ Wood, Leonie. Cured' cancer patients died, court told"'". The Sydney Morning Herald. 20 November 2010. 
  202. ^ Winter, Mark. "WebElements Periodic Table of the Elements | Francium | biological information". WebElements. Retrieved 15 February 2011. 


  1. ^ The symbol Na for sodium is derived from its Latin name, natrium; this is still the name for the element in some languages, such as German and Russian. In early English texts, the symbol So for the English name sodium is sometimes seen; this is wholly obsolete.
  2. ^ The symbol K for potassium is derived from its Latin name, kalium; this is still the name for the element in some languages, such as German and Russian. In early English texts, the symbol Po for the English name potassium is sometimes seen; this is wholly obsolete, and presently would collide with the symbol for polonium (also Po).
  3. ^ Caesium is the spelling recommended by the International Union of Pure and Applied Chemistry (IUPAC).[1] The American Chemical Society (ACS) has used the spelling cesium since 1921,[2][3] following Webster’s Third New International Dictionary.
  4. ^ In both the old IUPAC and the CAS systems for group numbering, this group is known as group IA (pronounced as "group one A", as the "I" is a Roman numeral).[11]
  5. ^ The number given in parentheses refers to the measurement uncertainty. This uncertainty applies to the least significant figure(s) of the number prior to the parenthesized value (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.[34]
  6. ^ The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.[36]
  7. ^ 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.[35][36]
  8. ^ The quoted values are for the tetracoordinate ions, except for Rb+ and Cs+ where they are for the hexacoordinate ion.
  9. ^ Bunsen quotes Aulus Gellius Noctes Atticae II, 26 by Nigidius Figulus: Nostris autem veteribus "caesia" dicta est, quae a Graecis glaukopis, ut Nigidius ait, "de colore caeli quasi caelia.
  10. ^ In the 1869 version of Mendeleev's periodic table, copper and silver were placed in their own group, aligned with hydrogen and mercury, while gold was tentatively placed under uranium and the undiscovered eka-aluminium in the boron group.
  11. ^ Some synthetic elements, like technetium and plutonium, have later been found in nature.
  12. ^ The asterisk denotes an excited state.


Francium has no biological role[203] and is most likely to be toxic due to its extreme radioactivity, causing radiation poisoning,[32] but since the greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,[176] it is unlikely that most people will ever encounter francium.

Rubidium has no known biological role, but may help stimulate metabolism,[30][197][198] and, similarly to caesium,[30][31] replace potassium in the body causing potassium deficiency.[30][198] Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium, allowing the caesium to replace the potassium in the body, causing potassium deficiency.[31] Exposure to large amounts of caesium compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered in natural sources, caesium is not a major chemical environmental pollutant.[199] The median lethal dose (LD50) value for caesium chloride in mice is 2.3 g per kilogram, which is comparable to the LD50 values of potassium chloride and sodium chloride.[200] Caesium chloride has been promoted as an alternative cancer therapy,[201] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.[202] Radioisotopes of caesium require special precautions: the improper handling of caesium-137 gamma ray sources can lead to release of this radioisotope and radiation injuries. Perhaps the best-known case is the Goiânia accident of 1987, in which an improperly-disposed-of radiation therapy system from an abandoned clinic in the city of Goiânia, Brazil, was scavenged from a junkyard, and the glowing caesium salt sold to curious, uneducated buyers. This led to four deaths and serious injuries from radiation exposure. Together with caesium-134, iodine-131, and strontium-90, caesium-137 was among the isotopes distributed by the Chernobyl disaster which constitute the greatest risk to health.[64]

A wheel type radiotherapy device which has a long collimator to focus the radiation into a narrow beam. The caesium-137 chloride radioactive source is the blue square, and gamma rays are represented by the beam emerging from the aperture. This was the radiation source involved in the Goiânia accident, containing about 93 grams of caesium-137 chloride.

Potassium is the major cation (positive ion) inside animal cells,[27] while sodium is the major cation outside animal cells.[27][28] The concentration differences of these charged particles causes a difference in electric potential between the inside and outside of cells, known as the membrane potential. The balance between potassium and sodium is maintained by ion pumps in the cell membrane.[196] The cell membrane potential created by potassium and sodium ions allows the cell to generate an action potential—a "spike" of electrical discharge. The ability of cells to produce electrical discharge is critical for body functions such as neurotransmission, muscle contraction, and heart function.[196]

Sodium and potassium occur in all known biological systems, generally functioning as electrolytes inside and outside cells.[27][28] Sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day.[191] Sodium chloride (also known as common salt) is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods.[192] The DRI for sodium is 1.5 grams per day,[193] but most people in the United States consume more than 2.3 grams per day,[1] the minimum amount that promotes hypertension;[194] this in turn causes 7.6 million premature deaths worldwide.[195]

Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested.[29] Lithium carbonate is used as a mood stabiliser in psychiatry to treat bipolar disorder (manic-depression) in daily doses of about 0.5 to 2 grams, although there are side-effects.[29] Excessive ingestion of lithium causes drowsiness, slurred speech and vomiting, among other symptoms,[29] and poisons the central nervous system,[29] which is dangerous as the required dosage of lithium to treat bipolar disorder is only slightly lower than the toxic dosage.[29][188] Its biochemistry, the way it is handled by the human body and studies using rats and goats suggest that it is an essential trace element, although the natural biological function of lithium in humans has yet to be identified.[189][190]

Biological role and precautions

Francium has no commercial applications,[161][162][185] but because of francium's relatively simple atomic structure, among other things, it has been used in spectroscopy experiments, leading to more information regarding energy levels and the coupling constants between subatomic particles.[186] Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels, similar to those predicted by quantum theory.[187]

Rubidium and caesium are often used in atomic clocks.[22] Caesium atomic clocks are extraordinarily accurate; if a clock had been made at the time of the dinosaurs, it would be off by less than four seconds (after 80 million years).[23] For that reason, caesium atoms are used as the definition of the second.[24] Rubidium ions are often used in purple fireworks,[183] and caesium is often used in drilling fluids in the petroleum industry.[23][184]

FOCS 1, a caesium atomic clock in Switzerland
FOCS 1, a caesium atomic clock in Switzerland

Potassium compounds are often used as fertilisers[9]:73[180] as potassium is an important element for plant nutrition. Other potassium ions are often used to hold anions. Potassium hydroxide is a very strong base, and is used to control the pH of various substances.[181][182]

Pure sodium has many applications, including use in sodium-vapour lamps, which produce very efficient light compared to other types of lighting,[25][26] and can help smooth the surface of other metals.[178][179] Being a strong reducing agent, it is often used to reduce many other metals, such as titanium and zirconium, from their chlorides.[9]:74 Sodium compounds have many applications as well, the most well-known compound being table salt. Sodium is also used in soap as salts of fatty acids.

All of the discovered alkali metals excluding francium have many applications. Lithium is often used in batteries, and lithium oxide can help process silica. Lithium can also be used to make lubricating greases, air treatment, and aluminium production.[177]


Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with calcium chloride typically added to lower the melting point of the mixture. The heavier alkali metals, however, is more typically isolated in a different way, where a reducing agent (typically sodium for potassium and magnesium or calcium for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes fractional distillation for purification.[33]

From their silicate ores, all the alkali metals may be obtained the same way: sulfuric acid is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the hydroxide. The remaining insoluble alkali metal carbonate is then precipitated selectively; the salt is then dissolved in hydrochloric acid. The result is then left to evaporate and the alkali metal can then be isolated through electrolysis.[33]

Francium-223, the only naturally occurring isotope of francium,[35][36] is produced naturally as the product of the alpha decay of actinium-227. Francium can be found in trace amounts in uranium and thorium minerals;[160] it has been calculated that at most there are 30 g of francium in the earth's crust at any given time.[163] As a result of its extreme rarity in nature, most francium is synthesized in the nuclear reaction 197Au + 18O210Fr + 5 n, yielding francium-209, francium-210, and francium-211.[175] The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,[176] which were synthesized using the nuclear reaction given above.[176]

For several years in the 1950s and 1960s, a by-product of the potassium production called Alkarb was a main source for rubidium. Alkarb contained 21% rubidium while the rest was potassium and a small fraction of caesium.[171] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.[172] Today, a common method for separating rubidium from potassium and caesium is the fractional crystallization of a rubidium and caesium alum (Cs, Rb)Al(SO4)2·12H2O, which yields pure rubidium alum after approximately 30 different reactions.[172][173] The limited applications and the lack of a mineral rich in rubidium limits the production of rubidium compounds to 2 to 4 tonnes per year.[172] Caesium, however, is not produced from the above reaction. Instead, the mining of pollucite ore is the main method of obtaining pure caesium, extracted from the ore mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.[172][174] Both metals are produced as by-products of lithium production: after 1958, when interest in lithium's thermonuclear properties increased sharply, the production of rubidium and caesium also increased correspondingly.[9]:71

A shiny gray 5-centimeter piece of matter with a rough surface.
This sample of uraninite contains about 100,000 atoms (3.3×10−20 g) of francium-223 at any given time.[161]

Potassium occurs in many minerals, such as sylvite (potassium chloride).[6] It is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide,[166] found extensively in places such as Canada, Russia, Belarus, Germany, Israel, United States, and Jordan, in a method similar to how sodium was produced in the late 1800s and early 1900s.[167] It can also be produced from seawater. Sodium occurs mostly in seawater and dried seabed,[6] but is now produced through electrolysis of sodium chloride by lowering the melting point of the substance to below 700 °C through the use of a Downs cell.[168][169] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.[170]

Lithium salts have to be extracted from the water of mineral springs, brine pools, and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride.[165]

The production of pure alkali metals is difficult due to their extreme reactivity with commonly used substances, such as water. The alkali metals are so reactive that they cannot be displaced by other elements and must be isolated through high-energy methods such as electrolysis.[6][33]

Salt flats are rich in lithium, such as these in Salar del Hombre Muerto, Argentina (left) and Uyuni, Bolivia (right). The lithium-rich brine is concentrated by pumping it into solar evaporation ponds (visible in Argentina image).

Production and isolation

Francium-223, the only naturally occurring isotope of francium,[35][36] is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals.[160] In a given sample of uranium, there is estimated to be only one francium atom for every 1018 uranium atoms.[161][162] It has been calculated that there is at most 30 g of francium in the earth's crust at any time, due to its extremely short half-life of 22 minutes.[163][164]

Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite,[159] although none of these contain only rubidium and no other alkali metals.[9]:70 Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.[23]

Despite its chemical similarity, lithium typically does not occur together with sodium or potassium due to its smaller size.[9]:69 Due to its relatively low reactivity, it can be found in seawater in large amounts; it is estimated that seawater is approximately 0.14 to 0.25 parts per million (ppm)[156][157] or 25 micromolar.[158] Its diagonal relationship with magnesium often allows it to replace magnesium in ferromagnesium minerals, where its crustal concentration is about 18 ppm, comparable to that of gallium and niobium. Commercially, the most important lithium mineral is spodumene, which occurs in large deposits worldwide.[9]:69

Sodium and potassium are very abundant in earth, both being among the ten most common elements in Earth's crust;[17][154] sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall[155] and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.[155] Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite.[155] Many of these solid deposits occur as a result of ancient seas evaporating, which still occurs now in places such as Utah's Great Salt Lake and the Dead Sea.[9]:69 Despite their near-equal abundance in Earth's crust, sodium is far more common than potassium in the ocean, both because potassium's larger size makes its salts less soluble, and because potassium is bound by silicates in soil and what potassium leaches is absorbed far more readily by plant life than sodium.[9]:69

The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are lithophiles and therefore remain close to the Earth's surface because they combine readily with oxygen and so associate strongly with silica, forming relatively low-density minerals that do not sink down into the Earth's core. Potassium, rubidium and caesium are also incompatible elements due to their large ionic radii.[153]

The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the solar system. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98×1024 kg. It is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to mass segregation, the core region is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.[152]

Spodumene, an important lithium mineral

On Earth

The Oddo-Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability.[149][150][151] All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesized in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesized in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements.[148]

Estimated abundances of the chemical elements in the Solar system. Hydrogen and helium are most common, from the Big Bang. The next three elements (lithium, beryllium, and boron) are rare because they are poorly synthesized in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers, and (2) a general decrease in abundance, as elements become heavier. Iron is especially common because it represents the minimum energy nuclide that can be made by fusion of helium in supernovae.[148]

In the Solar System


It is highly unlikely[19] that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making sufficient amounts of 254Es, which is favoured for production of ultraheavy elements because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms,[145] to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and has only been produced in laboratories. However, given that ununennium is only the first period 8 element on the extended periodic table, it may well be discovered in the near future through other reactions; indeed, another attempt to synthesise ununennium by bombarding a berkelium target with titanium ions is under way at the GSI Helmholtz Centre for Heavy Ion Research in Darmstadt, Germany.[146] Currently, none of the period 8 elements have been discovered yet, and it is also possible, due to drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible.[95][147] No attempts at synthesis have been made for any heavier alkali metals, such as unhexpentium, due to their extremely high atomic number.[21]:1737–1739

+ 48
* → no atoms[note 12]

The next element below francium (eka-francium) is very likely to be ununennium (Uue), element 119,[21]:1729–1730 although this is not completely certain due to relativistic effects.[20] The synthesis of ununennium was first attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.[19][144]

There were at least four erroneous and incomplete discoveries[60][61][141][142] before Marguerite Perey of the Curie Institute in Paris, France discovered francium in 1939 by purifying a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227.[18] Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%.[143] It was the last element discovered in nature, rather than by synthesis.[note 11]

After 1869, Dmitri Mendeleev proposed his periodic table placing lithium at the top of a group with sodium, potassium, rubidium, caesium, and thallium.[139] Two years later, Mendeleev revised his table, placing hydrogen in group 1 above lithium, and also moving thallium to the boron group. In this 1871 version, copper, silver, and gold were placed twice, once as part of group IB, and once as part of a "group VIII" encompassing today's groups 8 to 11.[119][note 10] After the introduction of the 18-column table, the group IB elements were moved to their current position in the d-block, while alkali metals were left in group IA. Later the group's name was changed to group 1 in 1988.[140] The trivial name "alkali metals" comes from the fact that the hydroxides of the group 1 elements are all strong alkalis when dissolved in water.[6]

Mendeleev's periodic system proposed in 1871 showing hydrogen and the alkali metals as part of his group I

Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music.[137][138] His version put all the alkali metals then known (lithium to caesium), as well as copper, silver, and thallium (which show the +1 oxidation state characteristic of the alkali metals), together into a group. His table placed hydrogen with the halogens.[115]

Rubidium and caesium were the first elements to be discovered using the spectroscope, invented in 1859 by Robert Bunsen and Gustav Kirchhoff.[133] The next year, they discovered caesium in the mineral water from Bad Dürkheim, Germany. Their discovery of rubidium came the following year in Heidelberg, Germany, finding it in the mineral lepidolite.[134] The names of rubidium and caesium come from the most prominent lines in their emission spectra: a bright red line for rubidium (from the Latin word rubidus, meaning dark red or bright red), and a sky-blue line for caesium (derived from the Latin word caesius, meaning sky-blue).[135][note 9][136]

A sample of lepidolite
Lepidolite, the rubidium mineral from which rubidium was first isolated

Petalite (LiAlSi4O10) was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada in a mine on the island of Utö, Sweden.[127][128][129] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jacob Berzelius, detected the presence of a new element while analysing petalite ore.[130][131] This new element was noted by him to form compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline than the other alkali metals.[132] Berzelius gave the unknown material the name "lithion/lithina", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium, which was known partly for its high abundance in animal blood. He named the metal inside the material "lithium".[40][128][131] Lithium, sodium, and potassium were part of the discovery of periodicity, as they are among a series of triads of elements in the same group that were noted by Johann Wolfgang Döbereiner in 1850 as having similar properties.[115]

Johann Wolfgang Döbereiner was among the first to notice similarities between what are now known as the alkali metals.

Pure potassium was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented voltaic pile. Previous attempts at electrolysis of the aqueous salt were unsuccessful due to potassium's extreme reactivity.[9]:68 Potassium was the first metal that was isolated by electrolysis.[124] Later that same year, Davy reported extraction of sodium from the similar substance caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.[122][123][125][126] Later that year, the first pieces of pure molten sodium metal were similarly prepared by Humphry Davy through the electrolysis of molten caustic soda (now called sodium hydroxide).[125]

[123][122] did include the alkali in his list of chemical elements in 1789.Antoine Lavoisier The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus [121] was able to prove this difference in 1736.Henri Louis Duhamel du Monceau and [120]

This article was sourced from Creative Commons Attribution-ShareAlike License; additional terms may apply. World Heritage Encyclopedia content is assembled from numerous content providers, Open Access Publishing, and in compliance with The Fair Access to Science and Technology Research Act (FASTR), Wikimedia Foundation, Inc., Public Library of Science, The Encyclopedia of Life, Open Book Publishers (OBP), PubMed, U.S. National Library of Medicine, National Center for Biotechnology Information, U.S. National Library of Medicine, National Institutes of Health (NIH), U.S. Department of Health & Human Services, and, which sources content from all federal, state, local, tribal, and territorial government publication portals (.gov, .mil, .edu). Funding for and content contributors is made possible from the U.S. Congress, E-Government Act of 2002.
Crowd sourced content that is contributed to World Heritage Encyclopedia is peer reviewed and edited by our editorial staff to ensure quality scholarly research articles.
By using this site, you agree to the Terms of Use and Privacy Policy. World Heritage Encyclopedia™ is a registered trademark of the World Public Library Association, a non-profit organization.

Copyright © World Library Foundation. All rights reserved. eBooks from Project Gutenberg are sponsored by the World Library Foundation,
a 501c(4) Member's Support Non-Profit Organization, and is NOT affiliated with any governmental agency or department.