Periodic table: Difference between revisions
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===Oxidation number=== |
===Oxidation number=== |
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With some minor exceptions, [[oxidation number]]s among the elements show four main trends according to their periodic table geographic location: left; middle; right; and south. On the left (groups 1 to 3), the highest most stable oxidation number is the group number, with lower oxidation states being less stable. In the middle (groups 4 to 11), higher oxidation states become more stable going down each group. Group 12 is an exception to this trend; they behave as if they were located on the left side of the table. On the right, higher oxidation states tend to become less stable going down a group.<ref name=Fernelius>{{cite journal |last1= Fernelius |first1=W. |last2=C. |title= Some reflections on the periodic table and its use|url= |journal= Journal of Chemical Education |volume=63 |issue= 3|pages=263–266 |doi=10.1021/ed063p263|access-date= |year=1986 |bibcode=1986JChEd..63..263F }}</ref> The shift between these trends is continuous: for example, group 3 also has lower oxidation states most stable in its lightest member (scandium, with CsScCl<sub>3</sub> for example known in the +2 state),<ref name="MeyerCorbett1981">{{cite journal|last1=Meyer|first1=Gerd.|last2=Corbett|first2=John D.|title=Reduced ternary halides of scandium: RbScX3 (X = chlorine, bromine) and CsScX3 (X = chlorine, bromine, iodine)|journal=Inorganic Chemistry|volume=20|issue=8|year=1981|pages=2627–2631|issn=0020-1669|doi=10.1021/ic50222a047}}</ref> and group 12 is predicted to have [[copernicium]] more readily showing oxidation states above +2.<ref name=CRNL |
With some minor exceptions, [[oxidation number]]s among the elements show four main trends according to their periodic table geographic location: left; middle; right; and south. On the left (groups 1 to 3), the highest most stable oxidation number is the group number, with lower oxidation states being less stable. In the middle (groups 4 to 11), higher oxidation states become more stable going down each group. Group 12 is an exception to this trend; they behave as if they were located on the left side of the table. On the right, higher oxidation states tend to become less stable going down a group.<ref name=Fernelius>{{cite journal |last1= Fernelius |first1=W. |last2=C. |title= Some reflections on the periodic table and its use|url= |journal= Journal of Chemical Education |volume=63 |issue= 3|pages=263–266 |doi=10.1021/ed063p263|access-date= |year=1986 |bibcode=1986JChEd..63..263F }}</ref> The shift between these trends is continuous: for example, group 3 also has lower oxidation states most stable in its lightest member (scandium, with CsScCl<sub>3</sub> for example known in the +2 state),<ref name="MeyerCorbett1981">{{cite journal|last1=Meyer|first1=Gerd.|last2=Corbett|first2=John D.|title=Reduced ternary halides of scandium: RbScX3 (X = chlorine, bromine) and CsScX3 (X = chlorine, bromine, iodine)|journal=Inorganic Chemistry|volume=20|issue=8|year=1981|pages=2627–2631|issn=0020-1669|doi=10.1021/ic50222a047}}</ref> and group 12 is predicted to have [[copernicium]] more readily showing oxidation states above +2.<ref name=CRNL/> |
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The lanthanides and actinides positioned along the south of the table are distinguished by having the +3 oxidation state in common; this is the most stable state for the lanthanides. The early actinides show a pattern of oxidation states somewhat similar to those of their period 6 and 7 transition metal congeners; the later actinides are more similar to the lanthanides, though the last ones (excluding lawrencium) have an increasingly important +2 oxidation state that becomes the most stable state for nobelium.<ref>{{cite book |last=Wiberg|first1=N.|title= Inorganic Chemistry|year=2001|pages=1644–1645|publisher=Academic Press |
The lanthanides and actinides positioned along the south of the table are distinguished by having the +3 oxidation state in common; this is the most stable state for the lanthanides. The early actinides show a pattern of oxidation states somewhat similar to those of their period 6 and 7 transition metal congeners; the later actinides are more similar to the lanthanides, though the last ones (excluding lawrencium) have an increasingly important +2 oxidation state that becomes the most stable state for nobelium.<ref>{{cite book |last=Wiberg|first1=N.|title= Inorganic Chemistry|year=2001|pages=1644–1645|publisher=Academic Press |
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From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block. In general, groups at the peripheries of blocks display similarities to the groups of the neighbouring blocks as well as to the other groups in their own blocks, as expected as most periodic trends are continuous.<ref name="MacKay">{{cite book |last1=MacKay |first1=K. M.|last2=MacKay |first2=R. A. |last3=Henderson |first3=W. |date=2002 |title=Introduction to Modern Inorganic Chemistry|edition=6th|publisher =Nelson Thornes|location=Cheltenham|isbn=978-0-7487-6420-4|pages=194–196}}</ref> These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, lanthanides, and heavy group 4 and 5 elements show some behaviour similar to the alkaline earth metals<ref>{{cite book |last=Remy |first=H.|date=1956 |title=Treatise on Inorganic Chemistry |volume=2|location=Amsterdam |publisher=Elsevier |page=30|editor-last=Kleinberg|editor-first=J.}}</ref> or, more generally, ''s'' block metals<ref>{{cite book |last1=Phillips |first1=C. S. G.|last2=Williams|first2=R. J. P.|date=1966 |title=Inorganic Chemistry|location=Oxford |publisher=Clarendon Press |pages=4–5}}</ref><ref>{{cite book |last=King |first=R. B.|date=1995 |title=Inorganic chemistry of main group elements|location=New York |publisher=Wiley-VCH |page=289}}</ref><ref name=Greenwood957/> but have some of the physical properties of ''d'' block transition metals |
From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block. In general, groups at the peripheries of blocks display similarities to the groups of the neighbouring blocks as well as to the other groups in their own blocks, as expected as most periodic trends are continuous.<ref name="MacKay">{{cite book |last1=MacKay |first1=K. M.|last2=MacKay |first2=R. A. |last3=Henderson |first3=W. |date=2002 |title=Introduction to Modern Inorganic Chemistry|edition=6th|publisher =Nelson Thornes|location=Cheltenham|isbn=978-0-7487-6420-4|pages=194–196}}</ref> These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, lanthanides, and heavy group 4 and 5 elements show some behaviour similar to the alkaline earth metals<ref>{{cite book |last=Remy |first=H.|date=1956 |title=Treatise on Inorganic Chemistry |volume=2|location=Amsterdam |publisher=Elsevier |page=30|editor-last=Kleinberg|editor-first=J.}}</ref> or, more generally, ''s'' block metals<ref>{{cite book |last1=Phillips |first1=C. S. G.|last2=Williams|first2=R. J. P.|date=1966 |title=Inorganic Chemistry|location=Oxford |publisher=Clarendon Press |pages=4–5}}</ref><ref>{{cite book |last=King |first=R. B.|date=1995 |title=Inorganic chemistry of main group elements|location=New York |publisher=Wiley-VCH |page=289}}</ref><ref name=Greenwood957/> but have some of the physical properties of ''d'' block transition metals.<ref name=Greenwood947/> In fact, the metals all the way up to group 6 are united by being class-A cations ([[HSAB theory|"hard" acids]]) that form more stable complexes with ligands whose donor atoms are the most electronegative nonmetals nitrogen, oxygen, and fluorine; metals later in the table form a transition to class-B cations ("soft" acids) that form more stable complexes with ligands whose donor atoms are the less electronegative heavier elements of groups 15 through 17.<ref>Greenwood and Earnshaw, p. 909</ref> |
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Meanwhile, lutetium behaves chemically as a lanthanide (with which it is often classified) but shows a mix of lanthanide and transition metal physical properties (as does yttrium).<ref>{{cite book |last1=Spedding |first1=F. H. |last2=Beadry |first2=B. J. |editor-last=Hampel |editor-first=C. A. |title=The Encyclopedia of the Chemical Elements |chapter-url=https://archive.org/details/encyclopediaofch00hamp |chapter-url-access=registration |publisher=Reinhold Book Corporation |date=1968 |pages=[https://archive.org/details/encyclopediaofch00hamp/page/374 374–378] |chapter=Lutetium}}</ref><ref>{{cite journal| last = Settouti | first = N.| last2 = Aourag | first2 = H. | date = 2014| title = A Study of the Physical and Mechanical Properties of Lutetium Compared with Those of Transition Metals: A Data Mining Approach| journal = JOM| volume = 67| issue = 1| pages = 87–93| doi = 10.1007/s11837-014-1247-x| bibcode = 2015JOM....67a..87S| url = https://www.semanticscholar.org/paper/54dbfdaead5d96c65fe32381422b8cd0927c1ddd}}</ref> Lawrencium, as an analogue of lutetium, would presumably display like characteristics.{{#tag:ref|While Lr is thought to have a p rather than d electron in its ground-state electron configuration, and would therefore be expected to be a volatile metal capable of forming a +1 cation in solution like thallium, no evidence of either of these properties has been able to be obtained despite experimental attempts to do so.<ref name=Silva1642>{{cite book|doi=10.1007/978-94-007-0211-0_13|title=The Chemistry of the Actinide and Transactinide Elements|pages=1621–1651|date=2011|isbn=978-94-007-0210-3|publisher=Springer |place=Netherlands|author=Silva, Robert J.|editor= Morss, Lester R.|editor2= Edelstein, Norman M.|editor3= Fuger, Jean |chapter=Chapter 13. Fermium, Mendelevium, Nobelium, and Lawrencium}}</ref> It was originally expected to have a d electron in its electron configuration<ref name=Silva1642/> and this may still be the case for metallic lawrencium, whereas gas phase atomic lawrencium is very likely thought to have a p electron.<ref name=Sato>{{cite journal |last=Sato |first=T. K. |last2=Asai |first2=M. |first3=A. |last3=Borschevsky |first4=T. |last4=Stora |first5=N. |last5=Sato |first6=Y. |last6=Kaneya |first7=K. |last7=Tsukada |first8=Ch. E. |last8=Düllman |first9=K. |last9=Eberhardt |first10=E. |last10=Eliav |first11=S. |last11=Ichikawa |first12=U. |last12=Kaldor |first13=J. V. |last13=Kratz |first14=S. |last14=Miyashita |first15=Y. |last15=Nagame |first16=K. |last16=Ooe |first17=A. |last17=Osa |first18=D. |last18=Renisch |first19=J. |last19=Runke |first20=M. |last20=Schädel |first21=P. |last21=Thörle-Pospiech |first22=A. |last22=Toyoshima |first23=N. |last23=Trautmann |date=9 April 2015 |title=Measurement of the first ionization potential of lawrencium, element 103 |journal=Nature |volume=520 |issue=7546 |pages=209–211 |doi=10.1038/nature14342 |pmid=25855457 |bibcode=2015Natur.520..209S |url=http://cds.cern.ch/record/2008656/files/TKSato-Lr-IP_prep_nature.pdf |access-date=25 October 2017 |archive-url=https://web.archive.org/web/20181030071405/http://cds.cern.ch/record/2008656/files/TKSato-Lr-IP_prep_nature.pdf |archive-date=30 October 2018 |url-status=live }}</ref>|group=n}} The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals.<ref>{{cite book |last=Steele |first=D. |title=The Chemistry of the Metallic Elements |publisher= Pergamon Press |location=Oxford |page=67}}</ref> The volatile group 12 metals, zinc, cadmium and mercury are sometimes regarded as linking the ''d'' block to the ''p'' block. Notionally they are ''d'' block elements but they have few transition metal properties and are more like their ''p'' block neighbors in group 13.<ref>{{cite book |last1=Greenwood|first1=N. N.|last2=Earnshaw|first2=A.|title=Chemistry of the Elements |publisher=Elsevier Science Ltd.|location=Oxford|edition=2nd|page=1206|isbn=978-0-7506-3365-9}}</ref><ref>{{cite book |last1=MacKay |first1=K. M.|last2=MacKay |first2=R. A. |last3=Henderson |first3=W. |date=2002 |title=Introduction to Modern Inorganic Chemistry|edition=6th|publisher =Nelson Thornes |location=Cheltenham |isbn=978-0-7487-6420-4 |pages=194–196, 385}}</ref> The relatively inert noble gases, in group 18, bridge the most reactive groups of elements in the periodic table—the halogens in group 17 and the alkali metals in group 1.<ref name="MacKay"/> |
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==History== |
==History== |
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Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely settled. They are commonly [[lanthanum]] and [[actinium]], and less often [[lutetium]] and [[lawrencium]]. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the ''f'' block elements start and end.<ref>{{cite book|last=Thyssen|first=P.|last2=Binnemans|first2=K.|editor1-last=Gschneidner Jr.|editor1-first= K. A.|editor2-last=Bünzli|editor2-first=J-C.G|editor3-last=Vecharsky|editor3-first=Bünzli|date=2011|title=Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis|journal=Handbook on the Physics and Chemistry of Rare Earths|publisher=Elsevier|location=Amsterdam|volume=41|pages=1–94|isbn=978-0-444-53590-0|doi=10.1016/B978-0-444-53590-0.00001-7}}</ref>{{#tag:ref|The detachment of the lanthanides from the main body of the periodic table has been attributed to the Czech chemist [[Bohuslav Brauner]] who, in 1902, allocated all of them ("Ce etc.") to one position in group 4, below zirconium. This arrangement was referred to as the "asteroid hypothesis", in analogy to asteroids occupying a single orbit in the solar system. Before this time the lanthanides were generally (and unsuccessfully) placed throughout groups I to VIII of the older 8-column form of periodic table. Although predecessors of Brauner's 1902 arrangement are recorded from as early as 1895, he is known to have referred to the "chemistry of asteroids" in an 1881 letter to Mendeleev. Other authors assigned all of the lanthanides to either group 3, groups 3 and 4, or groups 2, 3 and 4. In 1922 [[Niels Bohr]] continued the detachment process by locating the lanthanides between the s- and d-blocks. In 1949 [[Glenn T. Seaborg]] (re)introduced the form of periodic table that is popular today, in which the lanthanides and actinides appear as footnotes. Seaborg first published his table in a classified report dated 1944. It was published again by him in 1945 in ''[[Chemical and Engineering News]],'' and in the years up to 1949 several authors commented on, and generally agreed with, Seaborg's proposal. In that year he noted that the best method for presenting the actinides seemed to be by positioning them below, and as analogues of, the lanthanides. See: Thyssen P. and Binnemans K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In K. A. Gschneider Jr. (ed). ''Handbook on the Physics and Chemistry of the Rare Earths.'' '''41'''. Amsterdam: Elsevier, pp. 1–94; Seaborg G. T. (1994). Origin of the Actinide Concept'. In K. A. Gschneider Jr. (ed). ''Handbook on the Physics and Chemistry of the Rare Earths''. '''18'''. Amsterdam: Elsevier, pp. 1–27.|group=n}} It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks".<ref>{{cite journal |last=Stewart |first=P. J.|date=2008 |title=The Flyleaf Table: An Alternative|journal=Journal of Chemical Education|volume=85 |issue=11 |page=1490 |doi=10.1021/ed085p1490|bibcode = 2008JChEd..85.1490S |doi-access=free }}</ref> A third common variant shows the two positions below [[yttrium]] as being occupied by the lanthanides and the actinides. A fourth variant shows group 3 bifurcating after Sc-Y, into an La-Ac branch, and an Lu-Lr branch.<ref name="McGraw-Hill"/> |
Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely settled. They are commonly [[lanthanum]] and [[actinium]], and less often [[lutetium]] and [[lawrencium]]. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the ''f'' block elements start and end.<ref>{{cite book|last=Thyssen|first=P.|last2=Binnemans|first2=K.|editor1-last=Gschneidner Jr.|editor1-first= K. A.|editor2-last=Bünzli|editor2-first=J-C.G|editor3-last=Vecharsky|editor3-first=Bünzli|date=2011|title=Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis|journal=Handbook on the Physics and Chemistry of Rare Earths|publisher=Elsevier|location=Amsterdam|volume=41|pages=1–94|isbn=978-0-444-53590-0|doi=10.1016/B978-0-444-53590-0.00001-7}}</ref>{{#tag:ref|The detachment of the lanthanides from the main body of the periodic table has been attributed to the Czech chemist [[Bohuslav Brauner]] who, in 1902, allocated all of them ("Ce etc.") to one position in group 4, below zirconium. This arrangement was referred to as the "asteroid hypothesis", in analogy to asteroids occupying a single orbit in the solar system. Before this time the lanthanides were generally (and unsuccessfully) placed throughout groups I to VIII of the older 8-column form of periodic table. Although predecessors of Brauner's 1902 arrangement are recorded from as early as 1895, he is known to have referred to the "chemistry of asteroids" in an 1881 letter to Mendeleev. Other authors assigned all of the lanthanides to either group 3, groups 3 and 4, or groups 2, 3 and 4. In 1922 [[Niels Bohr]] continued the detachment process by locating the lanthanides between the s- and d-blocks. In 1949 [[Glenn T. Seaborg]] (re)introduced the form of periodic table that is popular today, in which the lanthanides and actinides appear as footnotes. Seaborg first published his table in a classified report dated 1944. It was published again by him in 1945 in ''[[Chemical and Engineering News]],'' and in the years up to 1949 several authors commented on, and generally agreed with, Seaborg's proposal. In that year he noted that the best method for presenting the actinides seemed to be by positioning them below, and as analogues of, the lanthanides. See: Thyssen P. and Binnemans K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In K. A. Gschneider Jr. (ed). ''Handbook on the Physics and Chemistry of the Rare Earths.'' '''41'''. Amsterdam: Elsevier, pp. 1–94; Seaborg G. T. (1994). Origin of the Actinide Concept'. In K. A. Gschneider Jr. (ed). ''Handbook on the Physics and Chemistry of the Rare Earths''. '''18'''. Amsterdam: Elsevier, pp. 1–27.|group=n}} It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks".<ref>{{cite journal |last=Stewart |first=P. J.|date=2008 |title=The Flyleaf Table: An Alternative|journal=Journal of Chemical Education|volume=85 |issue=11 |page=1490 |doi=10.1021/ed085p1490|bibcode = 2008JChEd..85.1490S |doi-access=free }}</ref> A third common variant shows the two positions below [[yttrium]] as being occupied by the lanthanides and the actinides. A fourth variant shows group 3 bifurcating after Sc-Y, into an La-Ac branch, and an Lu-Lr branch.<ref name="McGraw-Hill"/> |
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Since 1921, many chemical and physical arguments have been made in support of lutetium and lawrencium<ref>{{cite book |last1=Thyssen |first1=P. |last2=Binnemanns|first2=K. |editor-last1=Gschneidner Jr.|editor-first1=K. A. |editor-last2=Büzli |editor-first2=J-C. J.|editor-last3=Pecharsky |editor-first3=V. K.| title=Handbook on the Physics and Chemistry of Rare Earths | volume=41| publisher=Elsevier |date=2011 |pages=80–81|chapter=1: Accommodation of the rare earths in the periodic table: A historical analysis |isbn=978-0-444-53590-0|location=Amsterdam}}</ref><ref>{{cite book |last1=Keeler |first1=J. |last2=Wothers |first2=P. |publisher=Oxford University|location=Oxford |date=2014 |title=Chemical Structure and Reactivity: An Integrated Approach| isbn=978-0-19-960413-5|page=259}}</ref> but the majority of authors seem either unconvinced by them or unaware of them.<ref name="finally">{{cite journal|last=Scerri|first=E.|authorlink=Eric Scerri|year=2012|journal=Chemistry International|volume=34|issue=4|url=http://www.iupac.org/publications/ci/2012/3404/ud.html|title=Mendeleev's Periodic Table Is Finally Completed and What To Do about Group 3?|url-status=live|archiveurl=https://web.archive.org/web/20170705051357/https://www.iupac.org/publications/ci/2012/3404/ud.html|archivedate=5 July 2017|doi=10.1515/ci.2012.34.4.28|doi-access=free}}</ref><ref name=exotic/> Most working chemists are not aware there is any controversy.<ref name=exotic>{{Cite journal |url=http://www.nature.com/news/exotic-atom-struggles-to-find-its-place-in-the-periodic-table-1.17275 |title=Exotic atom struggles to find its place in the periodic table |last1=Castelvecchi |first1=D. |date=8 April 2015 |journal=Nature |access-date=20 September 2015 |url-status=live |archiveurl=https://web.archive.org/web/20151005164540/http://www.nature.com/news/exotic-atom-struggles-to-find-its-place-in-the-periodic-table-1.17275 |archivedate=5 October 2015 |doi=10.1038/nature.2015.17275 }}</ref> In December 2015 an [[International Union of Pure and Applied Chemistry|IUPAC]] project was established to make a recommendation on the matter, considering only the first two alternatives as possibilities.<ref>{{cite web|url=http://iupac.org/projects/project-details/?project_nr=2015-039-2-200|title=The constitution of group 3 of the periodic table|publisher=IUPAC|access-date=30 July 2016|date=2015|url-status=live|archiveurl=https://web.archive.org/web/20160705053631/http://iupac.org/projects/project-details/?project_nr=2015-039-2-200|archivedate=5 July 2016}}</ref> |
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====Lanthanum and actinium==== |
====Lanthanum and actinium==== |
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Lanthanum and actinium are commonly depicted as the remaining group 3 members.<ref>{{cite book |last=Emsley |first=J. |date=2011 |title=Nature's Building Blocks |publisher=Oxford University | edition=new|location =Oxford |isbn=978-0-19-960563-7|page=651}}</ref>{{#tag:ref|For examples of this table see [[P. W. Atkins|Atkins]] et al. (2006). ''Shriver & Atkins Inorganic Chemistry'' (4th ed.). Oxford: Oxford University Press • Myers et al. (2004). ''Holt Chemistry''. Orlando: Holt, Rinehart & Winston • [[Raymond Chang (chemist)|Chang R.]] (2000). ''Essential Chemistry'' (2nd ed.). Boston: McGraw-Hill|group=n}} It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the ground-state electron configurations of the elements and the notion of the differentiating electron. The ground-state configurations of [[caesium]], [[barium]] and lanthanum are [Xe]6s<sup>1</sup>, [Xe]6s<sup>2</sup> and [Xe]5d<sup>1</sup>6s<sup>2</sup>. Lanthanum thus emerges with a 5d differentiating electron and on these grounds it was considered to be "in group 3 as the first member of the d-block for period 6".<ref name=Jensen1982>{{cite journal |title=The Positions of Lanthanum (Actinium) and Lutetium (Lawrencium) in the Periodic Table |author=William B. Jensen |journal=J. Chem. Educ. |year=1982 |volume=59 |issue = 8|pages=634–636 |doi=10.1021/ed059p634|bibcode=1982JChEd..59..634J }}</ref> A superficially consistent set of electron configurations is then seen in group 3: [[scandium]] [Ar]3d<sup>1</sup>4s<sup>2</sup>, yttrium [Kr]4d<sup>1</sup>5s<sup>2</sup> and lanthanum [Xe]5d<sup>1</sup>6s<sup>2</sup>. Still in period 6, [[ytterbium]] was assigned an electron configuration of [Xe]4f<sup>13</sup>5d<sup>1</sup>6s<sup>2</sup> and [[lutetium]] [Xe]4f<sup>14</sup>5d<sup>1</sup>6s<sup>2</sup>, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6".<ref name=Jensen1982/> Later [[electron spectroscopy|spectroscopic]] work found that the electron configuration of ytterbium was in fact [Xe]4f<sup>14</sup>6s<sup>2</sup>. This meant that ytterbium and lutetium—the latter with [Xe]4f<sup>14</sup>5d<sup>1</sup>6s<sup>2</sup>—both had 14 f-electrons, "resulting in a d- rather than an f- differentiating electron" for lutetium and making it an "equally valid candidate" with [Xe]5d<sup>1</sup>6s<sup>2</sup> lanthanum, for the group 3 periodic table position below yttrium.<ref name=Jensen1982/> Lanthanum has the advantage of incumbency since the 5d<sup>1</sup> electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium<ref name=Trifonov>{{cite book |last=Trifonov|first1=D. N.| |
Lanthanum and actinium are commonly depicted as the remaining group 3 members.<ref>{{cite book |last=Emsley |first=J. |date=2011 |title=Nature's Building Blocks |publisher=Oxford University | edition=new|location =Oxford |isbn=978-0-19-960563-7|page=651}}</ref>{{#tag:ref|For examples of this table see [[P. W. Atkins|Atkins]] et al. (2006). ''Shriver & Atkins Inorganic Chemistry'' (4th ed.). Oxford: Oxford University Press • Myers et al. (2004). ''Holt Chemistry''. Orlando: Holt, Rinehart & Winston • [[Raymond Chang (chemist)|Chang R.]] (2000). ''Essential Chemistry'' (2nd ed.). Boston: McGraw-Hill|group=n}} It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the ground-state electron configurations of the elements and the notion of the differentiating electron. The ground-state configurations of [[caesium]], [[barium]] and lanthanum are [Xe]6s<sup>1</sup>, [Xe]6s<sup>2</sup> and [Xe]5d<sup>1</sup>6s<sup>2</sup>. Lanthanum thus emerges with a 5d differentiating electron and on these grounds it was considered to be "in group 3 as the first member of the d-block for period 6".<ref name=Jensen1982>{{cite journal |title=The Positions of Lanthanum (Actinium) and Lutetium (Lawrencium) in the Periodic Table |author=William B. Jensen |journal=J. Chem. Educ. |year=1982 |volume=59 |issue = 8|pages=634–636 |doi=10.1021/ed059p634|bibcode=1982JChEd..59..634J }}</ref> A superficially consistent set of electron configurations is then seen in group 3: [[scandium]] [Ar]3d<sup>1</sup>4s<sup>2</sup>, yttrium [Kr]4d<sup>1</sup>5s<sup>2</sup> and lanthanum [Xe]5d<sup>1</sup>6s<sup>2</sup>. Still in period 6, [[ytterbium]] was assigned an electron configuration of [Xe]4f<sup>13</sup>5d<sup>1</sup>6s<sup>2</sup> and [[lutetium]] [Xe]4f<sup>14</sup>5d<sup>1</sup>6s<sup>2</sup>, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6".<ref name=Jensen1982/> Later [[electron spectroscopy|spectroscopic]] work found that the electron configuration of ytterbium was in fact [Xe]4f<sup>14</sup>6s<sup>2</sup>. This meant that ytterbium and lutetium—the latter with [Xe]4f<sup>14</sup>5d<sup>1</sup>6s<sup>2</sup>—both had 14 f-electrons, "resulting in a d- rather than an f- differentiating electron" for lutetium and making it an "equally valid candidate" with [Xe]5d<sup>1</sup>6s<sup>2</sup> lanthanum, for the group 3 periodic table position below yttrium.<ref name=Jensen1982/> Lanthanum has the advantage of incumbency since the 5d<sup>1</sup> electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium<ref name=Trifonov>{{cite book |last=Trifonov|first1=D. N.| |
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title= Rare-earth elements and their position in the periodic system|type=translated from Russian|year=1970|pages=201–202|publisher=Indian National Scientific Documentation Centre|location=New Delhi}}</ref> (though similar logic would also lead to thorium getting the 6d<sup>2</sup> position, having incumbency over [[rutherfordium]]). |
title= Rare-earth elements and their position in the periodic system|type=translated from Russian|year=1970|pages=201–202|publisher=Indian National Scientific Documentation Centre|location=New Delhi}}</ref> (though similar logic would also lead to thorium getting the 6d<sup>2</sup> position, having incumbency over [[rutherfordium]]).{{cn|date=May 2020}} |
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Ground-state gas-phase configurations consider only isolated atoms as opposed to bonding atoms in compounds (the latter being more relevant for chemistry), which often show different configurations.<ref>{{cite book |last=Scerri |first=Eric |author-link=Eric Scerri |editor1-first=Eric |editor1-last=Scerri |editor2-first=Grant |editor2-last=Fisher |date=2016 |title=Essays in the Philosophy of Chemistry |publisher=Oxford University Press |chapter=The Changing Views of a Philosopher of Chemistry on the Question of Reduction}}</ref> Moreover, the lowest levels of two different configurations often are separated by only very small energies, that are minuscule compared to the spreading of ''J''-levels of each configuration (e.g. terbium, where the 285 cm<sup>−1</sup> difference between [Xe]4f<sup>8</sup>5d<sup>1</sup>6s<sup>2</sup> and the ground state [Xe]4f<sup>9</sup>5s<sup>2</sup> is much less than 1% of this spreading), making which configuration happens to be the ground state chemically quite irrelevant.<ref name=Jorgensen/> This form also creates an inconsistency in the treatment of thorium, which has no f-electrons in the ground-state (being [Rn]6d<sup>2</sup>7s<sup>2</sup>), similar to actinium as [Rn]6d<sup>1</sup>7s<sup>2</sup>; yet this form places thorium in the f-block but not actinium.<ref name=Jensen2009/> Considering only ground-state gas-phase configurations, thorium [Rn]6d<sup>2</sup>7s<sup>2</sup> by itself is just as good a homologue to zirconium [Kr]4d<sup>2</sup>5s<sup>2</sup> as lanthanum [Xe]5d<sup>1</sup>6s<sup>2</sup> is to scandium [Ar]3d<sup>1</sup>4s<sup>2</sup>;<ref name=Jorgensen/> yet thorium is invariably placed in the f-block. Thorium thus demonstrates that the possession of an f electron in the ground-state gas-phase configuration of an element is not necessary for it to belong to the f-block. Additionally, this form necessitates a split d-block if expanded to a 32-column periodic table.<ref>{{cite journal |last1=Scerri |first1=Eric |date=2009 |title=Which Elements Belong in Group 3? |url=https://pubs.acs.org/doi/pdf/10.1021/ed086p1188 |journal=Journal of Chemical Education |volume=86 |issue=10 |page=1188 |doi=10.1021/ed086p1188 |access-date=16 May 2020}}</ref> |
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In terms of chemical behaviour,<ref name=Harrington>{{cite book|last1=Greenwood|first1=N. N.|last2=Harrington|first2=T. J.|title=The chemistry of the transition elements|year=1973|page=[https://archive.org/details/chemistryoftrans0000earn/page/50 50]|publisher=Clarendon Press|location=Oxford|isbn=978-0-19-855435-6|url=https://archive.org/details/chemistryoftrans0000earn/page/50}}</ref> and trends going down group 3 (if Sc-Y-La is chosen) for properties such as melting point, electronegativity and ionic radius,<ref name= Aylward>{{cite book |last1= Aylward |first1=G.|last2= Findlay |first2=T.|title=SI chemical data|year=2008|publisher=John Wiley & Sons|location=Milton, Queensland|edition=6th|isbn=978-0-470-81638-7}}</ref><ref name=Wiberg>{{cite book |last=Wiberg|first1=N.|title= Inorganic Chemistry|year=2001|page=119|publisher=Academic Press |
In terms of chemical behaviour,<ref name=Harrington>{{cite book|last1=Greenwood|first1=N. N.|last2=Harrington|first2=T. J.|title=The chemistry of the transition elements|year=1973|page=[https://archive.org/details/chemistryoftrans0000earn/page/50 50]|publisher=Clarendon Press|location=Oxford|isbn=978-0-19-855435-6|url=https://archive.org/details/chemistryoftrans0000earn/page/50}}</ref> and trends going down group 3 (if Sc-Y-La is chosen) for properties such as melting point, electronegativity and ionic radius,<ref name= Aylward>{{cite book |last1= Aylward |first1=G.|last2= Findlay |first2=T.|title=SI chemical data|year=2008|publisher=John Wiley & Sons|location=Milton, Queensland|edition=6th|isbn=978-0-470-81638-7}}</ref><ref name=Wiberg>{{cite book |last=Wiberg|first1=N.|title= Inorganic Chemistry|year=2001|page=119|publisher=Academic Press |
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|isbn=978-0-12-352651-9|location=San Diego}}</ref> scandium, yttrium, lanthanum and actinium are similar to their group 1–2 counterparts. In this variant, the number of ''f'' electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block.<ref name=Wulfsberg>{{cite book |last=Wulfsberg|first1=G.|article=Periodic table: Trends in the properties of the elements|title= Encyclopedia of Inorganic Chemistry|year=2006|page=3|publisher=John Wiley & Sons|location=New York|isbn=978-0-470-86210-0}}</ref> For example, the f-electron counts for the trivalent ions of the first three f-block elements are Ce 1, Pr 2 and Nd 3.<ref name=Cotton>{{cite book |last=Cotton|first=S.|title= Lanthanide and Actinide Chemistry|year=2007|page=150|publisher=John Wiley & Sons|location=Chichester|isbn=978-0-470-01006-8}}</ref> However, outside the lanthanides there does not exist a typical oxidation state across any period of a block |
|isbn=978-0-12-352651-9|location=San Diego}}</ref> scandium, yttrium, lanthanum and actinium are similar to their group 1–2 counterparts, but at variance with the other groups in the early d-block. In this variant, the number of ''f'' electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block.<ref name=Wulfsberg>{{cite book |last=Wulfsberg|first1=G.|article=Periodic table: Trends in the properties of the elements|title= Encyclopedia of Inorganic Chemistry|year=2006|page=3|publisher=John Wiley & Sons|location=New York|isbn=978-0-470-86210-0}}</ref> For example, the f-electron counts for the trivalent ions of the first three f-block elements are Ce 1, Pr 2 and Nd 3.<ref name=Cotton>{{cite book |last=Cotton|first=S.|title= Lanthanide and Actinide Chemistry|year=2007|page=150|publisher=John Wiley & Sons|location=Chichester|isbn=978-0-470-01006-8}}</ref> However, outside the lanthanides there does not exist a typical oxidation state across any period of a block,<ref name=Greenwood27/> and the reason for this singular behaviour of the lanthanides in fact has very little to do with the electron configurations of the elements concerned.<ref name=Jorgensen2>{{cite journal |last1=Jørgensen |first1=Christian K. |date=1983 |title=5f Electrons in Transthorium Chemistry |url=https://core.ac.uk/download/pdf/85219145.pdf |journal=Radiochimica Acta |volume=32 |pages=1–5 |doi=10.1524/ract.1983.32.13.1 |access-date=16 May 2020}}</ref> |
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====Lutetium and lawrencium==== |
====Lutetium and lawrencium==== |
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||[[File:Periodic table 14LaAc form---Group 3 = Sc-Y-Lu-Lr.jpg|right|x100px]]<br /><small>Lu and Lr below Y</small> |
||[[File:Periodic table 14LaAc form---Group 3 = Sc-Y-Lu-Lr.jpg|right|x100px]]<br /><small>Lu and Lr below Y</small> |
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In other tables, lutetium and lawrencium are the remaining group 3 members.{{#tag:ref|For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G. & Overton T. (2013). ''Descriptive Inorganic Chemistry'' (6th ed.). New York: W. H. Freeman and Company • Brown et al. (2009). ''Chemistry: The Central Science'' (11th ed.). Upper Saddle River, New Jersey: Pearson Education • Moore et al. (1978). ''Chemistry''. Tokyo: McGraw-Hill Kogakusha|group=n}} Early techniques for chemically separating scandium, yttrium and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".<ref name=Jensen1982/> Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s.{{#tag:ref|The phenomenon of different separation groups is caused by increasing basicity with increasing radius, and does not constitute a fundamental reason to show Lu, rather than La, below Y. Thus, among the Group 2 [[alkaline earth metal]]s, Mg (less basic) belongs in the "soluble group" and Ca, Sr and Ba (more basic) occur in the "ammonium carbonate group". Nevertheless, Mg, Ca, Sr and Ba are routinely collocated in Group 2 of the periodic table. See: Moeller et al. (1989). ''Chemistry with Inorganic Qualitative Analysis'' (3rd ed.). SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.|group=n}} Several physicists in the 1950s and '60s favoured lutetium, in light of a comparison of several of its physical properties with those of lanthanum.<ref name=Jensen1982/> |
In other tables, lutetium and lawrencium are the remaining group 3 members.{{#tag:ref|For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G. & Overton T. (2013). ''Descriptive Inorganic Chemistry'' (6th ed.). New York: W. H. Freeman and Company • Brown et al. (2009). ''Chemistry: The Central Science'' (11th ed.). Upper Saddle River, New Jersey: Pearson Education • Moore et al. (1978). ''Chemistry''. Tokyo: McGraw-Hill Kogakusha|group=n}} Early techniques for chemically separating scandium, yttrium and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".<ref name=Jensen1982/> Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s.{{#tag:ref|The phenomenon of different separation groups is caused by increasing basicity with increasing radius, and does not constitute a fundamental reason to show Lu, rather than La, below Y. Thus, among the Group 2 [[alkaline earth metal]]s, Mg (less basic) belongs in the "soluble group" and Ca, Sr and Ba (more basic) occur in the "ammonium carbonate group". Nevertheless, Mg, Ca, Sr and Ba are routinely collocated in Group 2 of the periodic table. See: Moeller et al. (1989). ''Chemistry with Inorganic Qualitative Analysis'' (3rd ed.). SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.|group=n}} Several physicists in the 1950s and '60s favoured lutetium, in light of a comparison of several of its physical properties with those of lanthanum.<ref name=Jensen1982/> Among the prominent adherents of this form have been [[Lev Landau]] and [[Evgeny Lifshitz]], who wrote in their ''[[Course of Theoretical Physics]]'' (1958):<ref>{{cite book |
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|author=[[Lev Landau|L. D. Landau]], [[Evgeny Lifshitz|E. M. Lifshitz]] |
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|year=1958 |
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|title=Quantum Mechanics: Non-Relativistic Theory |
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|edition=1st |volume=Vol. 3 |
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|publisher=[[Pergamon Press]] |
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|pages=256–7 |
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}}</ref> |
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{{cquote|In books of chemistry, lutetium is usually placed in the rare-earth elements. This, however, is incorrect, since the 4f shell is complete in lutetium; it must therefore be placed in the platinum group [which they considered to be La+Lu–Pt]...|source=Landau and Lifshitz}} |
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This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. It has been argued that this is not a valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block.<ref>{{cite magazine|url= https://eic.rsc.org/opinion/five-ideas-in-chemical-education-that-must-die-part-five/2010032.article|title= Five ideas in chemical education that must die – Group three |last1= Scerri|first1= E.|date=15 September 2015| magazine=[[Education in Chemistry]] |publisher=[[Royal Society of Chemistry]]|access-date= 19 Sep 2015|quote= It is high time that the idea of group 3 consisting of Sc, Y, La and Ac is abandoned|url-status=live|archiveurl= https://web.archive.org/web/20151223190417/http://www.rsc.org/blogs/eic/2015/09/periodic-table-group-3|archivedate= 23 December 2015|df= dmy-all}}</ref> Karl Gschneidner, analysing the melting points of the lanthanides in a 1971 article, reached the conclusion that it was likely that 4f, 5d, 6s, and 6p electrons were all involved in the bonding of lanthanide metals except for lutetium, where 4f electrons were not found to be involved.<ref name=Gschneidner>{{cite journal |last1=Gschneidner Jr. |first1=K. A. |date=December 1971 |title=On the nature of 4ƒ bonding in the lanthanide elements and their compounds |journal=Journal of the Less Common Metals |volume=25 |issue=4 |pages=405–422 |doi=10.1016/0022-5088(71)90184-6}}</ref> The fact that lanthanum was demonstrated to be a 4f-band metal (with about 0.17 electrons per atom in fcc lanthanum, which is metastable at standard conditions)<ref name=Glotzel>{{cite journal |last1=Glotzel |first1=D. |date=1978 |title=Ground-state properties of f band metals: lanthanum, cerium and thorium |journal=Journal of Physics F: Metal Physics |volume=8 |issue=7 |pages=L163–L168 |doi=10.1088/0305-4608/8/7/004|bibcode=1978JPhF....8L.163G }}</ref> whereas the 4f shell appears to have no influence on the metallic properties of lutetium, has been used as an argument to place lutetium in group 3 instead of lanthanum.<ref name=Wittig>{{cite book |last=Wittig |first=Jörg |editor=H. J. Queisser |date=1973 |title=Festkörper Probleme: Plenary Lectures of the Divisions Semiconductor Physics, Surface Physics, Low Temperature Physics, High Polymers, Thermodynamics and Statistical Mechanics, of the German Physical Society, Münster, March 19–24, 1973 |chapter=The pressure variable in solid state physics: What about 4f-band superconductors? |location=Berlin, Heidelberg |publisher=Springer |page=375–396 |isbn=978-3-528-08019-8}}</ref> The 4f occupancy in solid lanthanum may explain some of its properties, such as its low melting point<ref name=Gschneidner/> and superconductivity.<ref name=Wittig/> Moreover, lanthanum has low-lying empty f orbitals which lutetium lacks,<ref name=Jensen2009>{{cite journal |last1=Jensen |first1=William B. |date=2009 |title=Misapplying the Periodic Law |url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/170.%20Misapplying%20the%20Periodic%20Law.pdf |journal=Journal of Chemical Education |volume=86 |issue=10 |pages=1186 |doi=10.1021/ed086p1186 |access-date=16 May 2020}}</ref> and it can have the 4f orbitals contributing measurably to the bonding, for example in [[lanthanum(III) fluoride]] (LaF<sub>3</sub>), which is not known for lutetium. While this contribution is small, it is greater for lanthanum than for any other lanthanide;<ref>{{cite journal |last1=Xu |first1=Wei |last2=Ji |first2=Wen-Xin |first3=Yi-Xiang |last3=Qiu |first4=W. H. Eugen |last4=Schwarz |first5=Shu-Guang |last5=Wang |date=2013 |title=On structure and bonding of lanthanoid trifluorides LnF<sub>3</sub> (Ln = La to Lu) |journal=Physical Chemistry Chemical Physics |volume=2013 |issue=15 |pages=7839–47 |doi=10.1039/C3CP50717C}}</ref> the order of involvement of 4f in lanthanum is similar to that of 5f in thorium.<ref name=Jorgensen>{{cite book |last=Jørgensen |first=Christian K. |date=1988 |title=Handbook on the Physics and Chemistry of Rare Earths |volume=11 |chapter-title=Influence of rare earths on chemical understanding and classification |pages=197–292 |doi=10.1016/S0168-1273(88)11007-6}}</ref> And while the trivalent lanthanides Pr<sup>3+</sup> through Yb<sup>3+</sup> show characteristic narrow bands with their positions almost completely independent on the ligands, the following 5d elements (along with the 3d and 4d elements) behave significantly differently; while both types of elements show electron-transfer bands, [[ligand field theory]] becomes important for these d elements.<ref name=Jorgensen2/> |
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Scandium, yttrium, and lutetium show a more consistent set of electron configurations matching the global trend on the periodic table: the 5d metals then all add a closed 4f<sup>14</sup> shell. For example, the shift from yttrium [Kr]4d<sup>1</sup>5s<sup>2</sup> to lutetium [Xe]4f<sup>14</sup>5d<sup>1</sup>6s<sup>2</sup> exactly parallels that from zirconium [Kr]4d<sup>2</sup>5s<sup>2</sup> to hafnium [Xe]4f<sup>14</sup>5d<sup>2</sup>6s<sup>2</sup>.<ref name=Jensen1982/> The inclusion of lutetium rather than lanthanum also homogenises the 5d transition series: trends in atomic size, coordination number, and relative abundance of metal–oxygen bonds all reveal that lutetium is closer than lanthanum to the behaviour of the uncontroversial 5d metals hafnium through mercury.<ref>{{cite journal |last1=Alvarez |first1=Santiago |date=11 February 2020 |title=The transition from 4f to 5d elements from the structural point of view |journal=[[CrystEngComm]] |volume=2020 |issue=Advance Article |doi=10.1039/D0CE00029A}}</ref> The same is true considering conduction band structures of the elements: lutetium has a transition-metal-like conduction band structure, but lanthanum does not.<ref>{{cite journal |last1=Merz |first1=H. |last2=Ulmer |first2=K. |date=1967 |title=Position of Lanthanum and Lutetium in the Periodic Table |journal=Physics Letters A |volume=26 |issue=1 |pages=6–7 |doi=10.1016/0375-9601(67)90527-0}}</ref> |
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As for lawrencium, its gas phase ground-state atomic electron configuration was confirmed in 2015 as [Rn]5f<sup>14</sup>7s<sup>2</sup>7p<sup>1</sup>. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for [[nihonium]] with its predicted configuration of [Rn]5f<sup>14</sup>6d<sup>10</sup>7s<sup>2</sup>7p<sup>1</sup>.<ref name=JensenLr>{{cite web|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |title=Some Comments on the Position of Lawrencium in the Periodic Table |last1=Jensen |first1=W. B. |date=2015 |access-date=20 September 2015 |url-status=dead |archiveurl=https://web.archive.org/web/20151223091325/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |archivedate=23 December 2015 |df= }}</ref> However, it is expected that in the condensed phase and in chemical environments lawrencium has the expected 6d occupancy, and simple modelling studies suggest it will behave like a lanthanide,<ref name=Xu>{{cite journal |last1=Xu |first1=W-H. |last2= Pyykkö|first2=P.|title=Is the chemistry of lawrencium peculiar? |journal=Physical Chemistry Chemical Physics |year=2016 |volume=18 |issue=26 |pages=17351–17355|doi=10.1039/C6CP02706G|pmid=27314425 |bibcode=2016PCCP...1817351X|hdl=10138/224395 |url=https://helda.helsinki.fi/bitstream/10138/224395/1/c6cp02706g.pdf }}</ref> in particular being a homologue of lutetium.<ref name= |
As for lawrencium, its gas phase ground-state atomic electron configuration was confirmed in 2015 as [Rn]5f<sup>14</sup>7s<sup>2</sup>7p<sup>1</sup>. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for [[nihonium]] with its predicted configuration of [Rn]5f<sup>14</sup>6d<sup>10</sup>7s<sup>2</sup>7p<sup>1</sup>.<ref name=JensenLr>{{cite web|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |title=Some Comments on the Position of Lawrencium in the Periodic Table |last1=Jensen |first1=W. B. |date=2015 |access-date=20 September 2015 |url-status=dead |archiveurl=https://web.archive.org/web/20151223091325/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |archivedate=23 December 2015 |df= }}</ref> However, it is expected that in the condensed phase and in chemical environments lawrencium has the expected 6d occupancy, and simple modelling studies suggest it will behave like a lanthanide,<ref name=Xu>{{cite journal |last1=Xu |first1=W-H. |last2= Pyykkö|first2=P.|title=Is the chemistry of lawrencium peculiar? |journal=Physical Chemistry Chemical Physics |year=2016 |volume=18 |issue=26 |pages=17351–17355|doi=10.1039/C6CP02706G|pmid=27314425 |bibcode=2016PCCP...1817351X|hdl=10138/224395 |url=https://helda.helsinki.fi/bitstream/10138/224395/1/c6cp02706g.pdf }}</ref> in particular being a homologue of lutetium. Lawrencium's return to +3 as the only stable oxidation state and being predicted to form a trivalent metal is distinct from the behaviour of the other late actinides fermium, mendelevium, and nobelium, which have a tendency towards forming lower oxidation states and form (or are predicted to form) divalent metals; it also makes an exception to the actinide contraction generally being larger than the analogous lanthanide contraction at the end of both series.<ref name="Silva">{{cite book |last=Silva |first=Robert J. |chapter=Fermium, Mendelevium, Nobelium, and Lawrencium |title=The Chemistry of the Actinide and Transactinide Elements |editor1-first=Lester R. |editor1-last=Morss |editor2-first=Norman M. |editor2-last=Edelstein |editor3-first=Jean |editor3-last=Fuger |edition=3rd |date=2006 |volume=3 |publisher=Springer |location=Dordrecht |pages=1621–1651 |chapter-url=http://radchem.nevada.edu/classes/rdch710/files/Fm%20to%20Lr.pdf |doi=10.1007/1-4020-3598-5_13 |isbn=978-1-4020-3555-5 |archive-url=https://web.archive.org/web/20100717155410/http://radchem.nevada.edu/classes/rdch710/files/Fm%20to%20Lr.pdf |archive-date=2010-07-17}}</ref> |
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While scandium, yttrium and lutetium (and lawrencium, so far as its chemistry is known) do often behave like trivalent versions of the group 1–2 metals, being hard class-A cations mostly restricted to the group oxidation state, they are not the only elements in the d-block or f-block that do so. The early transition metals zirconium and hafnium in group 4 also display such behaviour, as does the actinide thorium.<ref name=King>{{cite book|last=King|first=R. B.|title= Inorganic Chemistry of Main Group Elements|year=1995|page=289|publisher=Wiley-VCH|location=New York|isbn=978-1-56081-679-9}}</ref><ref name=Greenwood958>Greenwood and Earnshaw, p. 958</ref> The physical properties of the group 3 elements are affected by the presence of a d electron, which forms more localised bonds within the metals than the p electrons in the similar group 13 metals;<ref name=Greenwood947>Greenwood and Earnshaw, p. 947</ref> exactly the same situation is found comparing group 4 to group 14.<ref name=Greenwood957>Greenwood and Earnshaw, p. 957</ref> Trends going down group 3 (if Sc-Y-Lu is chosen) for properties such as melting point, electronegativity and ionic radius, are similar to those found among their group 4–8 counterparts.<ref name=Jensen1982/> In this variant, the number of ''f'' electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.<ref name=Jensen1982/> |
While scandium, yttrium and lutetium (and lawrencium, so far as its chemistry is known) do often behave like trivalent versions of the group 1–2 metals, being hard class-A cations mostly restricted to the group oxidation state, they are not the only elements in the d-block or f-block that do so. The early transition metals zirconium and hafnium in group 4, as well as niobium and tantalum in group 5, also display such behaviour, as does the actinide thorium. (The heavy group 4 elements and thorium are tetravalent; the heavy group 5 elements are pentavalent.)<ref name=King>{{cite book|last=King|first=R. B.|title= Inorganic Chemistry of Main Group Elements|year=1995|page=289|publisher=Wiley-VCH|location=New York|isbn=978-1-56081-679-9}}</ref><ref name=Greenwood958>Greenwood and Earnshaw, p. 958</ref> The physical properties of the group 3 elements are affected by the presence of a d electron, which forms more localised bonds within the metals than the p electrons in the similar group 13 metals;<ref name=Greenwood947>Greenwood and Earnshaw, p. 947</ref> exactly the same situation is found comparing group 4 to group 14.<ref name=Greenwood957>Greenwood and Earnshaw, p. 957</ref> Trends going down group 3 (if Sc-Y-Lu is chosen) for properties such as melting point, electronegativity and ionic radius, are similar to those found among their group 4–8 counterparts.<ref name=Jensen1982/> In this variant, the number of ''f'' electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.<ref name=Jensen1982/> |
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====Lanthanides and actinides==== |
====Lanthanides and actinides==== |
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A few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). |
A few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). |
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This variant, which is stated in the 2005 ''[[IUPAC nomenclature of inorganic chemistry#Nomenclature of Inorganic Chemistry|Red Book]]'' to be the IUPAC-agreed version as of 2005 (a number of later versions exist, and the last update is from 1 December 2018),<ref>{{cite book|title=Nomenclature of Inorganic Chemistry: IUPAC Recommendations 2005|last1=Connelly|first1=N. G.|last2=Damhus|first2=T.|last3=Hartshorn|first3=R. M.|last4=Hutton|first4=A. T.|year=2005|publisher=RSC Publishing|isbn=978-0-85404-438-2|page=vii|url=https://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|quote=Lesser omissions include ... the several different outdated versions of the periodic table. (That on the inside front cover is the current IUPAC-agreed version.)|access-date=26 November 2018|archive-url=https://web.archive.org/web/20181123034019/http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|archive-date=23 November 2018|url-status=live}}</ref>{{#tag:ref|Notwithstanding, an IUPAC member subsequently wrote that, "IUPAC has not approved any specific form of the periodic table, and an IUPAC-approved form does not exist, though even members of IUPAC themselves have published diagrams titled “IUPAC Periodic Table of the Elements". However, the only specific recommendation IUPAC has made concerning the periodic table covers the Group numbering of 1–18."<ref>{{cite journal |last=Leigh |first=G. J. |year=2009 |title=Periodic Tables and IUPAC |journal=[[Chemistry International]] |volume=31 |issue=1 |url=https://old.iupac.org/publications/ci/2009/3101/1_leigh.html |accessdate=27 November 2018 |doi=10.1515/ci.2009.31.1.4 |archive-url=https://web.archive.org/web/20181127022628/https://old.iupac.org/publications/ci/2009/3101/1_leigh.html |archive-date=27 November 2018 |url-status=live }}</ref>|group=n}} emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide ''f'' block (there can only be 14 elements in any row of the ''f'' block).{{#tag:ref|For examples of the group 3 = Ln and An table see Housecroft C. E. & Sharpe A. G. (2008). ''Inorganic Chemistry'' (3rd ed.). Harlow: Pearson Education • Halliday et al. (2005). ''Fundamentals of Physics'' (7th ed.). Hoboken, NewJersey: John Wiley & Sons • Nebergall et al. (1980). ''General Chemistry'' (6th ed.). Lexington: D. C. Heath and Company|group=n}} |
This variant, which is stated in the 2005 ''[[IUPAC nomenclature of inorganic chemistry#Nomenclature of Inorganic Chemistry|Red Book]]'' to be the IUPAC-agreed version as of 2005 (a number of later versions exist, and the last update is from 1 December 2018),<ref>{{cite book|title=Nomenclature of Inorganic Chemistry: IUPAC Recommendations 2005|last1=Connelly|first1=N. G.|last2=Damhus|first2=T.|last3=Hartshorn|first3=R. M.|last4=Hutton|first4=A. T.|year=2005|publisher=RSC Publishing|isbn=978-0-85404-438-2|page=vii|url=https://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|quote=Lesser omissions include ... the several different outdated versions of the periodic table. (That on the inside front cover is the current IUPAC-agreed version.)|access-date=26 November 2018|archive-url=https://web.archive.org/web/20181123034019/http://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|archive-date=23 November 2018|url-status=live}}</ref>{{#tag:ref|Notwithstanding, an IUPAC member subsequently wrote that, "IUPAC has not approved any specific form of the periodic table, and an IUPAC-approved form does not exist, though even members of IUPAC themselves have published diagrams titled “IUPAC Periodic Table of the Elements". However, the only specific recommendation IUPAC has made concerning the periodic table covers the Group numbering of 1–18."<ref>{{cite journal |last=Leigh |first=G. J. |year=2009 |title=Periodic Tables and IUPAC |journal=[[Chemistry International]] |volume=31 |issue=1 |url=https://old.iupac.org/publications/ci/2009/3101/1_leigh.html |accessdate=27 November 2018 |doi=10.1515/ci.2009.31.1.4 |archive-url=https://web.archive.org/web/20181127022628/https://old.iupac.org/publications/ci/2009/3101/1_leigh.html |archive-date=27 November 2018 |url-status=live }}</ref>|group=n}} emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide ''f'' block (there can only be 14 elements in any row of the ''f'' block).{{#tag:ref|For examples of the group 3 = Ln and An table see Housecroft C. E. & Sharpe A. G. (2008). ''Inorganic Chemistry'' (3rd ed.). Harlow: Pearson Education • Halliday et al. (2005). ''Fundamentals of Physics'' (7th ed.). Hoboken, NewJersey: John Wiley & Sons • Nebergall et al. (1980). ''General Chemistry'' (6th ed.). Lexington: D. C. Heath and Company|group=n}} However, this similarity does not extend to the 15 actinide elements (Ac–Lr), which show a much wider variety in their chemistries.<ref name=Jorgensen/> This form moreover reduces the f-block to a degenerate branch of group 3 of the d-block; it dates back to the 1920s when the lanthanides were thought to have their f electrons as core electrons, which is now known to be false. It is also false for the actinides, many of which show stable oxidation states above +3.<ref>{{cite journal |last1=Jensen |first1=William B. |date=2008 |title=The Periodic Table: Facts or Committees? |journal=Journal of Chemical Education |volume=85 |issue=11 |pages=1491–2 |doi=10.1021/ed085p1491.2}}</ref> |
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====La-Ac and Lu-Lr==== |
====La-Ac and Lu-Lr==== |
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===Elements with unknown chemical properties=== |
===Elements with unknown chemical properties=== |
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Although all elements up to oganesson have been discovered, of the elements above [[hassium]] (element 108), only [[copernicium]] (element 112), [[nihonium]] (element 113), and [[flerovium]] (element 114) have known chemical properties, and |
Although all elements up to oganesson have been discovered, of the elements above [[hassium]] (element 108), only [[copernicium]] (element 112), [[nihonium]] (element 113), and [[flerovium]] (element 114) have known chemical properties, and conclusive categorisation at present has not been reached.<ref name=CRNL>{{cite journal |last=Mewes |first=J.-M. |last2=Smits |first2=O. R. |last3=Kresse |first3=G. |last4=Schwerdtfeger |first4=P. |title=Copernicium is a Relativistic Noble Liquid |journal=Angewandte Chemie International Edition <!-- |volume= |issue= --> |date=2019 | |doi=10.1002/anie.201906966 |url=https://www.researchgate.net/publication/336389017}}</ref> Some of these may behave differently from what would be predicted by extrapolation, due to [[relativistic quantum chemistry|relativistic]] effects; for example, copernicium and flerovium have been predicted to possibly exhibit some noble-gas-like properties, even though neither is placed in group 18 with the other noble gases.<ref name=CRNL/><ref>{{cite book|title=The Chemistry of Superheavy Elements|last=Schändel|first=M.|year=2003|publisher=Kluwer Academic Publishers|location=Dordrecht|isbn=978-1-4020-1250-1|page=277}}</ref> The current experimental evidence still leaves open the question of whether copernicium and flerovium behave more like metals or noble gases.<ref name=CRNL/><ref>{{cite conference |url=http://www.epj-conferences.org/articles/epjconf/pdf/2016/26/epjconf-NS160-07003.pdf |title=Gas-phase chemistry of element 114, flerovium |last1=Yakushev |first1=Alexander |last2=Eichler |first2=Robert |date=2016 |conference=Nobel Symposium NS160 – Chemistry and Physics of Heavy and Superheavy Elements |doi=10.1051/epjconf/201613107003 |access-date=3 April 2018 |archive-url=https://web.archive.org/web/20170331024824/http://www.epj-conferences.org/articles/epjconf/pdf/2016/26/epjconf-NS160-07003.pdf |archive-date=31 March 2017 |url-status=live }}</ref> At the same time, [[oganesson]] (element 118) is expected to be a solid [[semiconductor]] at standard conditions, despite being in group 18.<ref name=semiconductor>{{cite journal |last1=Mewes |first1=Jan-Michael |last2=Smits |first2=Odile Rosette |first3=Paul |last3=Jerabek |first4=Peter |last4=Schwerdtfeger |date=25 July 2019 |title=Oganesson is a Semiconductor: On the Relativistic Band‐Gap Narrowing in the Heaviest Noble‐Gas Solids |url=https://onlinelibrary.wiley.com/doi/epdf/10.1002/anie.201908327 |journal=Angewandte Chemie |doi=10.1002/anie.201908327 |access-date=8 August 2019}}</ref> |
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===Further periodic table extensions=== |
===Further periodic table extensions=== |
Revision as of 17:03, 16 May 2020
![](http://upload.wikimedia.org/wikipedia/commons/thumb/2/2e/Simple_Periodic_Table_Chart-en.svg/330px-Simple_Periodic_Table_Chart-en.svg.png)
The periodic table, also known as the periodic table of elements, is a tabular display of the chemical elements, which are arranged by atomic number, electron configuration, and recurring chemical properties. The structure of the table shows periodic trends. The seven rows of the table, called periods, generally have metals on the left and nonmetals on the right. The columns, called groups, contain elements with similar chemical behaviours. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens; and group 18 are the noble gases. Also displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals.
The elements from atomic numbers 1 (hydrogen) through 118 (oganesson) have been discovered or synthesized, completing seven full rows of the periodic table.[1][2] The first 94 elements, hydrogen through plutonium, all occur naturally, though some are found only in trace amounts and a few were discovered in nature only after having first been synthesized.[n 1] Elements 95 to 118 have only been synthesized in laboratories or nuclear reactors.[3] The synthesis of elements having higher atomic numbers is currently being pursued: these elements would begin an eighth row, and theoretical work has been done to suggest possible candidates for this extension. Numerous synthetic radioisotopes of naturally occurring elements have also been produced in laboratories.
The organization of the periodic table can be used to derive relationships between the various element properties, and also to predict chemical properties and behaviours of undiscovered or newly synthesized elements. Russian chemist Dmitri Mendeleev published the first recognizable periodic table in 1869, developed mainly to illustrate periodic trends of the then-known elements. He also predicted some properties of unidentified elements that were expected to fill gaps within the table. Most of his forecasts proved to be correct. Mendeleev's idea has been slowly expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour. The modern periodic table now provides a useful framework for analyzing chemical reactions, and continues to be widely used in chemistry, nuclear physics and other sciences. Some discussion remains ongoing regarding the placement and categorisation of specific elements, the future extension and limits of the table, and whether there is an optimal form of the table.
Part of a series on the |
Periodic table |
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Overview
Group | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |
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Hydrogen & alkali metals |
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- Ca: 40.078 — Abridged value (uncertainty omitted here)[5]
- Po: [209] — mass number of the most stable isotope
Each chemical element has a unique atomic number (Z) representing the number of protons in its nucleus.[n 2] Most elements have differing numbers of neutrons among different atoms, with these variants being referred to as isotopes. For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and a very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element. Elements with no stable isotopes have the atomic masses of their most stable isotopes, where such masses are shown, listed in parentheses.[7]
In the standard periodic table, the elements are listed in order of increasing atomic number Z. A new row (period) is started when a new electron shell has its first electron. Columns (groups) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen and selenium are in the same column because they both have four electrons in the outermost p-subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.[8]
Since 2016, the periodic table has 118 confirmed elements, from element 1 (hydrogen) to 118 (oganesson). Elements 113, 115, 117 and 118, the most recent discoveries, were officially confirmed by the International Union of Pure and Applied Chemistry (IUPAC) in December 2015. Their proposed names, nihonium (Nh), moscovium (Mc), tennessine (Ts) and oganesson (Og) respectively, were made official in November 2016 by IUPAC.[9][10][11][12]
The first 94 elements occur naturally; the remaining 24, americium to oganesson (95–118), occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements.[3] No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form, nor has astatine (element 85); francium (element 87) has been only photographed in the form of light emitted from microscopic quantities (300,000 atoms).[13]
Grouping methods
Groups
A group or family is a vertical column in the periodic table. Groups usually have more significant periodic trends than periods and blocks, explained below. Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group generally have the same electron configurations in their valence shell.[14] Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number.[15] In some parts of the periodic table, such as the d-block and the f-block, horizontal similarities can be as important as, or more pronounced than, vertical similarities.[16][17][18]
Under an international naming convention, the groups are numbered numerically from 1 to 18 from the leftmost column (the alkali metals) to the rightmost column (the noble gases).[19] Previously, they were known by roman numerals. In America, the roman numerals were followed by either an "A" if the group was in the s- or p-block, or a "B" if the group was in the d-block. The roman numerals used correspond to the last digit of today's naming convention (e.g. the group 4 elements were group IVB, and the group 14 elements were group IVA). In Europe, the lettering was similar, except that "A" was used if the group was before group 10, and "B" was used for groups including and after group 10. In addition, groups 8, 9 and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII. In 1988, the new IUPAC naming system was put into use, and the old group names were deprecated.[20]
Some of these groups have been given trivial (unsystematic) names, as seen in the table below, although some are rarely used. Groups 3–10 have no trivial names and are referred to simply by their group numbers or by the name of the first member of their group (such as "the scandium group" for group 3),[19] since they display fewer similarities and/or vertical trends.
Elements in the same group tend to show patterns in atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group has a top-to-bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus.[21] There are exceptions to these trends: for example, in group 11, electronegativity increases farther down the group.[22]
IUPAC group | 1a | 2 | —b | 3c | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 |
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Mendeleev (I–VIII) | I | II | III | IV | V | VI | VII | VIII | I | II | III | IV | V | VI | VII | d | |||
CAS (US, A-B-A) | IA | IIA | IIIB | IVB | VB | VIB | VIIB | VIIIB | IB | IIB | IIIA | IVA | VA | VIA | VIIA | VIIIA | |||
Old IUPAC (Europe, A-B) | IA | IIA | IIIA | IVA | VA | VIA | VIIA | VIII | IB | IIB | IIIB | IVB | VB | VIB | VIIB | 0 | |||
Trivial namer | H and alkali metals | alkaline earth metals | triels | tetrels | pnictogens | chalcogens | halogens | noble gases | |||||||||||
Name by elementr | lithium group | beryllium group | scandium group | titanium group | vanadium group | chromium group | manganese group | iron group | cobalt group | nickel group | copper group | zinc group | boron group | carbon group | nitrogen group | oxygen group | fluorine group | helium or neon group | |
Period 1 | H | He | |||||||||||||||||
Period 2 | Li | Be | B | C | N | O | F | Ne | |||||||||||
Period 3 | Na | Mg | Al | Si | P | S | Cl | Ar | |||||||||||
Period 4 | K | Ca | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | |
Period 5 | Rb | Sr | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | |
Period 6 | Cs | Ba | La–Yb | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn |
Period 7 | Fr | Ra | Ac–No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og |
b The 14 f-block groups (columns) do not have a group number.
c The correct composition of group 3 is scandium (Sc), yttrium (Y), lutetium (Lu), and lawrencium (Lr), as shown here: this is endorsed by 1988[23] and 2021[24] IUPAC reports on the question. General inorganic chemistry texts often put scandium (Sc), yttrium (Y), lanthanum (La), and actinium (Ac) in group 3, so that Ce–Lu and Th–Lr become the f-block between groups 3 and 4; this was based on incorrectly measured electron configurations from history,[25] and Lev Landau and Evgeny Lifshitz already considered it incorrect in 1948.[26] Arguments can still occasionally be encountered in the contemporary literature purporting to defend it, but most authors consider them logically inconsistent.[27][28][29] Some sources follow a compromise that puts La–Lu and Ac–Lr as the f-block rows (despite that giving 15 f-block elements in each row, which contradicts quantum mechanics), leaving the heavier members of group 3 ambiguous.[24] See also Group 3 element#Composition.
d Group 18, the noble gases, were not discovered at the time of Mendeleev's original table. Later (1902), Mendeleev accepted the evidence for their existence, and they could be placed in a new "group 0", consistently and without breaking the periodic table principle.
r Group name as recommended by IUPAC.
Periods
A period is a horizontal row in the periodic table. Although groups generally have more significant periodic trends, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the lanthanides and actinides form two substantial horizontal series of elements.[30]
Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron, which causes the electron to be drawn closer to the nucleus.[31] This decrease in atomic radius also causes the ionization energy to increase when moving from left to right across a period. The more tightly bound an element is, the more energy is required to remove an electron. Electronegativity increases in the same manner as ionization energy because of the pull exerted on the electrons by the nucleus.[21] Electron affinity also shows a slight trend across a period. Metals (left side of a period) generally have a lower electron affinity than nonmetals (right side of a period), with the exception of the noble gases.[32]
Blocks
![](http://upload.wikimedia.org/wikipedia/commons/thumb/f/f2/Periodic_table_blocks_spdf_%2832_column%29.svg/350px-Periodic_table_blocks_spdf_%2832_column%29.svg.png)
Specific regions of the periodic table can be referred to as blocks in recognition of the sequence in which the electron shells of the elements are filled. Elements are assigned to blocks by what orbitals their valence electrons or vacancies lie in.[27] The s-block comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. The p-block comprises the last six groups, which are groups 13 to 18 in IUPAC group numbering (3A to 8A in American group numbering) and contains, among other elements, all of the metalloids. The d-block comprises groups 3 to 12 (or 3B to 2B in American group numbering) and contains all of the transition metals. The f-block, often offset below the rest of the periodic table, has no group numbers and comprises most of the lanthanides and actinides. A hypothetical g-block is expected to begin around element 121, a few elements away from what is currently known.[33]
Metals, metalloids and nonmetals
![](http://upload.wikimedia.org/wikipedia/commons/thumb/7/78/Periodic_table_%28metals%E2%80%93metalloids%E2%80%93nonmetals%2C_32_columns%29.png/300px-Periodic_table_%28metals%E2%80%93metalloids%E2%80%93nonmetals%2C_32_columns%29.png)
According to their shared physical and chemical properties, the elements can be classified into the major categories of metals, metalloids and nonmetals. Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals (other than noble gases). A majority of nonmetals are coloured or colourless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding. In between metals and nonmetals are metalloids, which have intermediate or mixed properties.[34]
Metal and nonmetals can be further classified into subcategories that show a gradation from metallic to non-metallic properties, when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals. Nonmetals may be simply subdivided into the polyatomic nonmetals, being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals, nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals, are examples of subsets of transition metals, also known[35] and occasionally denoted.[36]
Placing elements into categories and subcategories based just on shared properties is imperfect. There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes.[37] Beryllium, for example, is classified as an alkaline earth metal although its amphoteric chemistry and tendency to mostly form covalent compounds are both attributes of a chemically weak or post-transition metal. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals. Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories, or crystalline structures. Categorizing the elements in this fashion dates back to at least 1869 when Hinrichs[38] wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements.
Periodic trends and patterns
Electron configuration
![](http://upload.wikimedia.org/wikipedia/commons/thumb/9/95/Klechkovski_rule.svg/200px-Klechkovski_rule.svg.png)
The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more subshells (named s, p, d, f and g). As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram. The electron configuration for neon, for example, is 1s2 2s2 2p6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.[39][40]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/0/06/Periodic_trends.png/300px-Periodic_trends.png)
Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behaviour, some examples of which are shown in the diagrams below for atomic radii, ionization energy and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was developed, that led to the establishment of the periodic law (the properties of the elements recur at varying intervals) and the formulation of the first periodic tables.[39][40]
Atomic radii
![](http://upload.wikimedia.org/wikipedia/commons/thumb/b/bc/Empirical_atomic_radius_trends.png/300px-Empirical_atomic_radius_trends.png)
Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory.[41]
The electrons in the 4f-subshell, which is progressively filled from lanthanum (element 57) to ytterbium (element 70),[43] are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them.[44] Hence lutetium has virtually the same atomic radius (and chemistry) as yttrium, hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. This is an effect of the lanthanide contraction: a similar actinide contraction also exists. The effect of the lanthanide contraction is noticeable up to platinum (element 78), after which it is masked by a relativistic effect known as the inert pair effect.[45] The d-block contraction, which is a similar effect between the d-block and p-block, is less pronounced than the lanthanide contraction but arises from a similar cause.[44]
Ionization energy
![](http://upload.wikimedia.org/wikipedia/commons/thumb/1/1d/First_Ionization_Energy.svg/500px-First_Ionization_Energy.svg.png)
The first ionization energy is the energy it takes to remove one electron from an atom, the second ionization energy is the energy it takes to remove a second electron from the atom, and so on. For a given atom, successive ionization energies increase with the degree of ionization. For magnesium as an example, the first ionization energy is 738 kJ/mol and the second is 1450 kJ/mol. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table.[45]
Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas (complete electron shell) configuration. For magnesium again, the first two molar ionization energies of magnesium given above correspond to removing the two 3s electrons, and the third ionization energy is a much larger 7730 kJ/mol, for the removal of a 2p electron from the very stable neon-like configuration of Mg2+. Similar jumps occur in the ionization energies of other third-row atoms.[45]
Electronegativity
![](http://upload.wikimedia.org/wikipedia/commons/thumb/b/b4/Periodic_variation_of_Pauling_electronegativities.png/300px-Periodic_variation_of_Pauling_electronegativities.png)
Electronegativity is the tendency of an atom to attract a shared pair of electrons.[46] An atom's electronegativity is affected by both its atomic number and the distance between the valence electrons and the nucleus. The higher its electronegativity, the more an element attracts electrons. It was first proposed by Linus Pauling in 1932.[47] In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is the most electronegative of the elements,[n 4] while caesium is the least, at least of those elements for which substantial data is available.[22]
There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon respectively because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity.[22] The anomalously high electronegativity of lead, particularly when compared to thallium and bismuth, is an artifact of electronegativity varying with oxidation state: its electronegativity conforms better to trends if it is quoted for the +2 state instead of the +4 state.[48]
Electron affinity
![](http://upload.wikimedia.org/wikipedia/commons/thumb/6/6c/Electron_affinity_of_the_elements.svg/400px-Electron_affinity_of_the_elements.svg.png)
The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion. Although electron affinity varies greatly, some patterns emerge. Generally, nonmetals have more positive electron affinity values than metals. Chlorine most strongly attracts an extra electron. The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values.[51]
Electron affinity generally increases across a period. This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and is therefore more stable.[51]
A trend of decreasing electron affinity going down groups would be expected. The additional electron will be entering an orbital farther away from the nucleus. As such this electron would be less attracted to the nucleus and would release less energy when added. In going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congenors. Largely, this is due to the poor shielding by d and f electrons. A uniform decrease in electron affinity only applies to group 1 atoms.[52]
Metallic character
The lower the values of ionization energy, electronegativity and electron affinity, the more metallic character the element has. Conversely, nonmetallic character increases with higher values of these properties.[53] Given the periodic trends of these three properties, metallic character tends to decrease going across a period (or row) and, with some irregularities (mostly) due to poor screening of the nucleus by d and f electrons, and relativistic effects,[54] tends to increase going down a group (or column or family). Thus, the most metallic elements (such as caesium and francium) are found at the bottom left of traditional periodic tables and the most nonmetallic elements (oxygen, fluorine, chlorine) at the top right. The combination of horizontal and vertical trends in metallic character explains the stair-shaped dividing line between metals and nonmetals found on some periodic tables, and the practice of sometimes categorizing several elements adjacent to that line, or elements adjacent to those elements, as metalloids.[55][56]
Oxidation number
With some minor exceptions, oxidation numbers among the elements show four main trends according to their periodic table geographic location: left; middle; right; and south. On the left (groups 1 to 3), the highest most stable oxidation number is the group number, with lower oxidation states being less stable. In the middle (groups 4 to 11), higher oxidation states become more stable going down each group. Group 12 is an exception to this trend; they behave as if they were located on the left side of the table. On the right, higher oxidation states tend to become less stable going down a group.[57] The shift between these trends is continuous: for example, group 3 also has lower oxidation states most stable in its lightest member (scandium, with CsScCl3 for example known in the +2 state),[58] and group 12 is predicted to have copernicium more readily showing oxidation states above +2.[59]
The lanthanides and actinides positioned along the south of the table are distinguished by having the +3 oxidation state in common; this is the most stable state for the lanthanides. The early actinides show a pattern of oxidation states somewhat similar to those of their period 6 and 7 transition metal congeners; the later actinides are more similar to the lanthanides, though the last ones (excluding lawrencium) have an increasingly important +2 oxidation state that becomes the most stable state for nobelium.[60]
Linking or bridging groups
From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block. In general, groups at the peripheries of blocks display similarities to the groups of the neighbouring blocks as well as to the other groups in their own blocks, as expected as most periodic trends are continuous.[61] These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, lanthanides, and heavy group 4 and 5 elements show some behaviour similar to the alkaline earth metals[62] or, more generally, s block metals[63][64][65] but have some of the physical properties of d block transition metals.[66] In fact, the metals all the way up to group 6 are united by being class-A cations ("hard" acids) that form more stable complexes with ligands whose donor atoms are the most electronegative nonmetals nitrogen, oxygen, and fluorine; metals later in the table form a transition to class-B cations ("soft" acids) that form more stable complexes with ligands whose donor atoms are the less electronegative heavier elements of groups 15 through 17.[67]
Meanwhile, lutetium behaves chemically as a lanthanide (with which it is often classified) but shows a mix of lanthanide and transition metal physical properties (as does yttrium).[68][69] Lawrencium, as an analogue of lutetium, would presumably display like characteristics.[n 5] The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals.[72] The volatile group 12 metals, zinc, cadmium and mercury are sometimes regarded as linking the d block to the p block. Notionally they are d block elements but they have few transition metal properties and are more like their p block neighbors in group 13.[73][74] The relatively inert noble gases, in group 18, bridge the most reactive groups of elements in the periodic table—the halogens in group 17 and the alkali metals in group 1.[61]
History
First systemization attempts
![](http://upload.wikimedia.org/wikipedia/commons/thumb/3/3d/Discovery_of_chemical_elements.svg/500px-Discovery_of_chemical_elements.svg.png)
In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths.[75] Chemists spent the following century searching for a more precise classification scheme. In 1829, Johann Wolfgang Döbereiner observed that many of the elements could be grouped into triads based on their chemical properties. Lithium, sodium, and potassium, for example, were grouped together in a triad as soft, reactive metals. Döbereiner also observed that, when arranged by atomic weight, the second member of each triad was roughly the average of the first and the third.[76] This became known as the Law of Triads.[77] German chemist Leopold Gmelin worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. Jean-Baptiste Dumas published work in 1857 describing relationships between various groups of metals. Although various chemists were able to identify relationships between small groups of elements, they had yet to build one scheme that encompassed them all.[76] In 1857, German chemist August Kekulé observed that carbon often has four other atoms bonded to it. Methane, for example, has one carbon atom and four hydrogen atoms.[78] This concept eventually became known as valency, where different elements bond with different numbers of atoms.[79]
In 1862, the French geologist Alexandre-Émile Béguyer de Chancourtois published an early form of the periodic table, which he called the telluric helix or screw. He was the first person to notice the periodicity of the elements. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois showed that elements with similar properties seemed to occur at regular intervals. His chart included some ions and compounds in addition to elements. His paper also used geological rather than chemical terms and did not include a diagram. As a result, it received little attention until the work of Dmitri Mendeleev.[80]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/7/70/Periodic_table_Meyer_1864.png/300px-Periodic_table_Meyer_1864.png)
In 1864, Julius Lothar Meyer, a German chemist, published a table with 28 elements. Realizing that an arrangement according to atomic weight did not exactly fit the observed periodicity in chemical properties he gave valency priority over minor differences in atomic weight. A missing element between Si and Sn was predicted with atomic weight 73 and valency 4.[81] Concurrently, English chemist William Odling published an arrangement of 57 elements, ordered on the basis of their atomic weights. With some irregularities and gaps, he noticed what appeared to be a periodicity of atomic weights among the elements and that this accorded with "their usually received groupings".[82] Odling alluded to the idea of a periodic law but did not pursue it.[83] He subsequently proposed (in 1870) a valence-based classification of the elements.[84]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/e/e5/Newlands_periodiska_system_1866.png/300px-Newlands_periodiska_system_1866.png)
English chemist John Newlands produced a series of papers from 1863 to 1866 noting that when the elements were listed in order of increasing atomic weight, similar physical and chemical properties recurred at intervals of eight. He likened such periodicity to the octaves of music.[85][86] This so termed Law of Octaves was ridiculed by Newlands' contemporaries, and the Chemical Society refused to publish his work.[87] Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as germanium.[88] The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.[89]
In 1867, Gustavus Hinrichs, a Danish born academic chemist based in America, published a spiral periodic system based on atomic spectra and weights, and chemical similarities. His work was regarded as idiosyncratic, ostentatious and labyrinthine and this may have militated against its recognition and acceptance.[90][91]
Mendeleev's table
![](http://upload.wikimedia.org/wikipedia/commons/thumb/4/4d/Mendeleev%27s_1869_periodic_table.svg/250px-Mendeleev%27s_1869_periodic_table.svg.png)
Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer independently published their periodic tables in 1869 and 1870, respectively.[92] Mendeleev's table, dated March 1 [O.S. February 17] 1869,[93] was his first published version. That of Meyer was an expanded version of his (Meyer's) table of 1864.[94] They both constructed their tables by listing the elements in rows or columns in order of atomic weight and starting a new row or column when the characteristics of the elements began to repeat.[95]
The recognition and acceptance afforded to Mendeleev's table came from two decisions he made. The first was to leave gaps in the table when it seemed that the corresponding element had not yet been discovered.[96] Mendeleev was not the first chemist to do so, but he was the first to be recognized as using the trends in his periodic table to predict the properties of those missing elements, such as gallium and germanium.[97] The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as tellurium and iodine, to better classify them into chemical families.
Mendeleev published in 1869, using atomic weight to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough to allow Mendeleev to accurately predict the properties of missing elements.
Mendeleev took the unusual step of naming missing elements using the Sanskrit numerals eka (1), dvi (2), and tri (3) to indicate that the element in question was one, two, or three rows removed from a lighter congener. It has been suggested that Mendeleev, in doing so, was paying homage to ancient Sanskrit grammarians, in particular Pāṇini, who devised a periodic alphabet for the language.[98]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/c/c8/Henry_Moseley_%281887-1915%29.jpg/220px-Henry_Moseley_%281887-1915%29.jpg)
Following the discovery of the atomic nucleus by Ernest Rutherford in 1911, it was proposed that the integer count of the nuclear charge is identical to the sequential place of each element in the periodic table. In 1913, English physicist Henry Moseley using X-ray spectroscopy confirmed this proposal experimentally. Moseley determined the value of the nuclear charge of each element and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge.[99] Nuclear charge is identical to proton count and determines the value of the atomic number (Z) of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley predicted, in 1913, that the only elements still missing between aluminium (Z = 13) and gold (Z = 79) were Z = 43, 61, 72, and 75, all of which were later discovered. The atomic number is the absolute definition of an element and gives a factual basis for the ordering of the periodic table.[100]
Second version and further development
![](http://upload.wikimedia.org/wikipedia/commons/thumb/a/aa/Periodic_table_by_Mendeleev%2C_1871.svg/500px-Periodic_table_by_Mendeleev%2C_1871.svg.png)
![](http://upload.wikimedia.org/wikipedia/commons/thumb/c/c9/ShortPT20b.png/500px-ShortPT20b.png)
In 1871, Mendeleev published his periodic table in a new form, with groups of similar elements arranged in columns rather than in rows, and those columns numbered I to VIII corresponding with the element's oxidation state. He also gave detailed predictions for the properties of elements he had earlier noted were missing, but should exist.[101] These gaps were subsequently filled as chemists discovered additional naturally occurring elements.[102] It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939, but it was technically only the last element to be discovered in nature as opposed to by synthesis.[103] Plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971.[104]
The popular[105] periodic table layout, also known as the common or standard form (as shown at various other points in this article), is attributable to Horace Groves Deming. In 1923, Deming, an American chemist, published short (Mendeleev style) and medium (18-column) form periodic tables.[106][n 6] Merck and Company prepared a handout form of Deming's 18-column medium table, in 1928, which was widely circulated in American schools. By the 1930s Deming's table was appearing in handbooks and encyclopedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company.[107][108][109]
With the development of modern quantum mechanical theories of electron configurations within atoms, it became apparent that each period (row) in the table corresponded to the filling of a quantum shell of electrons. Larger atoms have more electron sub-shells, so later tables have required progressively longer periods.[110]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/4/47/Glenn_Seaborg_-_1964.jpg/150px-Glenn_Seaborg_-_1964.jpg)
In 1945, Glenn Seaborg, an American scientist, made the suggestion that the actinide elements, like the lanthanides, were filling an f sub-level. Before this time the actinides were thought to be forming a fourth d-block row. Seaborg's colleagues advised him not to publish such a radical suggestion as it would most likely ruin his career. As Seaborg considered he did not then have a career to bring into disrepute, he published anyway. Seaborg's suggestion was found to be correct and he subsequently went on to win the 1951 Nobel Prize in chemistry for his work in synthesizing actinide elements.[111][112][n 7]
Although minute quantities of some transuranic elements occur naturally,[3] they were all first discovered in laboratories. Their production has expanded the periodic table significantly, the first of these being neptunium, synthesized in 1939.[113] Because many of the transuranic elements are highly unstable and decay quickly, they are challenging to detect and characterize when produced. There have been controversies concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights.[114] In 2010, a joint Russia–US collaboration at Dubna, Moscow Oblast, Russia, claimed to have synthesized six atoms of tennessine (element 117), making it the most recently claimed discovery. It, along with nihonium (element 113), moscovium (element 115), and oganesson (element 118), are the four most recently named elements, whose names all became official on 28 November 2016.[115]
Different periodic tables
The long- or 32-column table
![](http://upload.wikimedia.org/wikipedia/commons/thumb/6/67/32-column_periodic_table-a.png/650px-32-column_periodic_table-a.png)
The modern periodic table is sometimes expanded into its long or 32-column form by reinstating the footnoted f-block elements into their natural position between the s- and d-blocks, as proposed by Alfred Werner.[116] Unlike the 18-column form this arrangement results in "no interruptions in the sequence of increasing atomic numbers".[117] The relationship of the f-block to the other blocks of the periodic table also becomes easier to see.[118] Jensen advocates a form of table with 32 columns on the grounds that the lanthanides and actinides are otherwise relegated in the minds of students as dull, unimportant elements that can be quarantined and ignored.[119] Despite these advantages the 32-column form is generally avoided by editors on account of its undue rectangular ratio compared to a book page ratio,[120] and the familiarity of chemists with the modern form, as introduced by Seaborg.[121]
Template:Periodic table (32 columns, detailed cells)
Tables with different structures
Within 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table.[119][122][123] As well as numerous rectangular variations, other periodic table formats have been shaped, for example,[n 8] like a circle, cube, cylinder, building, spiral, lemniscate,[124] octagonal prism, pyramid, sphere, or triangle. Such alternatives are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables.[123]
![](http://upload.wikimedia.org/wikipedia/commons/thumb/c/ce/Elementspiral_%28polyatomic%29.svg/220px-Elementspiral_%28polyatomic%29.svg.png)
A popular[125] alternative structure is that of Otto Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the centre and the transition metals, lanthanides, and actinides occupying peninsulas.[126]
Most periodic tables are two-dimensional;[3] three-dimensional tables are known to as far back as at least 1862 (pre-dating Mendeleev's two-dimensional table of 1869). More recent examples include Courtines' Periodic Classification (1925),[127] Wringley's Lamina System (1949),[128] Giguère's Periodic helix (1965)[129] and Dufour's Periodic Tree (1996).[130] Going one further, Stowe's Physicist's Periodic Table (1989)[131] has been described as being four-dimensional (having three spatial dimensions and one colour dimension).[132]
The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum.[133] Towards the chemistry end of the continuum can be found, as an example, Rayner-Canham's "unruly"[134] Inorganic Chemist's Periodic Table (2002),[135] which emphasizes trends and patterns, and unusual chemical relationships and properties. Near the physics end of the continuum is Janet's Left-Step Periodic Table (1928). This has a structure that shows a closer connection to the order of electron-shell filling and, by association, quantum mechanics.[136] A somewhat similar approach has been taken by Alper,[137] albeit criticized by Eric Scerri as disregarding the need to display chemical and physical periodicity.[138] Somewhere in the middle of the continuum is the ubiquitous common or standard form of periodic table. This is regarded as better expressing empirical trends in physical state, electrical and thermal conductivity, and oxidation numbers, and other properties easily inferred from traditional techniques of the chemical laboratory.[139] Its popularity is thought to be a result of this layout having a good balance of features in terms of ease of construction and size, and its depiction of atomic order and periodic trends.[83][140]
f1 | f2 | f3 | f4 | f5 | f6 | f7 | f8 | f9 | f10 | f11 | f12 | f13 | f14 | d1 | d2 | d3 | d4 | d5 | d6 | d7 | d8 | d9 | d10 | p1 | p2 | p3 | p4 | p5 | p6 | s1 | s2 | |||
1s | H | He | ||||||||||||||||||||||||||||||||
2s | Li | Be | ||||||||||||||||||||||||||||||||
2p 3s | B | C | N | O | F | Ne | Na | Mg | ||||||||||||||||||||||||||
3p 4s | Al | Si | P | S | Cl | Ar | K | Ca | ||||||||||||||||||||||||||
3d 4p 5s | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | Rb | Sr | ||||||||||||||||
4d 5p 6s | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | Cs | Ba | ||||||||||||||||
4f 5d 6p 7s | 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 | ||
5f 6d 7p 8s | Ac | Th | Pa | U | Np | Pu | Am | Cm | Bk | Cf | Es | Fm | Md | No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og | Uue | Ubn | ||
f-block | d-block | p-block | s-block |
Open questions and controversies
Placement of hydrogen and helium
Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium (1s22s1) and beryllium (1s22s2).[141] While such a placement is common for hydrogen, it is rarely used for helium outside of the context of electron configurations: When the noble gases (then called "inert gases") were first discovered around 1900, they were known as "group 0", reflecting no chemical reactivity of these elements known at that point, and helium was placed on the top of that group, as it did share the extreme chemical inertness seen throughout the group. As the group changed its formal number, many authors continued to assign helium directly above neon, in group 18; one of the examples of such placing is the current IUPAC table.[142]
The position of hydrogen in group 1 is reasonably well settled. Its usual oxidation state is +1 as is the case for its heavier alkali metal congeners. Like lithium, it has a significant covalent chemistry.[143][144] It can stand in for alkali metals in typical alkali metal structures.[145] It is capable of forming alloy-like hydrides, featuring metallic bonding, with some transition metals.[146]
Nevertheless, it is sometimes placed elsewhere. A common alternative is at the top of group 17[138] given hydrogen's strictly univalent and largely non-metallic chemistry, and the strictly univalent and non-metallic chemistry of fluorine (the element otherwise at the top of group 17). Sometimes, to show hydrogen has properties corresponding to both those of the alkali metals and the halogens, it is shown at the top of the two columns simultaneously.[147] Another suggestion is above carbon in group 14: placed that way, it fits well into the trends of increasing ionization potential values and electron affinity values, and is not too far from the electronegativity trend, even though hydrogen cannot show the tetravalence characteristic of the heavier group 14 elements.[148] Finally, hydrogen is sometimes placed separately from any group; this is based on its general properties being regarded as sufficiently different from those of the elements in any other group.
The other period 1 element, helium, is most often placed in group 18 with the other noble gases, as its extraordinary inertness is extremely close to that of the other light noble gases neon and argon.[149] Nevertheless, it is occasionally placed separately from any group as well.[150] The property that distinguishes helium from the rest of the noble gases is that in its closed electron shell, helium has only two electrons in the outermost electron orbital, while the rest of the noble gases have eight. Some authors, such as Henry Bent (the eponym of Bent's rule), Wojciech Grochala, and Felice Grandinetti, have argued that helium would be correctly placed in group 2, over beryllium; Charles Janet's left-step table also contains this assignment. The normalized ionization potentials and electron affinities show better trends with helium in group 2 than in group 18; helium is expected to be slightly more reactive than neon (which breaks the general trend of reactivity in the noble gases, where the heavier ones are more reactive); predicted helium compounds often lack neon analogues even theoretically, but sometimes have beryllium analogues; and helium over beryllium better follows the trend of first-row anomalies in the table (s >> p > d > f).[151][152][153]
Group 3 and its elements in periods 6 and 7
Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely settled. They are commonly lanthanum and actinium, and less often lutetium and lawrencium. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the f block elements start and end.[154][n 9] It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks".[155] A third common variant shows the two positions below yttrium as being occupied by the lanthanides and the actinides. A fourth variant shows group 3 bifurcating after Sc-Y, into an La-Ac branch, and an Lu-Lr branch.[34]
Since 1921, many chemical and physical arguments have been made in support of lutetium and lawrencium[156][157] but the majority of authors seem either unconvinced by them or unaware of them.[158][159] Most working chemists are not aware there is any controversy.[159] In December 2015 an IUPAC project was established to make a recommendation on the matter, considering only the first two alternatives as possibilities.[160]
Lanthanum and actinium
![]() La and Ac below Y |
Lanthanum and actinium are commonly depicted as the remaining group 3 members.[161][n 10] It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the ground-state electron configurations of the elements and the notion of the differentiating electron. The ground-state configurations of caesium, barium and lanthanum are [Xe]6s1, [Xe]6s2 and [Xe]5d16s2. Lanthanum thus emerges with a 5d differentiating electron and on these grounds it was considered to be "in group 3 as the first member of the d-block for period 6".[25] A superficially consistent set of electron configurations is then seen in group 3: scandium [Ar]3d14s2, yttrium [Kr]4d15s2 and lanthanum [Xe]5d16s2. Still in period 6, ytterbium was assigned an electron configuration of [Xe]4f135d16s2 and lutetium [Xe]4f145d16s2, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6".[25] Later spectroscopic work found that the electron configuration of ytterbium was in fact [Xe]4f146s2. This meant that ytterbium and lutetium—the latter with [Xe]4f145d16s2—both had 14 f-electrons, "resulting in a d- rather than an f- differentiating electron" for lutetium and making it an "equally valid candidate" with [Xe]5d16s2 lanthanum, for the group 3 periodic table position below yttrium.[25] Lanthanum has the advantage of incumbency since the 5d1 electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium[162] (though similar logic would also lead to thorium getting the 6d2 position, having incumbency over rutherfordium).[citation needed]
Ground-state gas-phase configurations consider only isolated atoms as opposed to bonding atoms in compounds (the latter being more relevant for chemistry), which often show different configurations.[163] Moreover, the lowest levels of two different configurations often are separated by only very small energies, that are minuscule compared to the spreading of J-levels of each configuration (e.g. terbium, where the 285 cm−1 difference between [Xe]4f85d16s2 and the ground state [Xe]4f95s2 is much less than 1% of this spreading), making which configuration happens to be the ground state chemically quite irrelevant.[164] This form also creates an inconsistency in the treatment of thorium, which has no f-electrons in the ground-state (being [Rn]6d27s2), similar to actinium as [Rn]6d17s2; yet this form places thorium in the f-block but not actinium.[165] Considering only ground-state gas-phase configurations, thorium [Rn]6d27s2 by itself is just as good a homologue to zirconium [Kr]4d25s2 as lanthanum [Xe]5d16s2 is to scandium [Ar]3d14s2;[164] yet thorium is invariably placed in the f-block. Thorium thus demonstrates that the possession of an f electron in the ground-state gas-phase configuration of an element is not necessary for it to belong to the f-block. Additionally, this form necessitates a split d-block if expanded to a 32-column periodic table.[166]
In terms of chemical behaviour,[167] and trends going down group 3 (if Sc-Y-La is chosen) for properties such as melting point, electronegativity and ionic radius,[168][169] scandium, yttrium, lanthanum and actinium are similar to their group 1–2 counterparts, but at variance with the other groups in the early d-block. In this variant, the number of f electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block.[170] For example, the f-electron counts for the trivalent ions of the first three f-block elements are Ce 1, Pr 2 and Nd 3.[171] However, outside the lanthanides there does not exist a typical oxidation state across any period of a block,[41] and the reason for this singular behaviour of the lanthanides in fact has very little to do with the electron configurations of the elements concerned.[172]
Lutetium and lawrencium
![]() Lu and Lr below Y |
In other tables, lutetium and lawrencium are the remaining group 3 members.[n 11] Early techniques for chemically separating scandium, yttrium and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".[25] Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s.[n 12] Several physicists in the 1950s and '60s favoured lutetium, in light of a comparison of several of its physical properties with those of lanthanum.[25] Among the prominent adherents of this form have been Lev Landau and Evgeny Lifshitz, who wrote in their Course of Theoretical Physics (1958):[173]
In books of chemistry, lutetium is usually placed in the rare-earth elements. This, however, is incorrect, since the 4f shell is complete in lutetium; it must therefore be placed in the platinum group [which they considered to be La+Lu–Pt]...
— Landau and Lifshitz
This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. It has been argued that this is not a valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block.[174] Karl Gschneidner, analysing the melting points of the lanthanides in a 1971 article, reached the conclusion that it was likely that 4f, 5d, 6s, and 6p electrons were all involved in the bonding of lanthanide metals except for lutetium, where 4f electrons were not found to be involved.[175] The fact that lanthanum was demonstrated to be a 4f-band metal (with about 0.17 electrons per atom in fcc lanthanum, which is metastable at standard conditions)[42] whereas the 4f shell appears to have no influence on the metallic properties of lutetium, has been used as an argument to place lutetium in group 3 instead of lanthanum.[176] The 4f occupancy in solid lanthanum may explain some of its properties, such as its low melting point[175] and superconductivity.[176] Moreover, lanthanum has low-lying empty f orbitals which lutetium lacks,[165] and it can have the 4f orbitals contributing measurably to the bonding, for example in lanthanum(III) fluoride (LaF3), which is not known for lutetium. While this contribution is small, it is greater for lanthanum than for any other lanthanide;[177] the order of involvement of 4f in lanthanum is similar to that of 5f in thorium.[164] And while the trivalent lanthanides Pr3+ through Yb3+ show characteristic narrow bands with their positions almost completely independent on the ligands, the following 5d elements (along with the 3d and 4d elements) behave significantly differently; while both types of elements show electron-transfer bands, ligand field theory becomes important for these d elements.[172]
Scandium, yttrium, and lutetium show a more consistent set of electron configurations matching the global trend on the periodic table: the 5d metals then all add a closed 4f14 shell. For example, the shift from yttrium [Kr]4d15s2 to lutetium [Xe]4f145d16s2 exactly parallels that from zirconium [Kr]4d25s2 to hafnium [Xe]4f145d26s2.[25] The inclusion of lutetium rather than lanthanum also homogenises the 5d transition series: trends in atomic size, coordination number, and relative abundance of metal–oxygen bonds all reveal that lutetium is closer than lanthanum to the behaviour of the uncontroversial 5d metals hafnium through mercury.[178] The same is true considering conduction band structures of the elements: lutetium has a transition-metal-like conduction band structure, but lanthanum does not.[179]
As for lawrencium, its gas phase ground-state atomic electron configuration was confirmed in 2015 as [Rn]5f147s27p1. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for nihonium with its predicted configuration of [Rn]5f146d107s27p1.[180] However, it is expected that in the condensed phase and in chemical environments lawrencium has the expected 6d occupancy, and simple modelling studies suggest it will behave like a lanthanide,[181] in particular being a homologue of lutetium. Lawrencium's return to +3 as the only stable oxidation state and being predicted to form a trivalent metal is distinct from the behaviour of the other late actinides fermium, mendelevium, and nobelium, which have a tendency towards forming lower oxidation states and form (or are predicted to form) divalent metals; it also makes an exception to the actinide contraction generally being larger than the analogous lanthanide contraction at the end of both series.[182]
While scandium, yttrium and lutetium (and lawrencium, so far as its chemistry is known) do often behave like trivalent versions of the group 1–2 metals, being hard class-A cations mostly restricted to the group oxidation state, they are not the only elements in the d-block or f-block that do so. The early transition metals zirconium and hafnium in group 4, as well as niobium and tantalum in group 5, also display such behaviour, as does the actinide thorium. (The heavy group 4 elements and thorium are tetravalent; the heavy group 5 elements are pentavalent.)[183][184] The physical properties of the group 3 elements are affected by the presence of a d electron, which forms more localised bonds within the metals than the p electrons in the similar group 13 metals;[66] exactly the same situation is found comparing group 4 to group 14.[65] Trends going down group 3 (if Sc-Y-Lu is chosen) for properties such as melting point, electronegativity and ionic radius, are similar to those found among their group 4–8 counterparts.[25] In this variant, the number of f electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.[25]
Lanthanides and actinides
![]() Markers below Y |
A few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). This variant, which is stated in the 2005 Red Book to be the IUPAC-agreed version as of 2005 (a number of later versions exist, and the last update is from 1 December 2018),[185][n 13] emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide f block (there can only be 14 elements in any row of the f block).[n 14] However, this similarity does not extend to the 15 actinide elements (Ac–Lr), which show a much wider variety in their chemistries.[164] This form moreover reduces the f-block to a degenerate branch of group 3 of the d-block; it dates back to the 1920s when the lanthanides were thought to have their f electrons as core electrons, which is now known to be false. It is also false for the actinides, many of which show stable oxidation states above +3.[187]
La-Ac and Lu-Lr
![](http://upload.wikimedia.org/wikipedia/commons/thumb/2/2d/Silberberg_style_periodic_table.png/220px-Silberberg_style_periodic_table.png)
In this variant, group 3 bifurcates after Sc-Y into a La-Ac branch, and a Lu-Lr branch. This arrangement is consistent with the hypothesis that arguments in favour of either Sc-Y-La-Ac or Sc-Y-Lu-Lr based on chemical and physical data are inconclusive.[188] As noted, trends going down Sc-Y-La-Ac match trends in groups 1−2[189] whereas trends going down Sc-Y-Lu-Lr better match trends in groups 4−10.[25]
The bifurcation of group 3 is a throwback to the Mendeleev eight column-form in which seven of the main groups each have two subgroups. Tables featuring a bifurcated group 3 have been periodically proposed since that time.[n 15]
Groups included in the transition metals
The definition of a transition metal, as given by IUPAC in the Gold Book, is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell.[190] By this definition all of the elements in groups 3–11 are transition metals. The IUPAC definition therefore excludes group 12, comprising zinc, cadmium and mercury, from the transition metals category. However, the 2005 IUPAC nomenclature as codified in the Red Book gives both the group 3–11 and group 3–12 definitions of the transition metals as alternatives.
Some chemists treat the categories "d-block elements" and "transition metals" interchangeably, thereby including groups 3–12 among the transition metals. In this instance the group 12 elements are treated as a special case of transition metal in which the d electrons are not ordinarily given up for chemical bonding (they can sometimes contribute to the valence bonding orbitals even so, as in zinc fluoride).[191] The 2007 report of mercury(IV) fluoride (HgF4), a compound in which mercury would use its d electrons for bonding, has prompted some commentators to suggest that mercury can be regarded as a transition metal.[192] Other commentators, such as Jensen,[193] have argued that the formation of a compound like HgF4 can occur only under highly abnormal conditions; indeed, its existence is currently disputed. As such, mercury could not be regarded as a transition metal by any reasonable interpretation of the ordinary meaning of the term.[193]
Still other chemists further exclude the group 3 elements from the definition of a transition metal. They do so on the basis that the group 3 elements do not form any ions having a partially occupied d shell and do not therefore exhibit properties characteristic of transition metal chemistry.[194] In this case, only groups 4–11 are regarded as transition metals. This categorisation is however not one of the alternatives considered by IUPAC. Though the group 3 elements show few of the characteristic chemical properties of the transition metals, the same is true of the heavy members of groups 4 and 5, which also are mostly restricted to the group oxidation state in their chemistry. Moreover, the group 3 elements show characteristic physical properties of transition metals (on account of the presence in each atom of a single d electron).[66]
Elements with unknown chemical properties
Although all elements up to oganesson have been discovered, of the elements above hassium (element 108), only copernicium (element 112), nihonium (element 113), and flerovium (element 114) have known chemical properties, and conclusive categorisation at present has not been reached.[59] Some of these may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, copernicium and flerovium have been predicted to possibly exhibit some noble-gas-like properties, even though neither is placed in group 18 with the other noble gases.[59][195] The current experimental evidence still leaves open the question of whether copernicium and flerovium behave more like metals or noble gases.[59][196] At the same time, oganesson (element 118) is expected to be a solid semiconductor at standard conditions, despite being in group 18.[197]
Further periodic table extensions
It is unclear whether new elements will continue the pattern of the current periodic table as period 8, or require further adaptations or adjustments. Seaborg expected the eighth period to follow the previously established pattern exactly, so that it would include a two-element s-block for elements 119 and 120, a new g-block for the next 18 elements, and 30 additional elements continuing the current f-, d-, and p-blocks, culminating in element 168, the next noble gas.[199] More recently, physicists such as Pekka Pyykkö have theorized that these additional elements do not exactly follow the Madelung rule, which predicts how electron shells are filled and thus affects the appearance of the present periodic table. There are currently several competing theoretical models for the placement of the elements of atomic number less than or equal to 172. In all of these it is element 172, rather than element 168, that emerges as the next noble gas after oganesson, although these must be regarded as speculative as no complete calculations have been done beyond element 123.[200][201]
Element with the highest possible atomic number
The number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than circa 256 (which would equate to between elements 99 and 100 in modern-day terms) did not exist.[202] A higher, more recent estimate is that the periodic table may end soon after the island of stability,[203] whose centre is predicted to lie between element 110 and element 126, as the extension of the periodic and nuclide tables is restricted by proton and neutron drip lines as well as decreasing stability towards spontaneous fission.[204][205] Other predictions of an end to the periodic table include at element 128 by John Emsley,[3] at element 137 by Richard Feynman,[206] at element 146 by Yogendra Gambhir,[207] and at element 155 by Albert Khazan.[3][n 16]
Bohr model
The Bohr model exhibits difficulty for atoms with atomic number greater than 137, as any element with an atomic number greater than 137 would require 1s electrons to be travelling faster than c, the speed of light.[208] Hence the non-relativistic Bohr model is inaccurate when applied to such an element.
Relativistic Dirac equation
The relativistic Dirac equation has problems for elements with more than 137 protons. For such elements, the wave function of the Dirac ground state is oscillatory rather than bound, and there is no gap between the positive and negative energy spectra, as in the Klein paradox.[209] More accurate calculations taking into account the effects of the finite size of the nucleus indicate that the binding energy first exceeds the limit for elements with more than 173 protons. For heavier elements, if the innermost orbital (1s) is not filled, the electric field of the nucleus will pull an electron out of the vacuum, resulting in the spontaneous emission of a positron.[210] This does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.[206]
Optimal form
The many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table.[211] The answer to this question is thought to depend on whether the chemical periodicity seen to occur among the elements has an underlying truth, effectively hard-wired into the universe, or if any such periodicity is instead the product of subjective human interpretation, contingent upon the circumstances, beliefs and predilections of human observers. An objective basis for chemical periodicity would settle the questions about the location of hydrogen and helium, and the composition of group 3. Such an underlying truth, if it exists, is thought to have not yet been discovered. In its absence, the many different forms of periodic table can be regarded as variations on the theme of chemical periodicity, each of which explores and emphasizes different aspects, properties, perspectives and relationships of and among the elements.[n 17]
Other
In celebration of the periodic table's 150th anniversary, the United Nations declared the year 2019 as the International Year of the Periodic Table, celebrating "one of the most significant achievements in science".[214]
See also
- Abundance of the chemical elements
- Atomic electron configuration table
- Element collecting
- List of chemical elements
- List of periodic table-related articles
- Names for sets of chemical elements
- Standard model
- Table of nuclides
- Template:Spectral lines of the elements
- The Mystery of Matter: Search for the Elements (PBS film)
- Timeline of chemical element discoveries
Notes
- ^ The elements discovered initially by synthesis and later in nature are technetium (Z = 43), promethium (61), astatine (85), neptunium (93), and plutonium (94).
- ^ An element zero (i.e. a substance composed purely of neutrons), is included in a few alternate presentations, for example, in the Chemical Galaxy.[6]
- ^ The noble gases, astatine, francium, and all elements heavier than americium were left out as there is no data for them.
- ^ While fluorine is the most electronegative of the elements under the Pauling scale, neon is the most electronegative element under other scales, such as the Allen scale.
- ^ While Lr is thought to have a p rather than d electron in its ground-state electron configuration, and would therefore be expected to be a volatile metal capable of forming a +1 cation in solution like thallium, no evidence of either of these properties has been able to be obtained despite experimental attempts to do so.[70] It was originally expected to have a d electron in its electron configuration[70] and this may still be the case for metallic lawrencium, whereas gas phase atomic lawrencium is very likely thought to have a p electron.[71]
- ^ An antecedent of Deming's 18-column table may be seen in Adams' 16-column Periodic Table of 1911. Adams omits the rare earths and the "radioactive elements" (i.e. the actinides) from the main body of his table and instead shows them as being "careted in only to save space" (rare earths between Ba and eka-Yt; radioactive elements between eka-Te and eka-I). See: Elliot Q. A. (1911). "A modification of the periodic table". Journal of the American Chemical Society. 33(5): 684–688 (687).
- ^ A second extra-long periodic table row, to accommodate known and undiscovered elements with an atomic weight greater than bismuth (thorium, protactinium and uranium, for example), had been postulated as far back as 1892. Most investigators considered that these elements were analogues of the third series transition elements, hafnium, tantalum and tungsten. The existence of a second inner transition series, in the form of the actinides, was not accepted until similarities with the electron structures of the lanthanides had been established. See: van Spronsen, J. W. (1969). The periodic system of chemical elements. Amsterdam: Elsevier. pp. 315–316, ISBN 0-444-40776-6.
- ^ See The Internet database of periodic tables for depictions of these kinds of variants.
- ^ The detachment of the lanthanides from the main body of the periodic table has been attributed to the Czech chemist Bohuslav Brauner who, in 1902, allocated all of them ("Ce etc.") to one position in group 4, below zirconium. This arrangement was referred to as the "asteroid hypothesis", in analogy to asteroids occupying a single orbit in the solar system. Before this time the lanthanides were generally (and unsuccessfully) placed throughout groups I to VIII of the older 8-column form of periodic table. Although predecessors of Brauner's 1902 arrangement are recorded from as early as 1895, he is known to have referred to the "chemistry of asteroids" in an 1881 letter to Mendeleev. Other authors assigned all of the lanthanides to either group 3, groups 3 and 4, or groups 2, 3 and 4. In 1922 Niels Bohr continued the detachment process by locating the lanthanides between the s- and d-blocks. In 1949 Glenn T. Seaborg (re)introduced the form of periodic table that is popular today, in which the lanthanides and actinides appear as footnotes. Seaborg first published his table in a classified report dated 1944. It was published again by him in 1945 in Chemical and Engineering News, and in the years up to 1949 several authors commented on, and generally agreed with, Seaborg's proposal. In that year he noted that the best method for presenting the actinides seemed to be by positioning them below, and as analogues of, the lanthanides. See: Thyssen P. and Binnemans K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 41. Amsterdam: Elsevier, pp. 1–94; Seaborg G. T. (1994). Origin of the Actinide Concept'. In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 18. Amsterdam: Elsevier, pp. 1–27.
- ^ For examples of this table see Atkins et al. (2006). Shriver & Atkins Inorganic Chemistry (4th ed.). Oxford: Oxford University Press • Myers et al. (2004). Holt Chemistry. Orlando: Holt, Rinehart & Winston • Chang R. (2000). Essential Chemistry (2nd ed.). Boston: McGraw-Hill
- ^ For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G. & Overton T. (2013). Descriptive Inorganic Chemistry (6th ed.). New York: W. H. Freeman and Company • Brown et al. (2009). Chemistry: The Central Science (11th ed.). Upper Saddle River, New Jersey: Pearson Education • Moore et al. (1978). Chemistry. Tokyo: McGraw-Hill Kogakusha
- ^ The phenomenon of different separation groups is caused by increasing basicity with increasing radius, and does not constitute a fundamental reason to show Lu, rather than La, below Y. Thus, among the Group 2 alkaline earth metals, Mg (less basic) belongs in the "soluble group" and Ca, Sr and Ba (more basic) occur in the "ammonium carbonate group". Nevertheless, Mg, Ca, Sr and Ba are routinely collocated in Group 2 of the periodic table. See: Moeller et al. (1989). Chemistry with Inorganic Qualitative Analysis (3rd ed.). SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.
- ^ Notwithstanding, an IUPAC member subsequently wrote that, "IUPAC has not approved any specific form of the periodic table, and an IUPAC-approved form does not exist, though even members of IUPAC themselves have published diagrams titled “IUPAC Periodic Table of the Elements". However, the only specific recommendation IUPAC has made concerning the periodic table covers the Group numbering of 1–18."[186]
- ^ For examples of the group 3 = Ln and An table see Housecroft C. E. & Sharpe A. G. (2008). Inorganic Chemistry (3rd ed.). Harlow: Pearson Education • Halliday et al. (2005). Fundamentals of Physics (7th ed.). Hoboken, NewJersey: John Wiley & Sons • Nebergall et al. (1980). General Chemistry (6th ed.). Lexington: D. C. Heath and Company
- ^
- 1922 Bohr's system, with bifurcations at Na, Mg, and Y
- 1939 Foster LS, [https://www.meta-synthesis.com/webbook/35_pt/pt_database.php?PT_id=1056 "Why not modernise textbooks also? I. The periodic table", Journal of Chemical Education, vol. 16, no. 9. In group 3 the box under Y is “La {58–70}* Lu”, where * = the rare earths.
- 1947 Stedman's planar arrangement of his conic system
- 1952 Coryell, shows a step pyramid style table, with solid and dashed tie lines to show primary and secondary relationships. Just two elements are shown by him as having two solid tie lines: yttrium, to La-Ac and to Lu-Ac; and silicon, to Ti-Zr-Hf and to Ge-Sn-Pb.
- 1964 Sanderson RT, "A rational periodic table", Journal of Chemical Education, vol. 41, no. 4, pp. 187–189
- 1974 Mazurs EG, Graphic representations of the periodic system during one hundred years, The University of Alabama Press, Alabama, p. 77
- 2020 Vernon RE, "Organising the metals and nonmetals", Foundations of Chemistry, open access, see Electronic supplementary material
- ^ Karol (2002, p. 63) contends that gravitational effects would become significant when atomic numbers become astronomically large, thereby overcoming other super-massive nuclei instability phenomena, and that neutron stars (with atomic numbers on the order of 1021) can arguably be regarded as representing the heaviest known elements in the universe. See: Karol P. J. (2002). "The Mendeleev–Seaborg periodic table: Through Z = 1138 and beyond". Journal of Chemical Education 79 (1): 60–63.
- ^ Scerri, one of the foremost authorities on the history of the periodic table,[212] whilst previously recognising the value of a plurality of periodic tables,[211] currently supports the concept of an optimal table.[213]
References
- ^ "Chemistry: Four elements added to periodic table". BBC News. 4 January 2016. Archived from the original on 4 January 2016.
- ^ St. Fleur, Nicholas (1 December 2016). "Four New Names Officially Added to the Periodic Table of Elements". New York Times. Archived from the original on 14 August 2017.
- ^ a b c d e f Emsley, J. (2011). Nature's Building Blocks: An A-Z Guide to the Elements (New ed.). New York, NY: Oxford University Press. ISBN 978-0-19-960563-7.
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- ^ Thyssen, P.; Binnemanns, K. (2011). "1: Accommodation of the rare earths in the periodic table: A historical analysis". In Gschneidner Jr., K. A.; Büzli, J-C. J.; Pecharsky, V. K. (eds.). Handbook on the Physics and Chemistry of Rare Earths. Vol. 41. Amsterdam: Elsevier. pp. 80–81. ISBN 978-0-444-53590-0.
- ^ Keeler, J.; Wothers, P. (2014). Chemical Structure and Reactivity: An Integrated Approach. Oxford: Oxford University. p. 259. ISBN 978-0-19-960413-5.
- ^ Scerri, E. (2012). "Mendeleev's Periodic Table Is Finally Completed and What To Do about Group 3?". Chemistry International. 34 (4). doi:10.1515/ci.2012.34.4.28. Archived from the original on 5 July 2017.
- ^ a b Castelvecchi, D. (8 April 2015). "Exotic atom struggles to find its place in the periodic table". Nature. doi:10.1038/nature.2015.17275. Archived from the original on 5 October 2015. Retrieved 20 September 2015.
- ^ "The constitution of group 3 of the periodic table". IUPAC. 2015. Archived from the original on 5 July 2016. Retrieved 30 July 2016.
- ^ Emsley, J. (2011). Nature's Building Blocks (new ed.). Oxford: Oxford University. p. 651. ISBN 978-0-19-960563-7.
- ^ Trifonov, D. N. (1970). Rare-earth elements and their position in the periodic system (translated from Russian). New Delhi: Indian National Scientific Documentation Centre. pp. 201–202.
- ^ Scerri, Eric (2016). "The Changing Views of a Philosopher of Chemistry on the Question of Reduction". In Scerri, Eric; Fisher, Grant (eds.). Essays in the Philosophy of Chemistry. Oxford University Press.
- ^ a b c d Jørgensen, Christian K. (1988). Handbook on the Physics and Chemistry of Rare Earths. Vol. 11. pp. 197–292. doi:10.1016/S0168-1273(88)11007-6.
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- ^ Scerri, Eric (2009). "Which Elements Belong in Group 3?". Journal of Chemical Education. 86 (10): 1188. doi:10.1021/ed086p1188. Retrieved 16 May 2020.
- ^ Greenwood, N. N.; Harrington, T. J. (1973). The chemistry of the transition elements. Oxford: Clarendon Press. p. 50. ISBN 978-0-19-855435-6.
- ^ Aylward, G.; Findlay, T. (2008). SI chemical data (6th ed.). Milton, Queensland: John Wiley & Sons. ISBN 978-0-470-81638-7.
- ^ Wiberg, N. (2001). Inorganic Chemistry. San Diego: Academic Press. p. 119. ISBN 978-0-12-352651-9.
- ^ Wulfsberg, G. (2006). "Periodic table: Trends in the properties of the elements". Encyclopedia of Inorganic Chemistry. New York: John Wiley & Sons. p. 3. ISBN 978-0-470-86210-0.
- ^ Cotton, S. (2007). Lanthanide and Actinide Chemistry. Chichester: John Wiley & Sons. p. 150. ISBN 978-0-470-01006-8.
- ^ a b Jørgensen, Christian K. (1983). "5f Electrons in Transthorium Chemistry" (PDF). Radiochimica Acta. 32: 1–5. doi:10.1524/ract.1983.32.13.1. Retrieved 16 May 2020.
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It is high time that the idea of group 3 consisting of Sc, Y, La and Ac is abandoned
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Lesser omissions include ... the several different outdated versions of the periodic table. (That on the inside front cover is the current IUPAC-agreed version.)
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Bibliography
- Ball, P. (2002). The Ingredients: A Guided Tour of the Elements. Oxford: Oxford University Press. ISBN 978-0-19-284100-1.
- Chang, R. (2002). Chemistry (7th ed.). New York: McGraw-Hill Higher Education. ISBN 978-0-19-284100-1.
- Gray, T. (2009). The Elements: A Visual Exploration of Every Known Atom in the Universe. New York: Black Dog & Leventhal Publishers. ISBN 978-1-57912-814-2.
- Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford: Pergamon Press. ISBN 978-0-08-022057-4.
- Huheey, J. E.; Keiter, E. A.; Keiter, R. L. (1993). Principles of structure and reactivity (4th ed.). New York: Harper Collins College Publishers. ISBN 978-0-06-042995-9.
- Moore, J. T. (2003). Chemistry For Dummies. For Dummies (1st ed.). New York: Wiley Publications. ISBN 978-0-7645-5430-8.
- Scerri, E. (2007). The periodic table: Its story and its significance. Oxford: Oxford University Press. ISBN 978-0-19-530573-9.
- Scerri, E. (2011). The periodic table: A very short introduction. Oxford: Oxford University Press. ISBN 978-0-19-958249-5.
- Venable, F. P. (1896). The Development of the Periodic Law. Easton, Pennsylvania: Chemical Publishing Company. OCLC 776059614.
Further reading
- Calvo, Miguel (2019). Construyendo la Tabla Periódica. Zaragoza, Spain: Prames. p. 407. ISBN 978-84-8321-908-9.
- Emsley, J. (2011). "The Periodic Table". Nature's Building Blocks: An A–Z Guide to the Elements (New ed.). Oxford: Oxford University Press. pp. 634–651. ISBN 978-0-19-960563-7.
- Fontani, Marco; Costa, Mariagrazia; Orna, Mary Virginia (2007). The Lost Elements: The Periodic Table's Shadow Side. Oxford: Oxford University Press. p. 508. ISBN 978-0-19-938334-4.
- Mazurs, E. G. (1974). Graphical Representations of the Periodic System During One Hundred Years. Alabama: University of Alabama Press. ISBN 978-0-19-960563-7.
- Rouvray, D.H.; King, R. B., eds. (2004). The Periodic Table: Into the 21st Century. Proceedings of the 2nd International Conference on the Periodic Table, part 1, Kananaskis Guest Ranch, Alberta, 14–20 July 2003. Baldock, Hertfordshire: Research Studies Press. ISBN 978-0-86380-292-8.
- Rouvray, D.H.; King, R. B., eds. (2006). The Mathematics of the Periodic Table. Proceedings of the 2nd International Conference on the Periodic Table, part 2, Kananaskis Guest Ranch, Alberta, 14–20 July 2003. New York: Nova Science. ISBN 978-1-59454-259-6.
- Scerri, E (n.d.). "Books on the Elements and the Periodic Table" (PDF). Retrieved 9 July 2018.
- Scerri, E.; Restrepo, G, eds. (2018). Mendeleev to Oganesson: A Multidisciplinary Perspective on the Periodic Table. Proceedings of the 3rd International Conference on the Periodic Table, Cuzco, Peru 14–16 August 2012. Oxford: Oxford University Press. ISBN 978-0-86380-292-8.
- van Spronsen, J. W. (1969). The Periodic System of Chemical Elements: A History of the First Hundred Years. Amsterdam: Elsevier. ISBN 978-0-444-40776-4.
- Verde, M., ed. (1971). Atti del convegno Mendeleeviano: Periodicità e simmetrie nella struttura elementare della materia [Proceedings of the Mendeleevian conference: Periodicity and symmetry in the elementary structure of matter]. 1st International Conference on the Periodic Table, Torino-Roma, 15–21 September 1969. Torino: Accademia delle Scienze di Torino.
External links
- Periodic Table featured topic page on Science History Institute Digital Collections featuring select visual representations of the periodic table of the elements, with an emphasis on alternative layouts including circular, cylindrical, pyramidal, spiral, and triangular forms.
- IUPAC Periodic Table of the Elements
- Dynamic periodic table, with interactive layouts
- Eric Scerri, leading philosopher of science specializing in the history and philosophy of the periodic table
- The INTERNET Database of Periodic Tables
- Periodic table of endangered elements
- Periodic table of samples
- Periodic table of videos
- WebElements
1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
1 | H | He | ||||||||||||||||||||||||||||||
2 | Li | Be | B | C | N | O | F | Ne | ||||||||||||||||||||||||
3 | Na | Mg | Al | Si | P | S | Cl | Ar | ||||||||||||||||||||||||
4 | K | Ca | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | ||||||||||||||
5 | Rb | Sr | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | ||||||||||||||
6 | 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 |
7 | 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 | Nh | Fl | Mc | Lv | Ts | Og |
|