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{{Short description|Tabular arrangement of the chemical elements ordered by atomic number}}
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{{About|the table used in chemistry|other uses|The Periodic Table (disambiguation){{!}}The Periodic Table}} {{About|the table used in chemistry and physics|other uses|Periodic table (disambiguation)}}
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] (in periodic tables), and a traditional ]. The ] actually fits between ] and ]; it is usually shown at the foot of the table to save horizontal space.]]
The '''periodic table''' is a tabular display of the ]s, organized on the basis of their ]s and chemical properties. Elements are presented in increasing atomic number. The main body of the table is a 18 × 7 grid, and elements with the same number of ]s are kept together in ], such as the ]s and the ]es. Due to this, there are gaps that form four distinct rectangular areas or ]. The ] is not included in the main table, but rather is usually floated below, as an inline f-block would make the table impractically wide. Using ], the periodic table can help predict the properties of various elements and the relations between properties. As a result, it provides a useful framework for analyzing chemical behavior, and is widely used in chemistry and other sciences.


{{Sidebar periodic table}}
Although precursors exist, the current table is generally credited to ], who developed it in 1869 to illustrate periodic trends in the properties of the then-known elements; the layout has been refined and extended as ] have been discovered and new theoretical models developed to explain chemical behavior. Mendeleev's presentation also predicted some properties of ] expected to fill gaps in his arrangement; most of these predictions were proved correct when those elements were discovered and found to have properties close to the predictions.
The '''periodic table''', also known as the '''periodic table of the elements''', is an ordered arrangement of the ]s into rows ("]") and columns ("]"). It is an ] of ] and is widely used in ] and other sciences. It is a depiction of the ], which states that when the elements are arranged in order of their ]s an approximate ] is evident. The table is divided into four roughly rectangular areas called ]. Elements in the same group tend to show similar chemical characteristics.


Vertical, horizontal and diagonal ] characterize the periodic table. ]lic character increases going down a group and from right to left across a period. ] character increases going from the bottom left of the periodic table to the top right.
All elements from atomic numbers 1 (]) to 118 (]) have been synthesized. Of these, all up to and including ]<!-- see discussion at ] if confounded --> exist naturally; the rest have only been artificially synthesised in laboratories, along with numerous synthetic ] of naturally occurring elements. Production of elements beyond ununoctium is being pursued, with the question of how the periodic table may need to be modified to accommodate these elements being ].


The first periodic table to become generally accepted was that of the Russian chemist ] in 1869; he formulated the periodic law as a dependence of chemical properties on ]. As not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to ]. The periodic law was recognized as a fundamental discovery in the late 19th century. It was explained early in the 20th century, with the discovery of ]s and associated pioneering work in ], both ideas serving to illuminate the internal structure of the atom. A recognisably modern form of the table was reached in 1945 with ]'s discovery that the ]s were in fact f-block rather than d-block elements. The periodic table and law are now a central and indispensable part of modern chemistry.
== History ==
{{Main|History of the periodic table}}
=== First attempts of systemization ===


The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 <!--THIS IS NOT A TYPO: uranium can fission spontaneously, and when the resulting neutrons strike other uranium atoms in the ore, they can be captured, and the subsequent beta decay produces tiny traces of neptunium and plutonium. See the note--> exist;{{efn|name=transuranium}} to go further, it was necessary to ] new elements in the laboratory. By 2010,<!--THE LAST FOUR WERE *NAMED* IN 2016, BUT ALL WERE ALREADY SYNTHESISED BY 2010--> the first 118 elements were known, thereby completing the first seven rows of the table;<ref>{{Cite web |title=Periodic Table of Elements |url=https://iupac.org/what-we-do/periodic-table-of-elements/ |access-date=11 May 2024 |website=IUPAC {{!}} International Union of Pure and Applied Chemistry |language=en-US}}</ref> however, chemical characterization is still needed for the heaviest elements to confirm that their properties match their positions. New discoveries will extend the table ], though it is not yet known how many more elements are possible; moreover, theoretical calculations suggest that this unknown region will not follow the patterns of the known part of the table. Some scientific discussion also continues regarding whether some elements are correctly positioned in today's table. Many ] of the periodic law exist, and there is some discussion as to whether there is an optimal form of the periodic table.
In 1789, ] published a list of 33 ]s. Although Lavoisier grouped the elements into ]es, ]s, ]s, and ],<ref>{{cite book |title=From elements to atoms: a history of chemical composition |last=Siegfried |first=Robert |year=2002 |page=92 |publisher=Library of Congress Cataloging-in-Publication Data |isbn=0-87169-924-9 |accessdate=25 January 2012 |location=Philadelphia, Pennsylvania}}</ref> chemists spent the following century searching for a more precise classification scheme. In 1829, ] observed that many of the elements could be grouped into ''triads'' (groups of three) based on their chemical properties. ], ], and ], for example, were grouped together as soft, ] 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.<ref name="Ball100">Ball, p. 100</ref> This became known as the ].<ref>{{cite book |last=Horvitz |first=Leslie |title=Eureka!: Scientific Breakthroughs That Changed The World |year=2002 |publisher=John Wiley |location=New York|isbn=978-0-471-23341-1 |oclc=50766822 |page=43}}</ref> German chemist ] worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. ] 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.<ref name="Ball100"/>


==Structure<span class="anchor" id="Detailed table"></span> ==
German chemist ] had observed in 1858 that ] has a tendency to bond with other elements in a ratio of one to four. ], for example, has one carbon atom and four hydrogen atoms. This concept eventually became known as '']''. In 1864, fellow German chemist ] published a table of the 49 known elements arranged by valency. The table revealed that elements with similar properties often shared the same valency.<ref>Ball, p. 101</ref>
{{Periodic table}}


] ]s showing probability density and phase (g orbitals and higher are not shown)]]
English chemist ] produced a series of papers in 1864 and 1865 that described his own classification of the elements: He noted that when listed in order of increasing atomic weight, similar physical and chemical properties recurred at intervals of eight, which he likened to the ]s of music.<ref>{{cite journal |title = On Relations Among the Equivalents |author = Newlands, John A. R. |journal = Chemical News |volume = 10 |pages = 94–95 |date = 1864-08-20 |url =http://web.lemoyne.edu/~giunta/EA/NEWLANDSann.HTML#newlands3}}</ref><ref>{{cite journal |title = On the Law of Octaves |author = Newlands, John A. R. |journal = Chemical News |volume = 12 |page = 83 |date = 1865-08-18 |url =http://web.lemoyne.edu/~giunta/EA/NEWLANDSann.HTML#newlands4}}</ref> This ''Law of Octaves'', however, was ridiculed by his contemporaries, and the ] refused to publish his work.<ref>{{cite book |last=Bryson|first=Bill |authorlink=Bill Bryson |title=] |publisher=Black Swan|year=2004 |pages=141–142 |isbn=978-0-552-15174-0}}</ref> Nonetheless, Newlands was able to draft an atomic table and use it to predict the existence of missing elements, such as ]. The Chemical Society only acknowledged the significance of his discoveries some five years after they credited Mendeleev.<ref>{{cite journal |last1=Brock |first1=W. H. |last2=Knight |first3=D. M. |title=The Atomic Debates: 'Memorable and Interesting Evenings in the Life of the Chemical Society' |journal=Isis |volume=56 |issue=1 |year=1965 |pages=5–25 |publisher=]}}</ref>
Each chemical element has a unique ] (''Z''{{--}} for "Zahl", German for "number") representing the number of ]s in its ].<ref name="neutronium">An ] (i.e. a substance composed purely of neutrons), is included in a few alternate presentations, for example, in the . See {{cite journal |last=Labarca |first=M. |title=An element of atomic number zero? |journal=New Journal of Chemistry |year=2016 |volume=40|issue=11|pages=9002–9006 |doi=10.1039/C6NJ02076C|hdl=11336/46854 |hdl-access=free |issn=1144-0546 }}</ref> Each distinct atomic number therefore corresponds to a class of atom: these classes are called the ]s.<ref>{{GoldBookRef |title=Chemical element |file=C01022}}</ref> The chemical elements are what the periodic table classifies and organizes. ] is the element with atomic number 1; ], atomic number 2; ], atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter ]; those for hydrogen, helium, and lithium are respectively H, He, and Li.<ref name="IUPAC-redbook"/> Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called ]s of the same chemical element.<ref name="IUPAC-redbook" /> Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined ]s, defined as the average mass of a naturally occurring atom of that element.<ref name="ciaaw">{{cite web |title=Standard Atomic Weights |website=Commission on Isotopic Abundances and Atomic Weights |date=2019 |publisher=International Union of Pure and Applied Chemistry |url=https://www.ciaaw.org/atomic-weights.htm |access-date=7 February 2021 |url-status=live |archive-date=8 August 2020 |archive-url=https://web.archive.org/web/20200808155924/https://www.ciaaw.org/atomic-weights.htm}}</ref>
All elements have multiple ]s, variants with the same number of protons but different numbers of ]s. For example, ] has three naturally occurring isotopes: all of its ]s 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. When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses.<ref name="Greenwood">Greenwood & Earnshaw, pp. 24–27</ref>


In the standard periodic table, the elements are listed in order of increasing atomic number. A new row (]) is started when a new ] has its first ]. Columns (]) are determined by the ] of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. ], ], and ] are in the same column because they all 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.<ref>Gray, p. 6</ref>
=== Mendeleev's table ===
]
]


Today, 118 elements are known, the first 94 of which are known to occur naturally on Earth at present.<ref name=ThorntonBurdette/>{{efn|The question of how many natural elements there are is quite complicated and is not fully resolved. The heaviest element that occurs in large quantities on Earth is element 92, ]. However, uranium can undergo ] in nature, and the resulting neutrons can strike other uranium atoms. If neutron capture then occurs, elements 93 and 94, ] and ], are formed via ];<ref name=ThorntonBurdette/> these are in fact more common than some of the rarest elements in the first 92, such as ], ], and ] (see ]). Theoretically, neutron capture on the resulting plutonium might produce even higher-numbered elements, but the quantities would be too small to be observed.<ref name=ThorntonBurdette/> In the early Solar System, shorter-lived elements had not yet decayed away, and consequently there were more than 94 naturally occurring elements. ] (element 96) is the longest-lived element beyond the first 94, and is probably still being brought to Earth via ]s, but it has not been found.<ref name=ThorntonBurdette>{{cite journal |last1=Thornton |first1=Brett F. |last2=Burdette |first2=Shawn C. |date=2019 |title=Neutron stardust and the elements of Earth |url=https://www.nature.com/articles/s41557-018-0190-9 |journal=Nature Chemistry |volume=11 |issue=1 |pages=4–10 |doi=10.1038/s41557-018-0190-9 |pmid=30552435 |bibcode=2019NatCh..11....4T |s2cid=54632815 |access-date=19 February 2022 |archive-date=14 August 2021 |archive-url=https://web.archive.org/web/20210814111535/https://www.nature.com/articles/s41557-018-0190-9 |url-status=live }}</ref> Elements up to 99 (]) have been observed in ].<ref name=gopka08>{{cite journal |last1=Gopka |first1=V.F. |last2=Yushchenko |first2=A.V. |last3=Yushchenko |first3=V.A. |last4=Panov |first4=I.V. |last5=Kim |first5=Ch. |date=15 May 2008 |title=Identification of absorption lines of short half-life actinides in the spectrum of Przybylski's star (HD 101065) |journal=Kinematics and Physics of Celestial Bodies |volume=24 |issue=2 |pages=89–98 |doi=10.3103/S0884591308020049 |bibcode = 2008KPCB...24...89G |s2cid=120526363 }}</ref> Elements up to 100 (]) probably occurred in the ] at ], ], but they have long since decayed away.<ref name="emsley">{{cite book |last=Emsley |first=John |date=2011 |title=Nature's Building Blocks: An A-Z guide to the elements |edition=New |publisher=Oxford University Press |location=New York, NY |isbn=978-0-19-960563-7}}</ref> Even heavier elements may be produced in the ] via ]e or ]s, but this has not been confirmed. It is not clear how far they would extend past 100 and how long they would last: calculations suggest that nuclides of mass number around 280 to 290 are formed in the r-process, but quickly ] to nuclides that suffer ], so that 99.9% of the produced ] nuclides would decay within a month.<ref>{{cite journal |last1=Panov |first1=I.V. |date=2017 |title=Formation of Superheavy Elements in Nature |journal=Physics of Atomic Nuclei |volume=81 |issue=1 |pages=57–65 |doi=10.1134/S1063778818010167|s2cid=125149409 }}</ref> If instead they were sufficiently long-lived, they might similarly be brought to Earth via cosmic rays, but again none have been found.<ref name=ThorntonBurdette/>|name=transuranium}} The remaining 24, americium to oganesson (95–118), occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are ] and 11 occur only in decay chains of primordial elements. A few of the latter are so rare that they were not discovered in nature, but were synthesized in the laboratory before it was determined that they do exist in nature after all: ] (element 43), ] (element 61), ] (element 85), ] (element 93), and ] (element 94).<ref name="emsley"/> No element heavier than ] (element 99) has ever been observed in macroscopic quantities in its pure form, nor has ]; ] (element 87) has been only photographed in the form of ] emitted from microscopic quantities (300,000 atoms).<ref>{{cite book| title = The Chemistry of the Actinide and Transactinide Elements| editor1-last = Morss| editor1-first=L. R.|editor2-first = N. M.|editor2-last = Edelstein| editor3-last = Fuger|editor3-first = J.| last = Silva|first = Robert J.| chapter = Fermium, Mendelevium, Nobelium and Lawrencium| publisher = ]| year = 2006| isbn = 978-1-4020-3555-5| location = Dordrecht| edition = 3rd| ref = CITEREFHaire2006}}</ref> Of the 94 natural elements, eighty have a stable isotope and one more (]) has an almost-stable isotope (with a ] of 2.01×10<sup>19</sup>&nbsp;years, over a billion times the ]).<ref name=Bi209alpha2>{{cite journal | last = Marcillac | first = Pierre de |author2=Noël Coron |author3=Gérard Dambier |author4=Jacques Leblanc |author5=Jean-Pierre Moalic |date=April 2003 | title = Experimental detection of α-particles from the radioactive decay of natural bismuth | journal = Nature | volume = 422 | pages = 876–878 | doi = 10.1038/nature01541 | pmid = 12712201 | issue = 6934 | bibcode=2003Natur.422..876D| s2cid = 4415582 }}</ref>{{efn|Some isotopes currently considered stable are theoretically expected to be radioactive with extremely long half-lives: for instance, all the stable isotopes of elements 62 (]), 63 (]), and all elements from 67 (]) onward are expected to undergo ] or ]. However, the predicted half-lives are extremely long (e.g. the alpha decay of <sup>208</sup>Pb to the ground state of <sup>204</sup>Hg is expected to have a half-life greater than 10<sup>120</sup> years), and the decays have never been observed.<ref name="bellidecay">{{cite journal |last1=Belli |first1=P. |last2=Bernabei |first2=R. |last3=Danevich |first3=F. A. |last4=Incicchitti |first4=A. |last5=Tretyak |first5=V. I. |display-authors=3 |title=Experimental searches for rare alpha and beta decays |journal=European Physical Journal A |date=2019 |volume=55 |issue=8 |pages=140–1–140–7 |doi=10.1140/epja/i2019-12823-2 |issn=1434-601X |arxiv=1908.11458|bibcode=2019EPJA...55..140B |s2cid=201664098 }}</ref><ref name="Tretyak2002">{{Cite journal
Russian chemistry professor ] and German chemist ] independently published their periodic tables in 1869 and 1870, respectively.<ref>{{cite journal|last=Mendelejew |first=Dimitri |year=1869 |title=Über die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente |journal=Zeitschrift für Chemie |pages=405–406 |language=German}}</ref> They both constructed their tables in a similar manner: By listing the elements in a row or column in order of atomic weight and starting a new row or column when the characteristics of the elements began to repeat.<ref>Ball, pp. 100–102</ref> The success of 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.<ref>{{cite book |author=Pullman, Bernard |title=The Atom in the History of Human Thought |publisher=Oxford University Press |year=1998 |page=227 |isbn=0-19-515040-6|others=Translated by Axel Reisinger}}</ref> 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 ] and germanium.<ref>Ball, p. 105</ref> The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as ] and ], to better classify them into chemical families. With the development of theories of atomic structure, it became apparent that Mendeleev had listed the elements in order of increasing atomic number.<ref>{{cite book |title=The Periodic Kingdom |author=Atkins, P. W. |publisher=HarperCollins Publishers, Inc. |year=1995 |page=87|isbn=0-465-07265-8}}</ref>


|last1=Tretyak |first1=V.I.
=== Further development ===
|last2=Zdesenko |first2=Yu.G.
|year=2002
|title=Tables of Double Beta Decay Data — An Update
|journal=] |volume=80 |issue=1 |pages=83–116
|doi=10.1006/adnd.2001.0873
|bibcode=2002ADNDT..80...83T }}</ref>}} Two more, ] and ], have isotopes undergoing ] with a half-life comparable to the ]. The stable elements plus bismuth, thorium, and uranium make up the 83 ] elements that survived from the Earth's formation.{{efn|The half-life of ]'s most stable isotope is just long enough that it should also be a primordial element. A 1971 study claimed to have detected primordial plutonium,<ref name="PU244">{{cite journal |first1=D. C. |last1=Hoffman |first2=F. O. |last2=Lawrence |first3=J. L. |last3=Mewherter |first4=F. M. |last4=Rourke |title=Detection of Plutonium-244 in Nature
|journal=] |volume=234 |pages= 132–134 |year=1971 |doi=10.1038/234132a0|bibcode = 1971Natur.234..132H |issue=5325|s2cid=4283169 }}</ref> but a more recent study from 2012 could not detect it.<ref name="PRC">{{cite journal|last=Lachner|first=J.|display-authors=etal|date=2012|title=Attempt to detect primordial <sup>244</sup>Pu on Earth|journal=Physical Review C|volume=85|issue=1|page=015801| doi=10.1103/PhysRevC.85.015801|bibcode=2012PhRvC..85a5801L}}</ref> Based on its likely initial abundance in the Solar System, present experiments as of 2022 are likely about an order of magnitude away from detecting live primordial <sup>244</sup>Pu.<ref>{{cite journal |last1=Wu |first1=Yang |last2=Dai |first2=Xiongxin |first3=Shan |last3=Xing |first4=Maoyi |last4=Luo |first5=Marcus |last5=Christl |first6=Hans-Arno |last6=Synal |first7=Shaochun |last7=Hou |date=2022 |title=Direct search for primordial <sup>244</sup>Pu in Bayan Obo bastnaesite |url=http://www.ccspublishing.org.cn/article/doi/10.1016/j.cclet.2022.03.036?pageType=en |journal=Chinese Chemical Letters |volume=33 |issue=7 |pages=3522–3526 |doi=10.1016/j.cclet.2022.03.036 |s2cid=247443809 |access-date=29 January 2024}}</ref>}} The remaining eleven natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of thorium and uranium.{{efn|Tiny traces of plutonium are also continually brought to Earth via cosmic rays.<ref name="WallnerFaestermann2015">{{cite journal |last1=Wallner |first1=A. |last2=Faestermann |first2=T. |last3=Feige |first3=J. |last4=Feldstein |first4=C. |last5=Knie |first5=K. |last6=Korschinek |first6=G. |last7=Kutschera |first7=W. |last8=Ofan |first8=A. |last9=Paul |first9=M. |last10=Quinto |first10=F. |last11=Rugel |first11=G. |last12=Steier |first12=P. |display-authors=6 |year=2015 |title=Abundance of live {{sup|244}}Pu in deep-sea reservoirs on Earth points to rarity of actinide nucleosynthesis |journal=Nature Communications |volume=6 |page=5956 |issn=2041-1723|doi=10.1038/ncomms6956|pmid=25601158 |pmc=4309418 |arxiv=1509.08054|bibcode=2015NatCo...6.5956W}}</ref>}} All 24 known artificial elements are radioactive.<ref name="IUPAC-redbook" />


=== Group names and numbers ===
In the years following publication of Mendeleev's periodic table, the gaps he identified were filled as chemists discovered additional naturally occurring elements.<ref>{{cite journal |last=Kaji |first=Masanori |year=2002 |title=D.I. Mendeleev's Concept of Chemical Elements and the Principle of Chemistry |journal=Bull. Hist. Chem. |volume=27 |issue=1 |pages=4-16 |publisher=Tokyo Institute of Technology |doi= |url=http://www.scs.illinois.edu/~mainzv/HIST/awards/OPA%20Papers/2005-Kaji.pdf |accessdate=11 June 2012 }}</ref> It is often stated that the last naturally occurring element to be discovered was ] (referred to by Mendeleev as ''eka-caesium'') in 1939.<ref>{{cite web |url=http://chemeducator.org/sbibs/s0010005/spapers/1050387gk.htm|title= Francium (Atomic Number 87), the Last Discovered Natural Element|author1=Adloff, Jean-Pierre |author2=Kaufman, George B. |date=25 September 2005 |publisher=The Chemical Educator |accessdate=26 March 2007}}</ref> However, ], produced synthetically in 1940, was identified in trace quantities as a naturally occurring primordial element in 1971,<ref>{{cite journal |doi = 10.1038/234132a0|url=http://www.nature.com/nature/journal/v234/n5325/abs/234132a0.html |title= Detection of Plutonium-244 in Nature|journal = Nature |pages = 132–134 |year = 1971 |last1 = Hoffman |first1 = D. C. |last2 = Lawrence |first2 = F. O.|last3 = Mewherter |first3 = J. L. |last4 = Rourke |first4 = F. M. |volume = 234 |bibcode = 1971Natur.234..132H|issue=5325}}</ref> and in 2011 it was found that all the elements up to ] can occur naturally in trace amounts in uranium ores.<ref name="emsley"/>
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). The f-block groups are ignored in this numbering.<ref name="IUPAC">{{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=51|url=https://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|access-date=26 November 2018|archive-url=https://web.archive.org/web/20181123034019/https://old.iupac.org/publications/books/rbook/Red_Book_2005.pdf|archive-date=23 November 2018|url-status=live}}</ref> Groups can also be named by their first element, e.g. the "scandium group" for group 3.<ref name="IUPAC"/> Previously, groups were known by ]. In the ], the Roman numerals were followed by either an "A" if the group was in the ] or ], or a "B" if the group was in the ]. The Roman numerals used correspond to the last digit of today's naming convention (e.g. the ]s were group IVB, and the ] were group IVA). In ], the lettering was similar, except that "A" was used for groups 1 through 7, and "B" was used for groups 11 through 17. 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 ] (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated.<ref name="Fluck">{{cite journal |last1=Fluck |first1=E. |year=1988 |title=New Notations in the Periodic Table |journal=] |volume=60 |pages=431–436|doi=10.1351/pac198860030431 |url=https://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |access-date=24 March 2012 |issue=3 |s2cid=96704008 |url-status=live |archive-url=https://web.archive.org/web/20120325152951/https://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-date=25 March 2012}}</ref>


{{Periodic table (group names)}}
With the development of modern ] theories of ] configurations within atoms, it became apparent that each row (or ''period'') in the table corresponded to the filling of a ] of electrons. In Mendeleev's original table, each period was the same length. However, because larger atoms have more ], modern tables have progressively longer periods further down the table.<ref>Ball, p. 111</ref>


=== Presentation forms<span class="anchor" id="The long- or 32-column table"></span> ===
Previously, the groups were known by ]. In America, the roman numerals were followed by either an "A" if the group was in the ] or ], or a "B" if the group was in the ]. The roman numerals used correspond to the last digit of today's naming convention (e.g. the ]s were group IVB, and the ] were group IVA). In Europe, the lettering was similar, except that "A" was used if the group was before ], 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.<ref>{{cite journal |last1=Fluck |first1=E.|year=1988 |title=New Notations in the Periodic Table |journal=]|volume=60 |number=3 |pages=431–436 |publisher=]|doi=10.1351/pac198860030431 |url=http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |accessdate=24 March 2012 }}</ref>
<div style="border:1px solid grey; float:right; text-align:center; padding:0.2em; margin:0; font-size:90%;">
{{Periodic table (32 columns, micro)}}
32 columns
{{Periodic table (18 columns, micro)}}
18 columns
</div>
For reasons of space,<ref name=Petrucci331/><ref>{{cite journal |last1=Pfeiffer |first1=Paul |date=1920 |title=Die Befruchtung der Chemie durch die Röntgenstrahlenphysik |url=https://link.springer.com/article/10.1007/BF02448807 |journal=Naturwissenschaften |language=de |volume=8 |issue=50 |pages=984–991 |bibcode=1920NW......8..984P |doi=10.1007/BF02448807 |s2cid=7071495}}</ref> the periodic table is commonly presented with the f-block elements cut out and positioned as a distinct part below the main body.<ref name="cartoon" /><ref name="Petrucci331" /><ref name="Fluck" /> This reduces the number of element columns from 32 to 18.<ref name=Petrucci331/>


Both forms represent the same periodic table.<ref name="IUPAC-redbook" /> The form with the f-block included in the main body is sometimes called the 32-column<ref name="IUPAC-redbook" /> or long form;<ref name="Thyssen" /> the form with the f-block cut out the 18-column<ref name="IUPAC-redbook" /> or medium-long form.<ref name="Thyssen" /> The 32-column form has the advantage of showing all elements in their correct sequence, but it has the disadvantage of requiring more space.<ref>Scerri, p. 375</ref> The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing ], the options can be shown equally (unprejudiced) in both forms.<ref name="2015IUPAC">{{cite web|url=https://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|archive-url=https://web.archive.org/web/20160705053631/https://iupac.org/projects/project-details/?project_nr=2015-039-2-200|archive-date=5 July 2016}}</ref>
The production of various ] has expanded the periodic table significantly, the first of these being ], synthesized in 1939.<ref>Ball, p. 123</ref> Because many of the transuranic elements are highly unstable and ] quickly, they are challenging to detect and characterize when produced, and there have been ] concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights. The most recently accepted and named elements are ] (114) and ] (116), both named on 31 May 2012.<ref>{{cite journal |author=Barber, Robert C.; Karol, Paul J; Nakahara, Hiromichi; Vardaci, Emanuele; Vogt, Erich W. |title=Discovery of the elements with atomic numbers greater than or equal to 113 (IUPAC Technical Report) |doi=10.1351/PAC-REP-10-05-01 |journal=Pure Appl. Chem. |year=2011 |volume=83 |issue=7|page=1485}}</ref> In 2010, a joint Russia–US collaboration at ], ], Russia, claimed to have synthesized six atoms of ], making it the most recently claimed discovery.<ref name=E117>{{ru icon}} {{cite web|title=Эксперимент по синтезу 117-го элемента получает продолжение|trans_title=Experiment on sythesis of the 117th element is to be continued|url=http://www.jinr.ru/news_article.asp?n_id=1195&language=rus|year=2012|publisher=JINR}}</ref>


Periodic tables usually at least show the elements' symbols; many also provide supplementary information about the elements, either via colour-coding or as data in the cells. The above table shows the names and atomic numbers of the elements, and also their blocks, natural occurrences and ]s. For the short-lived elements without standard atomic weights, the mass number of the most stable known isotope is used instead. Other tables may include properties such as state of matter, melting and boiling points, densities, as well as provide different classifications of the elements.{{efn|See for example }}
== Contents of the periodic table ==
=== Electron configurations ===
{{for|a larger version|Periodic table (large version)}}
{{periodic table}}


{{main|Electron configuration}}
All versions of the periodic table only include chemical elements, not ]s, ], or ]s, and each ] of a given element is represented in the same cell. In the standard periodic table, the elements are listed in order of increasing atomic number (the number of ]s in the ] of an atom). A new row (]) is started when a new ] has its first electron. Columns (]) are determined by the ] of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. ] and ] are in the same column because they both have 4 electrons in the outermost ]). The periods are longer further down in the periodic table, and the groups get longer on the right (although the ]s, the largest group, is on the far left, and the ]s, another large group, are next to the alkali metals). In general, elements with similar chemical properties 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.<ref>Gray, p. 6</ref>


The periodic table is a graphic description of the periodic law,<ref name="Scerri17"/> which states that the properties and atomic structures of the chemical elements are a ] of their ].<ref>{{cite Merriam-Webster |periodic law |access-date=29 March 2021}}</ref> Elements are placed in the periodic table according to their ]s,<ref name="Jensen2009"/> the periodic recurrences of which explain the ] in properties across the periodic table.<ref name="FIII19" />
As of 2012, the periodic table contains 118 confirmed chemical elements. Of these elements, 114 have been officially recognized and named by the ] (IUPAC). A total of 98 of these occur naturally, of which 84 are ]. The other 14 elements only occur in ]s of primordial elements.<ref name="emsley">{{cite book|last=Emsley|first=John|title=Nature's Building Blocks: An A-Z Guide to the Elements|edition=New|year=2011|publisher=Oxford University Press|location=New York, NY|isbn=978-0-19-960563-7}}</ref> All elements from einsteinium to ], as well as ], and ], while not occurring naturally in the universe, have been officially recognized by the IUPAC as being synthesized, while elements ], ], ] and ] have reportedly been synthesized in laboratories and are currently known only by their systematic element names, based off their atomic numbers.<ref>{{cite journal |last=Koppenol |first=W. H. |title=Naming of New Elements (IUPAC Recommendations 2002) |journal=Pure and Applied Chemistry |year=2002 |volume=74 |issue=5 |pages=787–791 |url=http://media.iupac.org/publications/pac/2002/pdf/7405x0787.pdf |format=PDF |doi=10.1351/pac200274050787}}</ref> No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form.<ref>{{cite book
| title = The Chemistry of the Actinide and Transactinide Elements
| editor1-last = Morss|editor2-first = Norman M.
| editor2-last = Edelstein
| editor3-last = Fuger|editor3-first = Jean
| last = Haire|first = Richard G.
| chapter = Fermium, Mendelevium, Nobelium and Lawrencium
| publisher = ]
| year = 2006
| isbn = 1-4020-3555-1
| location = Dordrecht, The Netherlands
| edition = 3rd
| ref = CITEREFHaire2006
}}</ref>
No elements past 118 have been synthesized as of 2012.<ref>Gray, p. 233</ref>


An electron can be thought of as inhabiting an ], which characterizes the probability it can be found in any particular region around the atom. Their energies are ], which is to say that they can only take discrete values. Furthermore, electrons obey the ]: different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain one or more orbitals. Each orbital can contain up to two electrons: they are distinguished by a quantity known as ], conventionally labelled "up" or "down".<ref>Petrucci et al., p. 323</ref>{{efn|Strictly speaking, one cannot draw an orbital such that the electron is guaranteed to be inside it, but it can be drawn to guarantee a 90% probability of this for example.<ref>Petrucci et al., p. 306</ref>}} In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimized by occupying the lowest-energy orbitals available.<ref>Petrucci et al., p. 322</ref> Only the outermost electrons (so-called ]s) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called ]s.<ref>{{cite book |last1=Ball |first1=David W. |last2=Key |first2=Jessie A. |date=2011 |title=Introductory Chemistry |edition=1st Canadian |place=Vancouver, British Columbia |publisher=BC Campus (opentextbc.ca) |isbn=978-1-77420-003-2 |url=https://opentextbc.ca/introductorychemistry/chapter/electronic-structure-and-the-periodic-table/ |access-date=15 August 2021 |archive-date=15 August 2021 |archive-url=https://web.archive.org/web/20210815072718/https://opentextbc.ca/introductorychemistry/chapter/electronic-structure-and-the-periodic-table/ |url-status=live |page=}}</ref>
In printed or other formally presented periodic tables, each element is provided a formatted cell that usually provides some of the basic properties of the element. Atomic number, ], and name are almost always included, and ]s, ], ] and ]s, crystal structure as a solid, origin, abbreviated electron configuration, electronegativity, and most common ]s are sometimes included as well.<ref>{{cite book |title=Kaplan SAT Subject Test: Chemistry 2009–2012 Edition |year=2009 |publisher=Kaplan Publishing |isbn=978-1-4195-5260-1}}</ref>


{| class="wikitable" style="float:right; margin:0.5em; text-align:center;"
By definition, each chemical element has a unique atomic number representing the number of protons in its nucleus, but most elements have differing numbers of ]s among different atoms; these are referred to as ]s. For example, all atoms of carbon have six protons and usually have six neutrons as well, but about 1% have seven neutrons, and a very small amount have eight neutrons; so carbon has three different naturally occurring isotopes. 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 listed in parentheses.<ref name="Greenwood">{{Greenwood&Earnshaw1st|pages=24–27}}</ref>
! style="text-align:right;" |ℓ =
! 0
! 1
! 2
! 3
! 4
! 5
! 6
! rowspan=2 | Shell capacity (2''n''<sup>2</sup>)<ref>{{cite web |date=6 May 2020 |title=Electron Configurations |url=https://www.chem.fsu.edu/chemlab/chm1045/e_config.html |access-date=17 April 2022 |website=www.chem.fsu.edu |publisher=Florida State University |archive-date=6 May 2022 |archive-url=https://web.archive.org/web/20220506074340/https://www.chem.fsu.edu/chemlab/chm1045/e_config.html |url-status=live }}</ref>
|-
! style="text-align:right;" | Orbital
! s
! p
! d
! f
! g
! h
! i
|-
! ''n'' = 1
| bgcolor="{{element color|s-block}}" | 1s
| colspan=6 |
| 2
|-
! ''n'' = 2
| bgcolor="{{element color|s-block}}" | 2s
| bgcolor="{{element color|p-block}}" | 2p
| colspan=5 |
| 8
|-
! ''n'' = 3
| bgcolor="{{element color|s-block}}" | 3s
| bgcolor="{{element color|p-block}}" | 3p
| bgcolor="{{element color|d-block}}" | 3d
| colspan=4 |
| 18
|-
! ''n'' = 4
| bgcolor="{{element color|s-block}}" | 4s
| bgcolor="{{element color|p-block}}" | 4p
| bgcolor="{{element color|d-block}}" | 4d
| bgcolor="{{element color|f-block}}" | 4f
| colspan=3 |
| 32
|-
! ''n'' = 5
| bgcolor="{{element color|s-block}}" | 5s
| bgcolor="{{element color|p-block}}" | 5p
| bgcolor="{{element color|d-block}}" | 5d
| bgcolor="{{element color|f-block}}" | 5f
| bgcolor="{{element color|g-block}}" | 5g
| colspan=2 |
| 50
|-
! ''n'' = 6
| bgcolor="{{element color|s-block}}" | 6s
| bgcolor="{{element color|p-block}}" | 6p
| bgcolor="{{element color|d-block}}" | 6d
| bgcolor="{{element color|f-block}}" | 6f
| bgcolor="{{element color|g-block}}" | 6g
| bgcolor="{{element color|h-block}}" | 6h
|
| 72
|-
! ''n'' = 7
| bgcolor="{{element color|s-block}}" | 7s
| bgcolor="{{element color|p-block}}" | 7p
| bgcolor="{{element color|d-block}}" | 7d
| bgcolor="{{element color|f-block}}" | 7f
| bgcolor="{{element color|g-block}}" | 7g
| bgcolor="{{element color|h-block}}" | 7h
| bgcolor="{{element color|i-block}}" | 7i
| 98
|-
! Subshell capacity (4ℓ+2)
| 2
| 6
| 10
| 14
| 18
| 22
| 26
|
|}
Elements are known with up to the first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, and it also contains three dumbbell-shaped 2p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has a capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32.<ref name="Petrucci331" /> Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the ground states of known elements.<ref name="Goudsmit" /> The subshell types are characterized by the ]s. Four numbers describe an orbital in an atom completely: the ] ''n'', the ] ℓ (the orbital type), the ] ''m''<sub>ℓ</sub>, and the ] ''m<sub>s</sub>''.<ref name="FIII19" />


==== Order of subshell filling ====
== Organization ==


] ]]
In the modern periodic table, the elements are placed progressively in each period from left to right in the sequence of their atomic numbers, with a new row started after a noble gas. The first element in the next row is always an alkali metal with an atomic number one greater than that of the noble gas (e.g. after ], a noble gas with the atomic number&nbsp;36, a new row is started by ], an alkali metal with the atomic number&nbsp;37). No gaps currently exist because all elements between ] and ununoctium (element&nbsp;118) have been discovered. Since the elements are sequenced by atomic number, sets of elements are sometimes specified by terms such as "through" (e.g. through ]), "beyond" (e.g. beyond ]), or "from ... through" (e.g. from ] through ]). The terms "light" and "heavy" are sometimes also used informally to indicate relative atomic numbers (not ]), as in "lighter than ]" or "heavier than ]", although technically the weight or mass of atoms of an element (their atomic weights or atomic masses) do not always increase ] with their atomic numbers. For instance ], element&nbsp;52, is on average heavier than ], element&nbsp;53.<ref name="Greenwood"/>
The sequence in which the subshells are filled is given in most cases by the ], also known as the Madelung or Klechkovsky rule (after ] and ] respectively). This rule was first observed empirically by Madelung, and Klechkovsky and later authors gave it theoretical justification.<ref name=Jolly>{{cite book |last1=Jolly |first1=William L. |title=Modern Inorganic Chemistry |edition=1st |publisher=McGraw-Hill |date=1984 |pages= |isbn=0-07-032760-2 |url=https://archive.org/details/trent_0116300649799/page/10 }}</ref><ref name=Ostrovsky/><ref name=Ostrovsky1981/><ref name=Wong/>{{efn|name=lowdin}} The shells overlap in energies, and the Madelung rule specifies the sequence of filling according to:<ref name="Ostrovsky">{{cite journal |last1=Ostrovsky |first1=V. N. |date=May 2001 |title=What and How Physics Contributes to Understanding the Periodic Law |journal=Foundations of Chemistry |volume=3 |issue=2 |pages=145–181 |doi=10.1023/A:1011476405933 |s2cid=15679915 }}</ref>
:1s ≪ 2s < 2p ≪ 3s < 3p ≪ 4s < 3d < 4p ≪ 5s < 4d < 5p ≪ 6s < 4f < 5d < 6p ≪ 7s < 5f < 6d < 7p ≪ ... <!--write in 8s and 5g when they get discovered-->
Here the sign ≪ means "much less than" as opposed to < meaning just "less than".<ref name="Ostrovsky"/> Phrased differently, electrons enter orbitals in order of increasing ''n'' + ℓ, and if two orbitals are available with the same value of ''n'' + ℓ, the one with lower ''n'' is occupied first.<ref name="Goudsmit" /><ref name="Wong">{{cite journal |title=Theoretical justification of Madelung's rule |journal=] |last=Wong |first=D. Pan |date=1979 |issue=11 |pages=714–718 |volume=56 |doi=10.1021/ed056p714 |bibcode = 1979JChEd..56..714W }}</ref> In general, orbitals with the same value of ''n'' + ℓ are similar in energy, but in the case of the s-orbitals (with ℓ = 0), quantum effects raise their energy to approach that of the next ''n'' + ℓ group. Hence the periodic table is usually drawn to begin each row (often called a period) with the filling of a new s-orbital, which corresponds to the beginning of a new shell.<ref name=Ostrovsky/><ref name=Ostrovsky1981>{{cite journal |last1=Ostrovsky |first1=V. N. |date=1981 |title=Dynamic symmetry of atomic potential |url= |journal=Journal of Physics B: Atomic and Molecular Physics |volume=14 |issue=23 |pages=4425–4439 |doi=10.1088/0022-3700/14/23/008 |bibcode=1981JPhB...14.4425O }}</ref><ref name="Petrucci331" /> Thus, with the exception of the first row, each period length appears twice:<ref name=Ostrovsky/>
:2, 8, 8, 18, 18, 32, 32, ...


The overlaps get quite close at the point where the d-orbitals enter the picture,<ref name="Petrucci328"/> and the order can shift slightly with atomic number<ref name=Cao/> and atomic charge.<ref name="Jorgensen"/>{{efn|
Hydrogen and ] are often placed in different places than their electron configurations would indicate; Hydrogen is usually placed above lithium, in accordance with its electron configuration, but is sometimes placed above ],<ref name="hydrogen" /> or even carbon,<ref name="hydrogen">{{cite journal |last=Cronyn |first=Marshall W. |title=The Proper Place for Hydrogen in the Periodic Table |journal=Journal of Chemical Education |volume=80 |issue=8 |year=2003 |month=August |pages=947–951|bibcode = 2003JChEd..80..947C |doi = 10.1021/ed080p947 }}</ref> as it also behaves similarly to them. Helium is almost always placed above ], as they are very similar chemically.<ref name="Gray12">Gray, p. 12</ref>
Once two to four electrons are removed, the d and f orbitals usually become lower in energy than the s ones:<ref name="Jorgensen"/>
:1s ≪ 2s < 2p ≪ 3s < 3p ≪ 3d < 4s < 4p ≪ 4d < 5s < 5p ≪ 4f < 5d < 6s < 6p ≪ 5f < 6d < 7s < 7p ≪ ...


and in the limit for extremely highly charged ions, orbitals simply fill in the order of increasing ''n'' instead. There is a gradual transition between the limiting situations of highly charged ions (increasing ''n'') and neutral atoms (Madelung's rule).<ref name="Goudsmit"/> Thus for example, the energy order for the 55th electron outside the xenon core proceeds as follows in the isoelectronic series of caesium (55 electrons):<ref name=elyashevich/>
The significance of atomic numbers to the organization of the periodic table was not appreciated until the existence and properties of protons and neutrons became understood. Mendeleev's periodic tables instead used atomic weights, information determinable to fair precision in his time, which worked well enough in most cases to give a powerfully predictive presentation far better than any other comprehensive portrayal of the chemical elements' properties then possible. Substitution of atomic numbers, once understood, gave a definitive, integer-based sequence for the elements, still used today even as new synthetic elements are being produced and studied.<ref>{{cite journal|journal=Nucl. Phys. A|volume=789|pages=142–154|year=2007|title=Predictions of alpha decay half lives of heavy and superheavy elements|last1=Samanta|first1=C.|last2=Chowdhury|first2=P. Roy|last3=Basu|first3=D.N.|doi=10.1016/j.nuclphysa.2007.04.001|bibcode=2007NuPhA.789..142S|arxiv = nucl-th/0703086 }}</ref>
:Cs<sup>0</sup>: 6s < 6p < 5d < 7s < 4f
:Ba<sup>+</sup>: 6s < 5d < 6p < 7s < 4f
:La<sup>2+</sup>: 5d < 4f < 6s < 6p < 7s
:Ce<sup>3+</sup>: 4f < 5d < 6s < 6p < 7s
and in the isoelectronic series of holmium (67 electrons), a Ho<sup>0</sup> atom is 4f<sup>11</sup>6s<sup>2</sup>, but Er<sup>+</sup> is 4f<sup>12</sup>6s<sup>1</sup>, Tm<sup>2+</sup> through W<sup>7+</sup> are 4f<sup>13</sup>, and from Re<sup>8+</sup> onward the configuration is 4f<sup>14</sup>5p<sup>5</sup> following the hydrogenic order.<ref name=rareearths/><ref>{{cite web |url=https://physics.nist.gov/cgi-bin/ASD/ie.pl?spectra=Ho-like&submit=Retrieve+Data&units=1&format=0&order=0&at_num_out=on&sp_name_out=on&ion_charge_out=on&el_name_out=on&seq_out=on&shells_out=on&level_out=on&ion_conf_out=on&e_out=0&unc_out=on&biblio=on |title=NIST Atomic Spectra Database: Ionization Energies Data: All Ho-like |author=NIST |date=2023 |website=nist.gov |publisher=NIST |access-date=5 January 2024 |quote=}}</ref>
:
Also, the ordering of the orbitals between each ≪ changes somewhat throughout each period. For example, the ordering in argon and potassium is 3p ≪ 4s < 4p ≪ 3d; by calcium it has become 3p ≪ 4s < 3d < 4p; from scandium to copper it is 3p ≪ 3d < 4s < 4p; and from zinc to krypton it is 3p < 3d ≪ 4s < 4p<ref name=Cao>{{cite journal |last1=Cao |first1=Changsu |last2=Vernon |first2=René E. |first3=W. H. Eugen |last3=Schwarz |first4=Jun |last4=Li |date=6 January 2021 |title=Understanding Periodic and Non-periodic Chemistry in Periodic Tables |journal=Frontiers in Chemistry |volume=8 |issue=813 |page=813 |doi=10.3389/fchem.2020.00813 |pmid=33490030 |pmc=7818537 |bibcode=2021FrCh....8..813S |doi-access=free }}</ref> as the d-orbitals fall into the core at gallium.<ref>{{cite journal |last1=Tossell |first1=J.A. |date=1 November 1977 |title=Theoretical studies of valence orbital binding energies in solid zinc sulfide, zinc oxide, and zinc fluoride |journal=Inorganic Chemistry |volume=16 |issue=11 |pages=2944–2949 |doi=10.1021/ic50177a056}}</ref><ref name=KW/> Deeply buried core shells in heavy atoms thus come closer to the hydrogenic order: around osmium (''Z'' {{=}} 76) 4f falls below 5p, and around bismuth (''Z'' {{=}} 83) 4f falls below 5s as well.<ref name=rareearths/>
}}


Starting from the simplest atom, this lets us build up the periodic table one at a time in order of atomic number, by considering the cases of single atoms. In ], there is only one electron, which must go in the lowest-energy orbital 1s. This ] is written 1s<sup>1</sup>, where the superscript indicates the number of electrons in the subshell. ] adds a second electron, which also goes into 1s, completely filling the first shell and giving the configuration 1s<sup>2</sup>.<ref name="FIII19">{{cite book |last1=Feynman |first1=Richard |last2=Leighton |first2=Robert B. |last3=Sands |first3=Matthew |date=1964 |title=The Feynman Lectures on Physics |url=https://feynmanlectures.caltech.edu/III_19.html |publisher=Addison–Wesley |volume=3 |chapter=19. The Hydrogen Atom and The Periodic Table |isbn=0-201-02115-3 |access-date=15 August 2021 |archive-date=19 October 2021 |archive-url=https://web.archive.org/web/20211019202245/https://www.feynmanlectures.caltech.edu/III_19.html |url-status=live }}</ref><ref name=jensenlaw>{{cite web|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/081.%20Periodic%20Table.pdf|archive-url=https://web.archive.org/web/20201110113324/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/081.%20Periodic%20Table.pdf|archive-date=10 November 2020|last1=Jensen|first1=William B.|author-link=William B. Jensen|title=The Periodic Law and Table|date=2000|access-date=10 December 2022}}</ref>{{efn|In fact, electron configurations represent a first-order approximation: an atom really exists in a superposition of multiple configurations, and electrons in an atom are indistinguishable.<ref name=Scerri2009/> The elements in the d- and f-blocks have multiple configurations separated by small energies and can change configuration depending on the chemical environment.<ref name=Jorgensen/> In some of the undiscovered g-block elements, mixing of configurations may become so important that the result can no longer be well-described by a single configuration.<ref name=nefedov/>}}
=== Grouping methods ===


Starting from the third element, ], the first shell is full, so its third electron occupies a 2s orbital, giving a 1s<sup>2</sup> 2s<sup>1</sup> configuration. The 2s electron is lithium's only valence electron, as the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms: such a shell is called a "]". The 1s subshell is a core shell for all elements from lithium onward. The 2s subshell is completed by the next element ] (1s<sup>2</sup> 2s<sup>2</sup>). The following elements then proceed to fill the 2p subshell. ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>1</sup>) puts its new electron in a 2p orbital; ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>2</sup>) fills a second 2p orbital; and with ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>3</sup>) all three 2p orbitals become singly occupied. This is consistent with ], which states that atoms usually prefer to singly occupy each orbital of the same type before filling them with the second electron. ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>4</sup>), ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>5</sup>), and ] (1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>6</sup>) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely.<ref name="FIII19" /><ref name=jensenlaw/>
==== Groups ====
{{Main|Group (periodic table)}}


Starting from element 11, ], the second shell is full, making the second shell a core shell for this and all heavier elements. The eleventh electron begins the filling of the third shell by occupying a 3s orbital, giving a configuration of 1s<sup>2</sup> 2s<sup>2</sup> 2p<sup>6</sup> 3s<sup>1</sup> for sodium. This configuration is abbreviated 3s<sup>1</sup>, where represents neon's configuration. ] ( 3s<sup>2</sup>) finishes this 3s orbital, and the following six elements ], ], ], ], ], and ] fill the three 3p orbitals ( 3s<sup>2</sup> 3p<sup>1</sup> through 3s<sup>2</sup> 3p<sup>6</sup>).<ref name="FIII19"/><ref name=jensenlaw/> This creates an analogous series in which the outer shell structures of sodium through argon are analogous to those of lithium through neon, and is the basis for the periodicity of chemical properties that the periodic table illustrates:<ref name="FIII19" /> at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat.<ref name="Scerri17">Scerri, p. 17</ref>
A ''group'' or ''family'' is a vertical column in the periodic table. Groups are considered the most important method of classifying the elements. In some groups, the elements have very similar properties and exhibit a clear trend in properties down the group. Under the international naming system, the groups are numbered numerically from 1 to 18 from the leftmost column (the alkali metals) to the rightmost column (the noble gases).<ref name="IUPAC">{{cite book |title=Nomenclature of Inorganic Chemistry: Recommendations 1990|last=Leigh |first=G. J. |year=1990 |publisher=Blackwell Science |isbn=0-632-02494-1}}</ref> The older naming systems differed slightly between Europe and America.<ref>{{cite web |last=Leigh |first=Jeffery |title=Periodic Tables and IUPAC|url=http://www.iupac.org/publications/ci/2009/3101/1_leigh.html |publisher=Chemistry International: The News Magazine of The International Union of Pure and Applied Chemistry (IUPAC) |accessdate=23 March 2011}}</ref>


The first 18 elements can thus be arranged as the start of a periodic table. Elements in the same column have the same number of valence electrons and have analogous valence electron configurations: these columns are called groups. The single exception is helium, which has two valence electrons like beryllium and magnesium, but is typically placed in the column of neon and argon to emphasise that its outer shell is full. (Some contemporary authors question even this single exception, preferring to consistently follow the valence configurations and place helium over beryllium.) There are eight columns in this periodic table fragment, corresponding to at most eight outer-shell electrons.<ref name="cartoon">{{cite book |last1=Gonick |first1=First |last2=Criddle |first2=Craig |date=2005 |title=The Cartoon Guide to Chemistry |publisher=Collins |pages=17–65 |isbn=0-06-093677-0}}</ref> A period begins when a new shell starts filling.<ref name="Petrucci331" /> Finally, the colouring illustrates the ]: the elements in the s-block (coloured red) are filling s-orbitals, while those in the p-block (coloured yellow) are filling p-orbitals.<ref name="Petrucci331" />
Some of these groups have been given trivial (unsystematic) names, such as the ]s, ], ]s, ]s, ]s, and ]es. However, some other groups, such as ], have no trivial names and are referred to simply by their group numbers, since they display fewer similarities and/or vertical trends.<ref name="IUPAC" />


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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 ], which is the most important factor in accounting for their similar properties.<ref>{{cite book |last=Scerri |first=E. R.|title=The Periodic Table: Its Story and its Significance |publisher=Oxford University Press |year=2007|isbn=978-0-19-530573-9 }}</ref>
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Starting the next row, for ] and ] the 4s subshell is the lowest in energy, and therefore they fill it.<ref name="FIII19"/><ref name=jensenlaw/> Potassium adds one electron to the 4s shell ( 4s<sup>1</sup>), and calcium then completes it ( 4s<sup>2</sup>). However, starting from ] ( 3d<sup>1</sup> 4s<sup>2</sup>) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling the electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at ] the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable for a chromium atom to have a 3d<sup>5</sup> 4s<sup>1</sup> configuration than an 3d<sup>4</sup> 4s<sup>2</sup> one. A similar anomaly occurs at ], whose atom has a 3d<sup>10</sup> 4s<sup>1</sup> configuration rather than the expected 3d<sup>9</sup> 4s<sup>2</sup>.<ref name="FIII19" /> These are violations of the Madelung rule. Such anomalies, however, do not have any chemical significance:<ref name="Jorgensen" /> most chemistry is not about isolated gaseous atoms,<ref>Wulfsberg, p. 27</ref> and the various configurations are so close in energy to each other<ref name="Petrucci328">Petrucci et al., p. 328</ref> that the presence of a nearby atom can shift the balance.<ref name="FIII19" /> Therefore, the periodic table ignores them and considers only idealized configurations.<ref name="Jensen2009">{{cite journal|author1-link=William B. Jensen |last1=Jensen |first1=William B. |date=2009 |title=Misapplying the Periodic Law |journal=Journal of Chemical Education |volume=86 |issue=10 |page=1186 |doi=10.1021/ed086p1186 |bibcode=2009JChEd..86.1186J |doi-access=free }}</ref>
Elements in the same group show patterns in ], ], and ]. 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.<ref name="For Dummies">Moore, p. 111</ref>


At ] ( 3d<sup>10</sup> 4s<sup>2</sup>), the 3d orbitals are completely filled with a total of ten electrons.<ref name="FIII19"/><ref name=jensenlaw/> Next come the 4p orbitals, completing the row, which are filled progressively by ] ( 3d<sup>10</sup> 4s<sup>2</sup> 4p<sup>1</sup>) through ] ( 3d<sup>10</sup> 4s<sup>2</sup> 4p<sup>6</sup>), in a manner analogous to the previous p-block elements.<ref name="FIII19" /><ref name=jensenlaw/> From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry.<ref name=KW/> The s- and p-block elements, which fill their outer shells, are called ]s; the d-block elements (coloured blue below), which fill an inner shell, are called ]s (or transition metals, since they are all metals).<ref name="Petrucci326">Petrucci et al., pp. 326–7</ref>
==== Periods ====
{{Main|Period (periodic table)}}


The next 18 elements fill the 5s orbitals (] and ]), then 4d (] through ], again with a few anomalies along the way), and then 5p (] through ]).<ref name=Petrucci331/><ref name=jensenlaw/> Again, from indium onward the 4d orbitals are in the core.<ref name=jensenlaw/><ref>{{cite journal |last1=Farberovich |first1=O. V. |last2=Kurganskii |first2=S. I. |last3=Domashevskaya |first3=E. P. |date=1980 |title=Problems of the OPW Method. II. Calculation of the Band Structure of ZnS and CdS |url= |journal=Physica Status Solidi B |volume=97 |issue=2 |pages=631–640 |doi=10.1002/pssb.2220970230 |bibcode=1980PSSBR..97..631F }}</ref> Hence the fifth row has the same structure as the fourth.<ref name="Petrucci331" />
A ''period'' or ''series'' is a horizontal row in the periodic table. Although groups are the most common way of classifying elements, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the ]s and ]s form two substantial horizontal series of elements.<ref>{{Cite book|last=Stoker|first=Stephen H.|title=General, organic, and biological chemistry|year=2007|page=68|publisher=Houghton Mifflin|location=New York|isbn=978-0-618-73063-6|oclc=52445586}}</ref>


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Elements in the same period show trends in atomic radius, ionization energy, ], 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.<ref>{{cite book |last=Mascetta |first=Joseph |title=Chemistry The Easy Way |year=2003|publisher=Hauppauge |location=New York |isbn=978-0-7641-1978-1 |oclc=52047235|edition=4th |page=50}}</ref>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.<ref name="For Dummies"/> 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.<ref>{{cite book |last1=Kotz |first1=John |last2=Treichel |first2=Paul|last3=Townsend |first3=John |title=Chemistry and Chemical Reactivity, Volume 2|edition=7th |year=2009|publisher=Thomson Brooks/Cole |location=Belmont |isbn=978-0-495-38712-1 |oclc=220756597 |page=324}}</ref>
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The sixth row of the table likewise starts with two s-block elements: ] and ].<ref name=jensenlaw/> After this, the first f-block elements (coloured green below) begin to appear, starting with ]. These are sometimes termed inner transition elements.<ref name="Petrucci326" /> As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations;<ref name="Petrucci328" /> this resulted in some dispute about where exactly the f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle.<ref name="Jensen2015">{{cite journal |last1=Jensen |first1=William B. |date=2015 |title=The positions of lanthanum (actinium) and lutetium (lawrencium) in the periodic table: an update |url=https://link.springer.com/article/10.1007/s10698-015-9216-1 |journal=Foundations of Chemistry |volume=17 |issue= |pages=23–31 |doi=10.1007/s10698-015-9216-1 |s2cid=98624395 |access-date=28 January 2021 |archive-date=30 January 2021 |archive-url=https://web.archive.org/web/20210130011116/https://link.springer.com/article/10.1007/s10698-015-9216-1 |url-status=live }}</ref> Even though lanthanum does not itself fill the 4f subshell as a single atom, because of repulsion between electrons,<ref name="Jorgensen">{{cite journal |last1=Jørgensen |first1=Christian |date=1973 |title=The Loose Connection between Electron Configuration and the Chemical Behavior of the Heavy Elements (Transuranics) |journal=Angewandte Chemie International Edition |volume=12 |issue=1 |pages=12–19 |doi=10.1002/anie.197300121}}</ref> its 4f orbitals are low enough in energy to participate in chemistry.<ref name="Hamilton">{{cite journal |last1=Hamilton |first1=David C. |date=1965 |title=Position of Lanthanum in the Periodic Table |journal=American Journal of Physics |volume=33 |issue=8 |pages=637–640 |doi=10.1119/1.1972042|bibcode=1965AmJPh..33..637H }}</ref><ref name=elyashevich>{{cite book |last=El'yashevich |first=M. A. |author-link= |date=1953 |title=Spectra of the Rare Earths |url= |location=Moscow |publisher=State Publishing House of Technical-Theoretical Literature |pages=382, 397 |isbn=}}</ref><ref name=Cp3Ln>{{cite journal | last1=Krinsky | first1=Jamin L. | last2=Minasian | first2=Stefan G. | last3=Arnold | first3=John | title=Covalent Lanthanide Chemistry Near the Limit of Weak Bonding: Observation of (CpSiMe<sub>3</sub>)<sub>3</sub>Ce−ECp* and a Comprehensive Density Functional Theory Analysis of Cp<sub>3</sub>Ln−ECp (E = Al, Ga) | journal=Inorganic Chemistry | publisher=American Chemical Society (ACS) | volume=50 | issue=1 | date=8 December 2010 | issn=0020-1669 | doi=10.1021/ic102028d | pages=345–357| pmid=21141834 }}</ref> At ], the seven 4f orbitals are completely filled with fourteen electrons; thereafter, a series of ten transition elements (] through ]) follows,<ref name=jensenlaw/><ref name="JensenLr">{{cite web|url=https://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 |archive-url=https://web.archive.org/web/20151223091325/https://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/251.%20Lawrencium.pdf |archive-date=23 December 2015 }}</ref><ref>{{cite journal |last1=Wang |first1=Fan |last2=Le-Min |first2=Li |date=2002 |title=镧系元素 4f 轨道在成键中的作用的理论研究 |trans-title=Theoretical Study on the Role of Lanthanide 4f Orbitals in Bonding |language=zh |journal=Acta Chimica Sinica |volume=62 |issue=8 |pages=1379–84}}</ref><ref name="LaF3">{{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|pmid=23598823 |bibcode=2013PCCP...15.7839X }}</ref> and finally six main-group elements (] through ]) complete the period.<ref name=jensenlaw/><ref name="Pyykko">{{cite journal
| title = Octacarbonyl Ion Complexes of Actinides +/− (An=Th, U) and the Role of f Orbitals in Metal–Ligand Bonding
| first1= Chaoxian |last1=Chi |first2=Sudip |last2=Pan | first3= Jiaye |last3=Jin |first4=Luyan |last4=Meng | first5= Mingbiao |last5=Luo |first6=Lili |last6=Zhao |first7=Mingfei |last7=Zhou |first8=Gernot |last8=Frenking
| journal = ]
| year = 2019
| volume = 25
| issue = 50
| pages = 11772–11784
| doi = 10.1002/chem.201902625
| pmid= 31276242 | pmc= 6772027 |doi-access=free }}</ref> From lutetium onwards the 4f orbitals are in the core,<ref name=jensenlaw/><ref name=Cp3Ln/> and from thallium onwards so are the 5d orbitals.<ref name=jensenlaw/><ref name=KW/><ref>{{cite journal |last1=Singh |first1=Prabhakar P. |date=1994 |title=Relativistic effects in mercury: Atom, clusters, and bulk |url= |journal=Physical Review B |volume=49 |issue=7 |pages=4954–4958 |doi=10.1103/PhysRevB.49.4954 |pmid=10011429 |bibcode=1994PhRvB..49.4954S }}</ref>

The seventh row is analogous to the sixth row: 7s fills (] and ]), then 5f (] to ]), then 6d (] to ]), and finally 7p (] to ]).<ref name=jensenlaw/> Starting from lawrencium the 5f orbitals are in the core,<ref name=jensenlaw/> and probably the 6d orbitals join the core starting from nihonium.<ref name=jensenlaw/><ref name=VI>{{cite journal |last1=Hu |first1=Shu-Xian |last2=Zou |first2=Wenli |date=23 September 2021 |title=Stable copernicium hexafluoride (CnF<sub>6</sub>) with an oxidation state of VI+ |journal=Physical Chemistry Chemical Physics |volume=2022 |issue=24 |pages=321–325 |doi=10.1039/D1CP04360A|pmid=34889909 |bibcode=2021PCCP...24..321H }}</ref>{{efn|Compounds that would use the 6d orbitals of nihonium as valence orbitals have been theoretically investigated, but they are all expected to be too unstable to observe.<ref name="Seth">{{cite journal |last1=Seth |first1=Michael |last2=Schwerdtfeger |first2=Peter |first3=Knut |last3=Fægri |date=1999 |title=The chemistry of superheavy elements. III. Theoretical studies on element 113 compounds |journal=Journal of Chemical Physics |volume=111 |issue=14 |pages=6422–6433 |doi=10.1063/1.480168 |bibcode=1999JChPh.111.6422S|s2cid=41854842 |doi-access=free |hdl=2292/5178 |hdl-access=free }}</ref>}} Again there are a few anomalies along the way:<ref name="Petrucci331">Petrucci et al., p. 331</ref> for example, as single atoms neither actinium nor ] actually fills the 5f subshell, and lawrencium does not fill the 6d shell, but all these subshells can still become filled in chemical environments.<ref>{{cite journal |last1=Kelley |first1=Morgan P. |last2=Deblonde |first2=Gauthier J.-P. |first3=Jing |last3=Su |first4=Corwin H. |last4=Booth |first5=Rebecca J. |last5=Abergel |first6=Enrique R. |last6=Batista |first7=Ping |last7=Yang |date=2018 |title=Bond Covalency and Oxidation State of Actinide Ions Complexed with Therapeutic Chelating Agent 3,4,3-LI(1,2-HOPO) |url= https://escholarship.org/uc/item/4tc1b0xz|journal=Inorganic Chemistry |volume=57 |issue=9 |pages=5352–5363 |doi=10.1021/acs.inorgchem.8b00345 |pmid=29624372 |osti=1458511 }}</ref><ref name="Johansson">{{cite journal|last1=Johansson |first1=B. |last2=Abuja |first2=R. |last3=Eriksson |first3=O. |last4=Wills |first4=J. M. |display-authors=3 |year=1995 |title=Anomalous fcc crystal structure of thorium metal. |journal=Physical Review Letters |volume=75 |issue=2 |pages=280–283 |doi=10.1103/PhysRevLett.75.280|pmid=10059654 |bibcode=1995PhRvL..75..280J|url=https://zenodo.org/record/1233903 }}</ref><ref name=XuPyykko>
{{cite journal |last1=Xu |first1=Wen-Hua |last2=Pyykkö |first2=Pekka |date=8 June 2016 |url=http://pubs.rsc.org/-/content/articlehtml/2016/cp/c6cp02706g |title=Is the chemistry of lawrencium peculiar |journal=Phys. Chem. Chem. Phys. |volume=2016 |issue=18 |pages=17351–5 |doi=10.1039/c6cp02706g |pmid=27314425 |access-date=24 April 2017|bibcode=2016PCCP...1817351X |hdl=10138/224395 |s2cid=31224634 |hdl-access=free }}</ref> For a very long time, the seventh row was incomplete as most of its elements do not occur in nature. The missing ] started to be synthesized in the laboratory in 1940, when neptunium was made.<ref name="Scerri354" /> (However, the first element to be discovered by synthesis rather than in nature was technetium in 1937.) The row was completed with the synthesis of ] in 2010<ref name="117s">{{cite journal |last1=Oganessian |first1=Yu.Ts. |author-link1=Yuri Oganessian |last2=Abdullin |first2=F.Sh. |last3=Bailey |first3=P.D. |last4=Benker |first4=D.E. |last5=Bennett |first5=M.E. |last6=Dmitriev |first6=S.N. |last7=Ezold |first7=J.G. |last8=Hamilton |first8=J.H. |last9=Henderson |first9=R.A. |first10=M.G. |last10=Itkis |first11=Yuri V. |last11=Lobanov |first12=A.N. |last12=Mezentsev |first13=K. J. |last13=Moody |first14=S.L. |last14=Nelson |first15=A.N. |last15=Polyakov |first16=C.E. |last16=Porter |first17=A.V. |last17=Ramayya |first18=F.D. |last18=Riley |first19=J.B. |last19=Roberto |first20=M. A. |last20=Ryabinin |first21=K.P. |last21=Rykaczewski |first22=R.N. |last22=Sagaidak |first23=D.A. |last23=Shaughnessy |first24=I.V. |last24=Shirokovsky |first25=M.A. |last25=Stoyer |first26=V.G. |last26=Subbotin |first27=R. |last27=Sudowe |first28=A.M. |last28=Sukhov |first29=Yu.S. |last29=Tsyganov |first30=Vladimir K. |last30=Utyonkov |first31=A.A. |last31=Voinov |first32=G.K. |last32=Vostokin |first33=P.A. |last33=Wilk |display-authors=6 |title=Synthesis of a new element with atomic number {{nowrap|''Z'' {{=}} 117}} |year=2010 |journal=Physical Review Letters |volume=104 |issue=14 |page=142502 |doi=10.1103/PhysRevLett.104.142502 |pmid=20481935 |bibcode=2010PhRvL.104n2502O |s2cid=3263480 |doi-access=free }}</ref> (the last element ] had already been made in 2002),<ref name="pp2002">{{cite journal|author=Oganessian, Yu. T.|display-authors=etal|title=Results from the first {{chem|249|Cf}}+{{chem|48|Ca}} experiment|url=https://www.jinr.ru/publish/Preprints/2002/287(D7-2002-287)e.pdf|journal=JINR Communication|date=2002|access-date=13 June 2009|archive-date=13 December 2004|archive-url=https://web.archive.org/web/20041213100709/https://www.jinr.ru/publish/Preprints/2002/287%28D7-2002-287%29e.pdf}}</ref> and the last elements in this seventh row were given names in 2016.<ref name="IUPAC-20161130">{{cite news |author=<!--Not stated--> |title=IUPAC Announces the Names of the Elements 113, 115, 117, and 118 |url=https://iupac.org/iupac-announces-the-names-of-the-elements-113-115-117-and-118/ |date=30 November 2016 |work=] |access-date=1 December 2016 |archive-date=30 November 2016 |archive-url=https://web.archive.org/web/20161130111959/https://iupac.org/iupac-announces-the-names-of-the-elements-113-115-117-and-118/ |url-status=live }}</ref>
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This completes the modern periodic table, with all seven rows completely filled to capacity.<ref name="IUPAC-20161130" /><!--when 8th row elements are discovered, replace them here and write "The eighth row finishes prematurely as we run out of elements discovered."-->

===Electron configuration table===
The following table shows the electron configuration of a neutral gas-phase atom of each element. Different configurations can be favoured in different chemical environments.<ref name="Jorgensen" /> The main-group elements have entirely regular electron configurations; the transition and inner transition elements show twenty irregularities due to the aforementioned competition between subshells close in energy level. For the last ten elements (109–118), experimental data is lacking<ref>{{cite journal |url=https://www.nist.gov/pml/periodic-table-elements |title=Periodic Table of the Elements |author=] (NIST) |date=August 2019 |journal=NIST |access-date=7 February 2021 |archive-date=8 February 2021 |archive-url=https://web.archive.org/web/20210208182536/https://www.nist.gov/pml/periodic-table-elements |url-status=live }}</ref> and therefore calculated configurations have been shown instead.<ref>{{cite journal |last1=Fricke |first1=B. |editor-last=Dunitz |editor-first=J. D. |year=1975 |journal=Structure and Bonding |volume=21 |pages=89–144 |title=Superheavy elements a prediction of their chemical and physical properties|publisher=Springer-Verlag |location=Berlin |doi=10.1007/BFb0116496|isbn=978-3-540-07109-9 }}</ref> Completely filled subshells have been greyed out.

{{Periodic table (electron configuration)}}


==== Blocks ==== == Variations ==
===Period 1===
{{Main|Block (periodic table)}}
{{Main|Period 1 element}}
Although the modern periodic table is standard today, the placement of the period 1 elements hydrogen and helium remains an open issue under discussion, and some variation can be found.<ref name=KW/><ref name="Lemonick">{{cite web |url=https://cen.acs.org/physical-chemistry/periodic-table/periodic-table-icon-chemists-still/97/i1 |title=The periodic table is an icon. But chemists still can't agree on how to arrange it |last=Lemonick |first=Sam |date=2019 |website=C&EN News |access-date=16 December 2020 |archive-date=28 January 2021 |archive-url=https://web.archive.org/web/20210128031450/https://cen.acs.org/physical-chemistry/periodic-table/periodic-table-icon-chemists-still/97/i1 |url-status=live }}</ref> Following their respective s<sup>1</sup> and s<sup>2</sup> electron configurations, hydrogen would be placed in group 1, and helium would be placed in group 2.<ref name="KW" /> The group 1 placement of hydrogen is common, but helium is almost always placed in group 18 with the other noble gases.<ref name="IUPAC-redbook" /> The debate has to do with conflicting understandings of the extent to which chemical or electronic properties should decide periodic table placement.<ref name=Lemonick/>


Like the group 1 metals, hydrogen has one electron in its outermost shell<ref name="Gray12">Gray, p. 12</ref> and typically loses its only electron in chemical reactions.<ref name="Vlasov" /> Hydrogen has some metal-like chemical properties, being able to displace some metals from their ].<ref name="Vlasov">{{cite book |last1=Vlasov |first1=L. |last2=Trifonov |first2=D. |translator-last1=Sobolev |translator-first1=D. |date=1970 |title=107 Stories About Chemistry |publisher=Mir Publishers |pages=23–27 |isbn=978-0-8285-5067-3}}</ref> But it forms a diatomic nonmetallic gas at standard conditions, unlike the alkali metals which are reactive solid metals. This and hydrogen's formation of ]s, in which it gains an electron, brings it close to the properties of the ]s which do the same<ref name=Vlasov/> (though it is rarer for hydrogen to form H<sup>−</sup> than H<sup>+</sup>).<ref name="raynercanham">{{cite book |last=Rayner-Canham |first=Geoffrey |date=2020 |title=The Periodic Table: Past, Present, Future |publisher=World Scientific |pages=53–70, 85–102 |isbn=978-981-12-1850-7}}</ref> Moreover, the lightest two halogens (] and ]) are gaseous like hydrogen at standard conditions.<ref name="Vlasov" /> Some properties of hydrogen are not a good fit for either group: hydrogen is neither highly oxidizing nor highly reducing and is not reactive with water.<ref name=raynercanham/> Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens, but matches neither group perfectly, and is thus difficult to place by its chemistry.<ref name="Vlasov" /> Therefore, while the electronic placement of hydrogen in group 1 predominates, some rarer arrangements show either hydrogen in group 17,<ref>{{Clayden}}</ref> duplicate hydrogen in both groups 1 and 17,<ref>{{cite journal |last=Seaborg |first= G.|title=The chemical and radioactive properties of the heavy elements |journal= Chemical & Engineering News|year=1945 |volume=23 |issue=23 |pages=2190–93|doi= 10.1021/cen-v023n023.p2190}}</ref><ref name="Kaesz" /> or float it separately from all groups.<ref name="Kaesz">{{cite journal |last1=Kaesz |first1=Herb |last2=Atkins |first2=Peter |date=2009 |title=A Central Position for Hydrogen in the Periodic Table |journal=Chemistry International |volume=25 |issue=6 |page=14 |doi=10.1515/ci.2003.25.6.14 |doi-access=free }}</ref><ref name="GE">Greenwood & Earnshaw, throughout the book</ref><ref name="KW" /> This last option has nonetheless been criticized by the chemist and philosopher of science ] on the grounds that it appears to imply that hydrogen is above the periodic law altogether, unlike all the other elements.<ref>{{cite journal |last1=Scerri |first1=Eric |date=2004 |title=The Placement of Hydrogen in the Periodic Table |url=http://publications.iupac.org/ci/2004/2603/ud2_scerri.html |journal=Chemistry International |volume=26 |issue=3 |pages=21–22 |doi=10.1515/ci.2004.26.3.21 |access-date=1 January 2023|doi-access=free }}</ref>
]s with the CAS (American Group Numbering System).]]


Helium is the only element that routinely occupies a position in the periodic table that is not consistent with its electronic structure. It has two electrons in its outermost shell, whereas the other noble gases have eight; and it is an s-block element, whereas all other noble gases are p-block elements. However it is unreactive at standard conditions, and has a full outer shell: these properties are like the noble gases in group 18, but not at all like the reactive alkaline earth metals of group 2. For these reasons helium is nearly universally placed in group 18<ref name="IUPAC-redbook" /> which its properties best match;<ref name="KW" /> a proposal to move helium to group 2 was rejected by IUPAC in 1988 for these reasons.<ref name=Fluck/> Nonetheless, helium is still occasionally placed in group 2 today,<ref name=shattered>{{cite book |last1=Thyssen |first1=Pieter |last2=Ceulemans |first2=Arnout |date=2017 |title=Shattered Symmetry: Group Theory from the Eightfold Way to the Periodic Table |url= |location= |publisher=Oxford University Press |pages=336, 360–381 |isbn=978-0-19-061139-2}}</ref> and some of its physical and chemical properties are closer to the group 2 elements and support the electronic placement.<ref name="Gray12" /><ref name="KW">{{cite book |last1=Keeler |first1=James |last2=Wothers |first2=Peter |date=2014 |title=Chemical Structure and Reactivity |url= |edition=2nd |location= |publisher=Oxford University Press |pages=257–260 |isbn=978-0-19-9604135}}</ref> Solid helium crystallises in a ] structure, which matches beryllium and magnesium in group 2, but not the other noble gases in group 18.<ref name=Kurushkin>{{cite journal |last1=Kurushkin |first1=Mikhail |date=2020 |title=Helium's placement in the Periodic Table from a crystal structure viewpoint |url=https://www.researchgate.net/publication/342152661 |journal=IUCrJ |volume=7 |issue=4 |pages=577–578 |doi=10.1107/S2052252520007769 |pmid=32695406 |pmc=7340260 |access-date=19 June 2020 |doi-access=free |bibcode=2020IUCrJ...7..577K |archive-date=19 October 2021 |archive-url=https://web.archive.org/web/20211019202250/https://www.researchgate.net/publication/342152661_Helium's_placement_in_the_Periodic_Table_from_a_crystal_structure_viewpoint |url-status=live }}</ref> Recent theoretical developments in noble gas chemistry, in which helium is expected to show slightly less inertness than neon and to form (HeO)(LiF)<sub>2</sub> with a structure similar to the analogous beryllium compound (but with no expected neon analogue), have resulted in more chemists advocating a placement of helium in group 2. This relates to the electronic argument, as the reason for neon's greater inertness is repulsion from its filled p-shell that helium lacks, though realistically it is unlikely that helium-containing molecules will be stable outside extreme low-temperature conditions (around 10&nbsp;]).<ref name="PTSS" /><ref name=grochala>{{cite journal |last1=Grochala |first1=Wojciech |date=1 November 2017 |title=On the position of helium and neon in the Periodic Table of Elements |journal=Foundations of Chemistry |volume=20 |pages=191–207 |issue=2018 |doi=10.1007/s10698-017-9302-7 |doi-access=free }}</ref><ref>{{cite journal |last1=Bent Weberg |first1=Libby |date=18 January 2019 |title="The" periodic table |url=https://cen.acs.org/articles/97/i3/Reactions.html |journal=Chemical & Engineering News |volume=97 |issue=3 |access-date=27 March 2020 |archive-date=1 February 2020 |archive-url=https://web.archive.org/web/20200201200009/https://cen.acs.org/articles/97/i3/Reactions.html |url-status=live }}</ref><ref>{{cite journal |last1=Grandinetti |first1=Felice |date=23 April 2013 |title=Neon behind the signs |journal=Nature Chemistry |volume=5 |issue=2013 |page=438 |doi=10.1038/nchem.1631 |pmid=23609097 |bibcode=2013NatCh...5..438G |doi-access=free }}</ref>
Because of the importance of the outermost electron shell, the different regions of the periodic table are sometimes referred to as ''periodic table blocks'', named according to the subshell in which the "last" electron resides.<ref name="Gray12"/> The ] comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. The ] comprises the last six groups which are groups 13 to 18 in IUPAC (3A&nbsp;to 8A in American) and contains, among others, all of the ]s. The d-block comprises groups 3 to 12 in IUPAC (or&nbsp;3B to 2B in American group numbering) and contains all of the transition metals. The f-block, usually offset below the rest of the periodic table, comprises the lanthanides and actinides.<ref>{{cite book |last=Jones |first=Chris |title=d- and f-block chemistry |year=2002 |publisher=J. Wiley & Sons |location=New York |isbn=978-0-471-22476-1|oclc=300468713 |page=2}}</ref>


The ] in the periodic table has additionally been cited to support moving helium to group 2. It arises because the first orbital of any type is unusually small, since unlike its higher analogues, it does not experience interelectronic repulsion from a smaller orbital of the same type. This makes the first row of elements in each block unusually small, and such elements tend to exhibit characteristic kinds of anomalies for their group. Some chemists arguing for the repositioning of helium have pointed out that helium exhibits these anomalies if it is placed in group 2, but not if it is placed in group 18: on the other hand, neon, which would be the first group 18 element if helium was removed from that spot, does exhibit those anomalies.<ref name="PTSS" /> The relationship between helium and beryllium is then argued to resemble that between hydrogen and lithium, a placement which is much more commonly accepted.<ref name=grochala/> For example, because of this trend in the sizes of orbitals, a large difference in atomic radii between the first and second members of each main group is seen in groups 1 and 13–17: it exists between neon and argon, and between helium and beryllium, but not between helium and neon. This similarly affects the noble gases' boiling points and solubilities in water, where helium is too close to neon, and the large difference characteristic between the first two elements of a group appears only between neon and argon. Moving helium to group 2 makes this trend consistent in groups 2 and 18 as well, by making helium the first group 2 element and neon the first group 18 element: both exhibit the characteristic properties of a ] first element of a group.<ref name=SB23/><ref>Siekierski and Burgess, p. 128</ref> The group 18 placement of helium nonetheless remains near-universal due to its extreme inertness.<ref>{{Cite book|title = Modeling Marvels: Computational Anticipation of Novel Molecules|url = https://books.google.com/books?id=IoFzgBSSCwEC|publisher = Springer Science & Business Media|date = 5 December 2008|isbn = 978-1-4020-6973-4|first = Errol G.|last = Lewars|pages = 69–71|url-status=live|archive-url = https://web.archive.org/web/20160519021952/https://books.google.com/books?id=IoFzgBSSCwEC|archive-date = 19 May 2016|df = dmy-all}}</ref> Additionally, tables that float both hydrogen and helium outside all groups may rarely be encountered.<ref name=GE/><ref name=KW/><ref name=jensenlaw/>
=== Variations and other conventions ===


===Group 3===
In presentations of the periodic table, the lanthanides and the actinides are customarily shown as two additional rows below the main body of the table,<ref>Gray, p. 11</ref> with placeholders or else a selected single element of each series (either ] or ], and either ] or ], respectively) shown in a single cell of the main table, between ] and ], and ] and ], respectively. This convention is entirely a matter of aesthetics and formatting practicality; a rarely used ] inserts the lanthanide and actinide series in their proper places, as parts of the table's sixth and seventh rows (periods).<ref name="Guenther">{{cite doi|10.1021/ed064p9}}</ref>
{{main|Group 3 element#Composition}}
{{Periodic table (micro)|mark=Sc,Y,Lu,Lr|title=Group&nbsp;3: Sc, Y, Lu, Lr ]|caption=Correct depiction of Group 3}}
{{Periodic table (micro)|form=Sc, Y, La, Ac|mark=Sc,Y,La,Ac|title=Group&nbsp;3: Sc, Y, La, Ac ]|caption=Incorrect depiction of Group 3}}
In many periodic tables, the f-block is shifted one element to the right, so that lanthanum and actinium become d-block elements in group 3, and Ce–Lu and Th–Lr form the f-block. Thus the d-block is split into two very uneven portions. This is a holdover from early mistaken measurements of electron configurations; modern measurements are more consistent with the form with lutetium and lawrencium in group 3, and with La–Yb and Ac–No as the f-block.<ref name="Jensen1982"/><ref name=wulfsberg53/>


The 4f shell is completely filled at ytterbium, and for that reason ] and ] in 1948 considered it incorrect to group lutetium as an f-block element.<ref name=Landau/> They did not yet take the step of removing lanthanum from the d-block as well, but ] realized in 1963 that lanthanum's low-temperature ] implied the activity of its 4f shell.<ref name=Kondo/> In 1965, David C. Hamilton linked this observation to its position in the periodic table, and argued that the f-block should be composed of the elements La–Yb and Ac–No.<ref name=Hamilton/> Since then, physical, chemical, and electronic evidence has supported this assignment.<ref name=Jensen1982/><ref name=Fluck/><ref name=wulfsberg53>Wulfsberg, p. 53: "As pointed out by W. B. Jensen, the metallurgical resemblance is much stronger for lutetium than for lanthanum, so we have adopted the metallurgist's convention of listing Lu (and by extension Lr) below Sc and Y. An important additional advantage of this is that the periodic table becomes more symmetrical, and it becomes easier to predict electron configurations. E. R. Scerri points out that recent determinations of the electron configurations of most of the ''f''-block elements now are more compatible with this placement of Lu and Lr."</ref> The issue was brought to wide attention by ] in 1982,<ref name=Jensen1982/> and the reassignment of lutetium and lawrencium to group 3 was supported by IUPAC reports dating from 1988 (when the 1–18 group numbers were recommended)<ref name="Fluck"/> and 2021.<ref name=2021IUPAC/> The variation nonetheless still exists because most textbook writers are not aware of the issue.<ref name=Jensen1982/>
]


A third form can sometimes be encountered in which the spaces below yttrium in group 3 are left empty, such as the table appearing on the IUPAC web site,<ref name="IUPAC-redbook" /> but this creates an inconsistency with quantum mechanics by making the f-block 15 elements wide (La–Lu and Ac–Lr) even though only 14 electrons can fit in an f-subshell.<ref name=2021IUPAC/> There is moreover some confusion in the literature on which elements are then implied to be in group 3.<ref name=2021IUPAC/><ref name=Thyssen/><ref name="JWP">{{cite journal |author=Barber, Robert C. |author2=Karol, Paul J |author3=Nakahara, Hiromichi |author4=Vardaci, Emanuele |author5=Vogt, Erich W. |title=Discovery of the elements with atomic numbers greater than or equal to 113 (IUPAC Technical Report) |doi=10.1351/PAC-REP-10-05-01 |journal=Pure Appl. Chem. |date=2011 |volume=83 |issue=7 |page=1485|doi-access=free }}</ref><ref name="Karol">{{cite journal |last1=Karol |first1=Paul J. |last2=Barber |first2=Robert C. |last3=Sherrill |first3=Bradley M. |last4=Vardaci |first4=Emanuele |last5=Yamazaki |first5=Toshimitsu |date=22 December 2015 |title=Discovery of the elements with atomic numbers Z = 113, 115 and 117 (IUPAC Technical Report) |journal=Pure Appl. Chem. |volume=88 |issue=1–2 |pages=139–153 |doi=10.1515/pac-2015-0502|doi-access=free }}</ref><ref>{{cite journal |last1=Pyykkö |first1=Pekka |date=2019 |title=An essay on periodic tables |url=http://www.chem.helsinki.fi/~pyykko/pekka/No330b.pdf |journal=Pure and Applied Chemistry |volume=91 |issue=12 |pages=1959–1967 |doi=10.1515/pac-2019-0801 |s2cid=203944816 |access-date=27 November 2022}}</ref> While the 2021 IUPAC report noted that 15-element-wide f-blocks are supported by some practitioners of a specialized branch of ] focusing on the properties of ]s, the project's opinion was that such interest-dependent concerns should not have any bearing on how the periodic table is presented to "the general chemical and scientific community".<ref name=2021IUPAC/> Other authors focusing on superheavy elements since clarified that the "15th entry of the f-block represents the first slot of the d-block which is left vacant to indicate the place of the f-block inserts", which would imply that this form still has lutetium and lawrencium (the 15th entries in question) as d-block elements in group 3.<ref name=smits/> Indeed, when IUPAC publications expand the table to 32 columns, they make this clear and place lutetium and lawrencium under yttrium in group 3.<ref>{{cite journal |last1=Leigh |first1=G. Jeffrey |date=2009 |title=Periodic Tables and IUPAC |url=https://publications.iupac.org/ci/2009/3101/1_leigh.html |journal=Chemistry International |volume=31 |issue=1 |pages=4–6 |doi=10.1515/ci.2009.31.1.4 |access-date=8 January 2024}}</ref><ref>{{cite book |editor-last=Leigh |editor-first=G. Jeffrey |date=1990 |title=Nomenclature of inorganic chemistry : recommendations 1990 |url=https://archive.org/details/nomenclatureofin0000unse/page/282/mode/2up |location= |publisher=Blackwell Scientific Publications |page=283 |isbn=0-632-02319-8}}</ref>
Many presentations of the periodic table show a dark stair-step diagonal line along the ]s, with metals to the left of the line and non-metals to the right.<ref name="Guenther"/> Various other groupings of the chemical elements are sometimes also highlighted on a periodic table, such as ]s, ]s, and metalloids. Other informal groupings of the elements exist, such as the ] and the ], but are rarely addressed in periodic tables.<ref name="Guenther"/>


Several arguments in favour of Sc-Y-La-Ac can be encountered in the literature,<ref>{{cite journal |last1= Vernon|first1= R|date= 2021|title=The location and composition of Group 3 of the periodic table|journal=Foundations of Chemistry |volume= 23|issue= 2|pages= 155–197|doi=10.1007/s10698-020-09384-2|s2cid= 254501533|doi-access= free}}</ref><ref>{{cite journal |last1=Cotton |first1=SA |last2=Raithby |first2=BR |last3=Shield |first3=A|date= 2022|title= A comparison of the structural chemistry of scandium, yttrium, lanthanum and lutetium: A contribution to the group 3 debate |journal= Coordination Chemistry Reviews |volume=455 |issue= |page= 214366 |doi= 10.1016/j.ccr.2021.214366|s2cid=245712597 |url=https://purehost.bath.ac.uk/ws/files/227604162/CCR_SC_Y_Ln_Manuscript_accepted_131221.pdf }}</ref> but they have been challenged as being logically inconsistent.<ref name=Jensen2015/><ref name=Scerri2009/><ref name=Chemey/> For example, it has been argued that lanthanum and actinium cannot be f-block elements because as individual gas-phase atoms, they have not begun to fill the f-subshells.<ref name=Lavelle>{{cite journal |last1=Lavelle |first1=Laurence |date=2008 |title=Lanthanum (La) and Actinium (Ac) Should Remain in the d-block |journal=Journal of Chemical Education |volume=85 |issue=11 |pages=1482–1483 |doi=10.1021/ed085p1482|bibcode=2008JChEd..85.1482L |doi-access=free }}</ref> But the same is true of thorium which is never disputed as an f-block element,<ref name=2021IUPAC/><ref name=Jensen1982/> and this argument overlooks the problem on the other end: that the f-shells complete filling at ytterbium and nobelium, matching the Sc-Y-Lu-Lr form, and not at lutetium and lawrencium as the Sc-Y-La-Ac form would have it.<ref name=johnson>{{cite book |last=Johnson |first=David |date=1984 |title=The Periodic Law |url=https://www.rsc.org/images/23_The_Periodic_Law_tcm18-30005.pdf |location= |publisher=The Royal Society of Chemistry |page= |isbn=0-85186-428-7}}</ref> Not only are such exceptional configurations in the minority,<ref name=johnson/> but they have also in any case never been considered as relevant for positioning any other elements on the periodic table: in gaseous atoms, the d-shells complete their filling at copper, palladium, and gold, but it is universally accepted by chemists that these configurations are exceptional and that the d-block really ends in accordance with the Madelung rule at zinc, cadmium, and mercury.<ref name="Thyssen"/> The relevant fact for placement<ref name=Jensen2009/><ref name=JensenLr/> is that lanthanum and actinium (like thorium) have valence f-orbitals that can become occupied in chemical environments, whereas lutetium and lawrencium do not:<ref name=jensenlaw/><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? |series=Advances in Solid State Physics |volume=13 |location=Berlin, Heidelberg |publisher=Springer |pages=375–396 |isbn=978-3-528-08019-8 |doi=10.1007/BFb0108579}}</ref><ref name=XuPyykko/> their f-shells are in the core, and cannot be used for chemical reactions.<ref name=Cp3Ln/><ref name=wulfsberg26/> Thus the relationship between yttrium and lanthanum is only a secondary relationship between elements with the same number of valence electrons but different kinds of valence orbitals, such as that between chromium and uranium; whereas the relationship between yttrium and lutetium is primary, sharing both valence electron count and valence orbital type.<ref name=jensenlaw/>
{{clear}}


== Periodic trends == == Periodic trends ==
{{Main|Periodic trends}} {{Main|Periodic trends}}
As chemical reactions involve the valence electrons,<ref name="cartoon" /> elements with similar outer electron configurations may be expected to react similarly and form compounds with similar proportions of elements in them.<ref name="Greenwood27" /> Such elements are placed in the same group, and thus there tend to be clear similarities and trends in chemical behaviour as one proceeds down a group.<ref>{{cite book |last=Messler|first=R. W.|title=The essence of materials for engineers|year=2010|publisher=Jones & Bartlett Publishers|location=Sudbury, MA|isbn=978-0-7637-7833-0|page=32}}</ref> As analogous configurations occur at regular intervals, the properties of the elements thus exhibit periodic recurrences, hence the name of the periodic table and the periodic law. These periodic recurrences were noticed well before the underlying theory that explains them was developed.<ref name="Myers">{{cite book |last=Myers|first=R.|title=The basics of chemistry|url=https://archive.org/details/basicschemistry00myer_641|url-access=limited|year=2003|pages=–67|publisher=Greenwood Publishing Group|location=Westport, CT|isbn=978-0-313-31664-7}}</ref><ref name="chang2">{{cite book|last=Chang|first=R.|title=Chemistry|url=https://archive.org/details/riimchemistry00chan/page/289|url-access=registration|year=2002|publisher=McGraw-Hill|location=New York|edition=7|isbn=978-0-07-112072-2|pages=}}</ref>


=== Atomic radius ===
]
Historically, the physical size of atoms was unknown until the early 20th century. The first calculated estimate of the atomic radius of hydrogen was published by physicist ] in 1910 to within an order of magnitude (a factor of 10) of the accepted value, the ] (~0.529 Å). In his model, Haas used a single-electron configuration based on the classical atomic model proposed by ] in 1904, often called the ].<ref>Haas, Arthur Erich (1884–1941) Uber die elektrodynamische Bedeutung des Planckschen Strahlungsgesetzes und uber eine neue Bestimmung des elektrischen Elementarquantums und der dimension des wasserstoffatoms. Sitzungsberichte der kaiserlichen Akademie der Wissenschaften in Wien. 2a, 119 pp 119–144 (1910). Haas AE. Die Entwicklungsgeschichte des Satzes von der Erhaltung der Kraft. Habilitation Thesis, Vienna, 1909. Hermann, A. Arthur Erich Haas, Der erste Quantenansatz für das Atom. Stuttgart, 1965 </ref>


] (the size of atoms) are dependent on the sizes of their outermost orbitals.<ref name=SB23>Siekierski and Burgess, pp. 23–26</ref> They generally decrease going left to right along the main-group elements, because the nuclear charge increases but the outer electrons are still in the same shell. However, going down a column, the radii generally increase, because the outermost electrons are in higher shells that are thus further away from the nucleus.<ref name="cartoon" /><ref name="chemguidear">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/atradius.html |title=Atomic and Ionic Radius |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=14 November 2020 |archive-url=https://web.archive.org/web/20201114002613/https://www.chemguide.co.uk/atoms/properties/atradius.html |url-status=live }}</ref> The first row of each block is abnormally small, due to an effect called ] or primogenic repulsion:<ref>{{cite journal |last1=Cao |first1=Chang-Su |last2=Hu |first2=Han-Shi |last3=Li |first3=Jun |last4=Schwarz |first4=W. H. Eugen |date=2019 |title=Physical origin of chemical periodicities in the system of elements |journal=Pure and Applied Chemistry |volume=91 |issue=12 |pages=1969–1999 |doi=10.1515/pac-2019-0901 |s2cid=208868546 |doi-access=free }}</ref> the 1s, 2p, 3d, and 4f subshells have no inner analogues. For example, the 2p orbitals do not experience strong repulsion from the 1s and 2s orbitals, which have quite different angular charge distributions, and hence are not very large; but the 3p orbitals experience strong repulsion from the 2p orbitals, which have similar angular charge distributions. Thus higher s-, p-, d-, and f-subshells experience strong repulsion from their inner analogues, which have approximately the same angular distribution of charge, and must expand to avoid this. This makes significant differences arise between the small 2p elements, which prefer ]ing, and the larger 3p and higher p-elements, which do not.<ref name=SB23/> Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical {{Not a typo|5g}} elements:<ref name="Kaupp">{{cite journal |last=Kaupp |first=Martin |date=1 December 2006 |title=The role of radial nodes of atomic orbitals for chemical bonding and the periodic table |journal=Journal of Computational Chemistry |volume=28 |issue=1 |pages=320–25 |doi=10.1002/jcc.20522 |pmid=17143872 |s2cid=12677737 |doi-access=free }}</ref> the degree of this first-row anomaly is highest for the s-block, is moderate for the p-block, and is less pronounced for the d- and f-blocks.<ref name="PTSS2" />
The primary determinant of an element's chemical properties is its electron configuration, particularly the valence shell electrons. For instance, any atoms with four valence electrons occupying p orbitals will exhibit some similarity. The type of orbital in which the atom's outermost electrons reside determines the "block" to which it belongs. The number of valence shell electrons determines the family, or group, to which the element belongs.<ref name="Gray12"/>
The total number of electron shells an atom has determines the period to which it belongs. Each shell is divided into different subshells, which as atomic number increases are filled roughly in the order depicted in the table at hand (according to the ]; see table below).<ref>Moore, p. 46</ref> Hence the structure of the periodic table. Since the outermost electrons determine chemical properties, those with the same number of valence electrons are generally grouped together.<ref>{{cite doi|10.1063/1.432900}}</ref>


In the transition elements, an inner shell is filling, but the size of the atom is still determined by the outer electrons. The increasing nuclear charge across the series and the increased number of inner electrons for shielding somewhat compensate each other, so the decrease in radius is smaller.<ref name="chemguidear" /> The 4p and 5d atoms, coming immediately after new types of transition series are first introduced, are smaller than would have been expected,<ref name="Greenwood29">Greenwood and Earnshaw, p. 29</ref> because the added core 3d and 4f subshells provide only incomplete shielding of the nuclear charge for the outer electrons. Hence for example gallium atoms are slightly smaller than aluminium atoms.<ref name=SB23/> Together with kainosymmetry, this results in an even-odd difference between the periods (except in the s-block){{efn|Properties of the p-block elements nevertheless do affect the succeeding s-block elements. The 3s shell in sodium is above a kainosymmetric 2p core, but the 4s shell in potassium is above the much larger 3p core. Hence while one would have already expected potassium atoms to be larger than sodium atoms, the size difference is greater than usual.<ref name=SB23/>}} that is sometimes known as secondary periodicity: elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods (except the first) differ in the opposite direction. Thus for example many properties in the p-block show a zigzag rather than a smooth trend along the group. For example, phosphorus and antimony in odd periods of group 15 readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth in even periods prefer to stay at +3.<ref name="PTSS2" /><ref>{{cite journal |last1=Imyanitov |first1=Naum S. |date=2018 |title=Is the periodic table appears doubled? Two variants of division of elements into two subsets. Internal and secondary periodicity |url= |journal=Foundations of Chemistry |volume=21 |issue= |pages=255–284 |doi=10.1007/s10698-018-9321-z |s2cid=254514910 |access-date=}}</ref> A similar situation holds for the d-block, with lutetium through tungsten atoms being slightly smaller than yttrium through molybdenum atoms respectively.<ref>{{cite journal |last1=Chistyakov |first1=V. M. |date=1968 |title=Biron's Secondary Periodicity of the Side d-subgroups of Mendeleev's Short Table |url=https://archive.org/details/sim_russian-journal-of-general-chemistry_1968-02_38_2/page/212/mode/2up |journal=Journal of General Chemistry of the USSR |volume=38 |issue=2 |pages=213–214 |doi= |access-date=6 January 2024}}</ref><ref name="Calc1">{{cite journal|author1=P. Pyykkö|author2=M. Atsumi|year=2009|title=Molecular Single-Bond Covalent Radii for Elements 1-118|journal=Chemistry: A European Journal|volume=15|issue=1|pages=186–197|doi=10.1002/chem.200800987|pmid=19058281}}</ref>
{| class="wikitable" style="margin:auto; float:left; margin:5px;"

]
Thallium and lead atoms are about the same size as indium and tin atoms respectively, but from bismuth to radon the 6p atoms are larger than the analogous 5p atoms. This happens because when atomic nuclei become highly charged, ] becomes needed to gauge the effect of the nucleus on the electron cloud. These ] result in heavy elements increasingly having differing properties compared to their lighter homologues in the periodic table. ] splits the p-subshell: one p-orbital is relativistically stabilized and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilized and expanded.<ref name=SB23/> Relativistic effects also explain why ] is golden and ] is a liquid at room temperature.<ref name="PekkaPyykko">{{cite journal |doi=10.1021/ar50140a002 |title=Relativity and the periodic system of elements |year=1979 |last1=Pyykkö |first1=Pekka |last2=Desclaux |first2=Jean Paul |journal=Accounts of Chemical Research |volume=12 |issue=8 |page=276}}</ref><ref name="Norrby">{{cite journal |doi=10.1021/ed068p110 |title=Why is mercury liquid? Or, why do relativistic effects not get into chemistry textbooks? |year=1991 |last1=Norrby |first1=Lars J. |journal=Journal of Chemical Education |volume=68 |issue=2 |page=110 |bibcode = 1991JChEd..68..110N}}</ref> They are expected to become very strong in the late seventh period, potentially leading to a collapse of periodicity.<ref name=actrev/> Electron configurations are only clearly known until element 108 (]), and experimental chemistry beyond 108 has only been done for elements 112 (]) through 115 (]), so the chemical characterization of the heaviest elements remains a topic of current research.<ref name="Schändel 2003 277">{{cite book|title=The Chemistry of Superheavy Elements|last=Schädel|first=M.|year=2003|publisher=Kluwer Academic Publishers|location=Dordrecht|isbn=978-1-4020-1250-1|page=277}}</ref><ref name=moscovium>{{cite journal |last1=Yakushev |first1=A. |last2=Khuyagbaatar |first2=J. |first3=Ch. E. |last3=Düllmann |first4=M. |last4=Block |first5=R. A. |last5=Cantemir |first6=D. M. |last6=Cox |first7=D. |last7=Dietzel |first8=F. |last8=Giacoppo |first9=Y. |last9=Hrabar |first10=M. |last10=Iliaš |first11=E. |last11=Jäger |first12=J. |last12=Krier |first13=D. |last13=Krupp |first14=N. |last14=Kurz |first15=L. |last15=Lens |first16=S. |last16=Löchner |first17=Ch. |last17=Mokry |first18=P. |last18=Mošať |first19=V. |last19=Pershina |first20=S. |last20=Raeder |first21=D. |last21=Rudolph |first22=J. |last22=Runke |first23=L. G. |last23=Sarmiento |first24=B. |last24=Schausten |first25=U. |last25=Scherer |first26=P. |last26=Thörle-Pospiesch |first27=N. |last27=Trautmann |first28=M. |last28=Wegrzecki |first29=P. |last29=Wieczorek |date=23 September 2024 |title=Manifestation of relativistic effects in the chemical properties of nihonium and moscovium revealed by gas chromatography studies |journal=Frontiers in Chemistry |volume=12 |issue= |pages= |doi=10.3389/fchem.2024.1474820 |doi-access=free |pmid=39391836 |pmc=11464923 |bibcode=2024FrCh...1274820Y }}</ref>

The trend that atomic radii decrease from left to right is also present in ], though it is more difficult to examine because the most common ions of consecutive elements normally differ in charge. Ions with the same electron configuration decrease in size as their atomic number rises, due to increased attraction from the more positively charged nucleus: thus for example ionic radii decrease in the series Se<sup>2−</sup>, Br<sup>−</sup>, Rb<sup>+</sup>, Sr<sup>2+</sup>, Y<sup>3+</sup>, Zr<sup>4+</sup>, Nb<sup>5+</sup>, Mo<sup>6+</sup>, Tc<sup>7+</sup>. Ions of the same element get smaller as more electrons are removed, because the attraction from the nucleus begins to outweigh the repulsion between electrons that causes electron clouds to expand: thus for example ionic radii decrease in the series V<sup>2+</sup>, V<sup>3+</sup>, V<sup>4+</sup>, V<sup>5+</sup>.<ref>Wulfsberg, pp. 33–34</ref>

=== Ionisation energy ===
]
The first ] of an atom is the energy required to remove an electron from it. This varies with the atomic radius: ionisation energy increases left to right and down to up, because electrons that are closer to the nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimized at the first element of each period – hydrogen and the ]s – and then generally rises until it reaches the ] at the right edge of the period.<ref name="cartoon" /> There are some exceptions to this trend, such as oxygen, where the electron being removed is paired and thus interelectronic repulsion makes it easier to remove than expected.<ref name="Greenwood294">Greenwood and Earnshaw, pp. 24–5</ref>

In the transition series, the outer electrons are preferentially lost even though the inner orbitals are filling. For example, in the 3d series, the 4s electrons are lost first even though the 3d orbitals are being filled. The shielding effect of adding an extra 3d electron approximately compensates the rise in nuclear charge, and therefore the ionisation energies stay mostly constant, though there is a small increase especially at the end of each transition series.<ref name="chemguideIE">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/ies.html |title=Ionisation Energy |last=Clark |first=Jim |date=2016 |website=Chemguide |access-date=30 March 2021 |archive-date=22 April 2021 |archive-url=https://web.archive.org/web/20210422032340/https://www.chemguide.co.uk/atoms/properties/ies.html |url-status=live }}</ref>

As metal atoms tend to lose electrons in chemical reactions, ionisation energy is generally correlated with chemical reactivity, although there are other factors involved as well.<ref name="chemguideIE" />

=== Electron affinity ===
]
The opposite property to ionisation energy is the ], which is the energy released when adding an electron to the atom.<ref name="chemguideea" /> A passing electron will be more readily attracted to an atom if it feels the pull of the nucleus more strongly, and especially if there is an available partially filled outer orbital that can accommodate it. Therefore, electron affinity tends to increase down to up and left to right. The exception is the last column, the noble gases, which have a full shell and have no room for another electron. This gives the ]s in the next-to-last column the highest electron affinities.<ref name="cartoon" />

Some atoms, like the noble gases, have no electron affinity: they cannot form stable gas-phase anions.<ref>{{cite journal |last1=Cárdenas |first1=Carlos |last2=Ayers |first2=Paul |first3=Frank |last3=De Proft |first4=David J. |last4=Tozer |first5=Paul |last5=Geerlings |date=2010 |title=Should negative electron affinities be used for evaluating the chemical hardness? |journal=Physical Chemistry Chemical Physics |volume=13 |issue=6 |pages=2285–2293 |doi=10.1039/C0CP01785J|pmid=21113528 }}</ref> (They can form metastable ] if the incoming electron arrives with enough kinetic energy, but these inevitably and rapidly ]: for example, the lifetime of the most long-lived He<sup>−</sup> level is about 359&nbsp;microseconds.)<ref>{{cite journal |last1=Schmidt |first1=H. T. |last2=Reinhed |first2=P. |first3=A. |last3=Orbán |first4=S. |last4=Rosén |first5=R. D. |last5=Thomas |first6=H. A. B. |last6=Johansson |first7=J. |last7=Werner |first8=D. |last8=Misra |first9=M. |last9=Björkhage |first10=L. |last10=Brännholm |first11=P. |last11=Löfgren |first12=L. |last12=Liljeby |first13=H. |last13=Cederquist |date=2012 |title=The lifetime of the helium anion |journal=Journal of Physics: Conference Series |volume=388 |issue= 1|pages=012006 |doi=10.1088/1742-6596/388/1/012006 |doi-access=free |bibcode=2012JPhCS.388a2006S }}</ref> The noble gases, having high ionisation energies and no electron affinity, have little inclination towards gaining or losing electrons and are generally unreactive.<ref name="cartoon" />

Some exceptions to the trends occur: oxygen and fluorine have lower electron affinities than their heavier homologues sulfur and chlorine, because they are small atoms and hence the newly added electron would experience significant repulsion from the already present ones. For the nonmetallic elements, electron affinity likewise somewhat correlates with reactivity, but not perfectly since other factors are involved. For example, fluorine has a lower electron affinity than chlorine (because of extreme interelectronic repulsion for the very small fluorine atom), but is more reactive.<ref name="chemguideea">{{cite web |url=https://www.chemguide.co.uk/atoms/properties/eas.html |title=Electron Affinity |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=23 April 2021 |archive-url=https://web.archive.org/web/20210423195854/https://www.chemguide.co.uk/atoms/properties/eas.html |url-status=live }}</ref>

===Valence and oxidation states===
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The ] of an element can be defined either as the number of hydrogen atoms that can combine with it to form a simple binary hydride, or as twice the number of oxygen atoms that can combine with it to form a simple binary oxide (that is, not a ] or a ]).<ref name=johnson/> The valences of the main-group elements are directly related to the group number: the hydrides in the main groups 1–2 and 13–17 follow the formulae MH, MH<sub>2</sub>, MH<sub>3</sub>, MH<sub>4</sub>, MH<sub>3</sub>, MH<sub>2</sub>, and finally MH. The highest oxides instead increase in valence, following the formulae M<sub>2</sub>O, MO, M<sub>2</sub>O<sub>3</sub>, MO<sub>2</sub>, M<sub>2</sub>O<sub>5</sub>, MO<sub>3</sub>, M<sub>2</sub>O<sub>7</sub>.{{efn|There are many lower oxides as well: for example, ] in group 15 forms two oxides, ] and ].<ref name="Greenwood27">Greenwood and Earnshaw, pp. 27–9</ref>}} Today the notion of valence has been extended by that of the ], which is the formal charge left on an element when all other elements in a compound have been removed as their ions.<ref name="Greenwood27" />

The electron configuration suggests a ready explanation from the number of electrons available for bonding;<ref name="Greenwood27" /> indeed, the number of valence electrons starts at 1 in group 1, and then increases towards the right side of the periodic table, only resetting at 3 whenever each new block starts. Thus in period 6, Cs–Ba have 1–2 valence electrons; La–Yb have 3–16; Lu–Hg have 3–12; and Tl–Rn have 3–8.<ref name=wulfsberg26>Wulfsberg, p. 26</ref> However, towards the right side of the d- and f-blocks, the theoretical maximum corresponding to using all valence electrons is not achievable at all;<ref>Wulfsberg, p. 28</ref> the same situation affects oxygen, fluorine, and the light noble gases up to krypton.<ref>Wulfsberg, p. 274</ref>
{| class="wikitable" style="margin:auto;text-align:center;"
|+ Number of valence electrons
!
! ]
! ]
! colspan=14 |
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
! ]
|- |-
! ]
! Subshell!!s!!f!!d!!p
| bgcolor="{{element color|s-block}}" | H<br />1
| colspan=30 style="border-width:0" |
| bgcolor="{{element color|s-block}}" | He<br />2
|- |-
! Period || || || || ! ]
| bgcolor="{{element color|s-block}}" | Li<br />1
| bgcolor="{{element color|s-block}}" | Be<br />2
| colspan=24 style="border-width:0" |
| bgcolor="{{element color|p-block}}" | B<br />3
| bgcolor="{{element color|p-block}}" | C<br />4
| bgcolor="{{element color|p-block}}" | N<br />5
| bgcolor="{{element color|p-block}}" | O<br />6
| bgcolor="{{element color|p-block}}" | F<br />7
| bgcolor="{{element color|p-block}}" | Ne<br />8
|- |-
! ]
! 1
| bgcolor="{{element color|s-block}}" | Na<br />1
| 1s || || ||
| bgcolor="{{element color|s-block}}" | Mg<br />2
| colspan=24 style="border-width:0" |
| bgcolor="{{element color|p-block}}" | Al<br />3
| bgcolor="{{element color|p-block}}" | Si<br />4
| bgcolor="{{element color|p-block}}" | P<br />5
| bgcolor="{{element color|p-block}}" | S<br />6
| bgcolor="{{element color|p-block}}" | Cl<br />7
| bgcolor="{{element color|p-block}}" | Ar<br />8
|- |-
! ]
! 2
| bgcolor="{{element color|s-block}}" | K<br />1
| 2s || || || 2p
| bgcolor="{{element color|s-block}}" | Ca<br />2
| colspan=14 style="border-width:0" |
| bgcolor="{{element color|d-block}}" | Sc<br />3
| bgcolor="{{element color|d-block}}" | Ti<br />4
| bgcolor="{{element color|d-block}}" | V<br />5
| bgcolor="{{element color|d-block}}" | Cr<br />6
| bgcolor="{{element color|d-block}}" | Mn<br />7
| bgcolor="{{element color|d-block}}" | Fe<br />8
| bgcolor="{{element color|d-block}}" | Co<br />9
| bgcolor="{{element color|d-block}}" | Ni<br />10
| bgcolor="{{element color|d-block}}" | Cu<br />11
| bgcolor="{{element color|d-block}}" | Zn<br />12
| bgcolor="{{element color|p-block}}" | Ga<br />3
| bgcolor="{{element color|p-block}}" | Ge<br />4
| bgcolor="{{element color|p-block}}" | As<br />5
| bgcolor="{{element color|p-block}}" | Se<br />6
| bgcolor="{{element color|p-block}}" | Br<br />7
| bgcolor="{{element color|p-block}}" | Kr<br />8
|- |-
! ]
! 3
| bgcolor="{{element color|s-block}}" | Rb<br />1
| 3s || || || 3p
| bgcolor="{{element color|s-block}}" | Sr<br />2
| colspan=14 style="border-width:0" |
| bgcolor="{{element color|d-block}}" | Y<br />3
| bgcolor="{{element color|d-block}}" | Zr<br />4
| bgcolor="{{element color|d-block}}" | Nb<br />5
| bgcolor="{{element color|d-block}}" | Mo<br />6
| bgcolor="{{element color|d-block}}" | Tc<br />7
| bgcolor="{{element color|d-block}}" | Ru<br />8
| bgcolor="{{element color|d-block}}" | Rh<br />9
| bgcolor="{{element color|d-block}}" | Pd<br />10
| bgcolor="{{element color|d-block}}" | Ag<br />11
| bgcolor="{{element color|d-block}}" | Cd<br />12
| bgcolor="{{element color|p-block}}" | In<br />3
| bgcolor="{{element color|p-block}}" | Sn<br />4
| bgcolor="{{element color|p-block}}" | Sb<br />5
| bgcolor="{{element color|p-block}}" | Te<br />6
| bgcolor="{{element color|p-block}}" | I<br />7
| bgcolor="{{element color|p-block}}" | Xe<br />8
|- |-
! ]
! 4
| bgcolor="{{element color|s-block}}" | Cs<br />1
| 4s || || 3d || 4p
| bgcolor="{{element color|s-block}}" | Ba<br />2
| bgcolor="{{element color|f-block}}" | La<br />3
| bgcolor="{{element color|f-block}}" | Ce<br />4
| bgcolor="{{element color|f-block}}" | Pr<br />5
| bgcolor="{{element color|f-block}}" | Nd<br />6
| bgcolor="{{element color|f-block}}" | Pm<br />7
| bgcolor="{{element color|f-block}}" | Sm<br />8
| bgcolor="{{element color|f-block}}" | Eu<br />9
| bgcolor="{{element color|f-block}}" | Gd<br />10
| bgcolor="{{element color|f-block}}" | Tb<br />11
| bgcolor="{{element color|f-block}}" | Dy<br />12
| bgcolor="{{element color|f-block}}" | Ho<br />13
| bgcolor="{{element color|f-block}}" | Er<br />14
| bgcolor="{{element color|f-block}}" | Tm<br />15
| bgcolor="{{element color|f-block}}" | Yb<br />16
| bgcolor="{{element color|d-block}}" | Lu<br />3
| bgcolor="{{element color|d-block}}" | Hf<br />4
| bgcolor="{{element color|d-block}}" | Ta<br />5
| bgcolor="{{element color|d-block}}" | W<br />6
| bgcolor="{{element color|d-block}}" | Re<br />7
| bgcolor="{{element color|d-block}}" | Os<br />8
| bgcolor="{{element color|d-block}}" | Ir<br />9
| bgcolor="{{element color|d-block}}" | Pt<br />10
| bgcolor="{{element color|d-block}}" | Au<br />11
| bgcolor="{{element color|d-block}}" | Hg<br />12
| bgcolor="{{element color|p-block}}" | Tl<br />3
| bgcolor="{{element color|p-block}}" | Pb<br />4
| bgcolor="{{element color|p-block}}" | Bi<br />5
| bgcolor="{{element color|p-block}}" | Po<br />6
| bgcolor="{{element color|p-block}}" | At<br />7
| bgcolor="{{element color|p-block}}" | Rn<br />8
|- |-
! ]
! 5
| bgcolor="{{element color|s-block}}" | Fr<br />1
| 5s || || 4d || 5p
| bgcolor="{{element color|s-block}}" | Ra<br />2
|-
| bgcolor="{{element color|f-block}}" | Ac<br />3
! 6
| bgcolor="{{element color|f-block}}" | Th<br />4
| 6s || 4f || 5d || 6p
| bgcolor="{{element color|f-block}}" | Pa<br />5
|-
| bgcolor="{{element color|f-block}}" | U<br />6
! 7
| bgcolor="{{element color|f-block}}" | Np<br />7
| 7s || 5f || 6d || 7p
| bgcolor="{{element color|f-block}}" | Pu<br />8
| bgcolor="{{element color|f-block}}" | Am<br />9
| bgcolor="{{element color|f-block}}" | Cm<br />10
| bgcolor="{{element color|f-block}}" | Bk<br />11
| bgcolor="{{element color|f-block}}" | Cf<br />12
| bgcolor="{{element color|f-block}}" | Es<br />13
| bgcolor="{{element color|f-block}}" | Fm<br />14
| bgcolor="{{element color|f-block}}" | Md<br />15
| bgcolor="{{element color|f-block}}" | No<br />16
| bgcolor="{{element color|d-block}}" | Lr<br />3
| bgcolor="{{element color|d-block}}" | Rf<br />4
| bgcolor="{{element color|d-block}}" | Db<br />5
| bgcolor="{{element color|d-block}}" | Sg<br />6
| bgcolor="{{element color|d-block}}" | Bh<br />7
| bgcolor="{{element color|d-block}}" | Hs<br />8
| bgcolor="{{element color|d-block}}" | Mt<br />9
| bgcolor="{{element color|d-block}}" | Ds<br />10
| bgcolor="{{element color|d-block}}" | Rg<br />11
| bgcolor="{{element color|d-block}}" | Cn<br />12
| bgcolor="{{element color|p-block}}" | Nh<br />3
| bgcolor="{{element color|p-block}}" | Fl<br />4
| bgcolor="{{element color|p-block}}" | Mc<br />5
| bgcolor="{{element color|p-block}}" | Lv<br />6
| bgcolor="{{element color|p-block}}" | Ts<br />7
| bgcolor="{{element color|p-block}}" | Og<br />8
|} |}
A full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone.<ref name="Greenwood113">Greenwood and Earnshaw, p. 113</ref> For example, magnesium forms Mg<sup>2+</sup> rather than Mg<sup>+</sup> cations when dissolved in water, because the latter would spontaneously ] into Mg<sup>0</sup> and Mg<sup>2+</sup> cations. This is because the ] of hydration (surrounding the cation with water molecules) increases in magnitude with the charge and radius of the ion. In Mg<sup>+</sup>, the outermost orbital (which determines ionic radius) is still 3s, so the hydration enthalpy is small and insufficient to compensate the energy required to remove the electron; but ionizing again to Mg<sup>2+</sup> uncovers the core 2p subshell, making the hydration enthalpy large enough to allow magnesium(II) compounds to form. For similar reasons, the common oxidation states of the heavier p-block elements (where the ns electrons become lower in energy than the np) tend to vary by steps of 2, because that is necessary to uncover an inner subshell and decrease the ionic radius (e.g. Tl<sup>+</sup> uncovers 6s, and Tl<sup>3+</sup> uncovers 5d, so once thallium loses two electrons it tends to lose the third one as well). Analogous arguments based on ] can be used for the less electronegative p-block elements.<ref name=sb45>Siekierski and Burgess, pp. 45–54</ref>{{efn|The normally "forbidden" intermediate oxidation states may be stabilized by forming ], as in <sup>2−</sup> (gallium in the +2 oxidation state) or ] (sulfur in the +5 oxidation state).<ref name=sb45/> Some compounds that appear to be in such intermediate oxidation states are actually mixed-valence compounds, such as ], which contains both Sb(III) and Sb(V).<ref name="Amador">{{cite journal | last1 = Amador | first1 = J. | last2 = Puebla | first2 = E. Gutierrez | last3 = Monge | first3 = M. A. | last4 = Rasines | first4 = I. | last5 = Valero | first5 = C. Ruiz | year = 1988 | title = Diantimony Tetraoxides Revisited | journal = Inorganic Chemistry | volume = 27 | issue = 8 | pages = 1367–1370 | doi = 10.1021/ic00281a011 }}</ref>}}


]
Progressing through a group from lightest element to heaviest element, the outer-shell electrons (those most readily accessible for participation in chemical reactions) are all in the same type of orbital, with a similar shape, but with increasingly higher energy and average distance from the nucleus. For instance, the outer-shell (or&nbsp;"valence") electrons of the first group, headed by hydrogen, all have one electron in an s orbital. In hydrogen, that s orbital is in the lowest possible energy state of any atom, the first-shell orbital (and represented by hydrogen's position in the first period of the table).<ref>{{cite book |last=Hornback |first=Joseph |title=Organic Chemistry |year=2006 |publisher=Thomson Brooks/Cole |location=Pacific Grove |isbn=978-0-534-49317-2 |oclc=66441248|edition=2nd |page=62}}</ref> In francium, the heaviest element of the group, the outer-shell electron is in the seventh-shell orbital, significantly further out on average from the nucleus than those electrons filling all the shells below it in energy. As another example, both carbon and lead have four electrons in their outer shell orbitals.<ref>Gray, pp. 25, 189</ref>
For transition metals, common oxidation states are nearly always at least +2 for similar reasons (uncovering the next subshell); this holds even for the metals with anomalous d<sup>x+1</sup>s<sup>1</sup> or d<sup>x+2</sup>s<sup>0</sup> configurations (except for ]), because repulsion between d-electrons means that the movement of the second electron from the s- to the d-subshell does not appreciably change its ionisation energy.<ref name=sb134>Siekierski and Burgess, pp. 134–137</ref> Because ionizing the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead.<ref name=sb45/> The lanthanides and late actinides generally show a stable +3 oxidation state, removing the outer s-electrons and then (usually) one electron from the (n−2)f-orbitals, that are similar in energy to ns.<ref name=sb178/> The common and maximum oxidation states of the d- and f-block elements tend to depend on the ionisation energies. As the energy difference between the (n−1)d and ns orbitals rises along each transition series, it becomes less energetically favourable to ionize further electrons. Thus, the early transition metal groups tend to prefer higher oxidation states, but the +2 oxidation state becomes more stable for the late transition metal groups. The highest formal oxidation state thus increases from +3 at the beginning of each d-block row, to +7 or +8 in the middle (e.g. ]), and then decrease to +2 at the end.<ref name=sb134/> The lanthanides and late actinides usually have high fourth ionisation energies and hence rarely surpass the +3 oxidation state, whereas early actinides have low fourth ionisation energies and so for example neptunium and plutonium can reach +7.<ref name=johnson/><ref name=sb134/><ref name=sb178>Siekierski and Burgess, pp. 178–180</ref> The very last actinides go further than the lanthanides towards low oxidation states: mendelevium is more easily reduced to the +2 state than thulium or even europium (the lanthanide with the most stable +2 state, on account of its half-filled f-shell), and nobelium outright favours +2 over +3, in contrast to ytterbium.<ref name=rareearths/>


As elements in the same group share the same valence configurations, they usually exhibit similar chemical behaviour. For example, the ]s in the first group all have one valence electron, and form a very homogeneous class of elements: they are all soft and reactive metals. However, there are many factors involved, and groups can often be rather heterogeneous. For instance, hydrogen also has one valence electron and is in the same group as the alkali metals, but its chemical behaviour is quite different. The stable elements of ] comprise a nonmetal (]), two semiconductors (] and ]), and two metals (] and ]); they are nonetheless united by having four valence electrons.<ref name="Scerri14">Scerri, pp. 14–15</ref> This often leads to similarities in maximum and minimum oxidation states (e.g. ] and ] in ] both have maximum oxidation state +6, as in ] and ], and minimum oxidation state −2, as in ]s and ]s); but not always (e.g. ] is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).<ref name=jensenlaw/>
As atomic number (i.e., charge on the atomic nucleus) increases, the ] between the nucleus and the electrons becomes greater, reducing the validity of the quantum mechanical orbital approximation model, which considers each atomic orbital as a separate entity.<ref>{{cite journal |last1=Bethe |first1=Hans |last2=Salpeter |first2=E. E. |year=1957 |title=Quantum Mechanics of One- and Two-Electron Atoms |journal=Encyclopedia of Physics |volume=35 |pages=347–385 |url=http://www.imamu.edu.sa/Scientific_selections/abstracts/Physics/Quantum%20mechanics%20of%20one%20and%20two%20electron%20atoms.pdf |accessdate=11 April 2012 }}</ref>


=== Electronegativity ===
{{clear}}
]
Another important property of elements is their ]. Atoms can form ]s to each other by sharing electrons in pairs, creating an overlap of valence orbitals. The degree to which each atom attracts the shared electron pair depends on the atom's electronegativity<ref name="Greenwood25" /> – the tendency of an atom towards gaining or losing electrons.<ref name="cartoon" /> The more electronegative atom will tend to attract the electron pair more, and the less electronegative (or more electropositive) one will attract it less. In extreme cases, the electron can be thought of as having been passed completely from the more electropositive atom to the more electronegative one, though this is a simplification. The bond then binds two ions, one positive (having given up the electron) and one negative (having accepted it), and is termed an ].<ref name="cartoon" />


Electronegativity depends on how strongly the nucleus can attract an electron pair, and so it exhibits a similar variation to the other properties already discussed: electronegativity tends to fall going up to down, and rise going left to right. The alkali and alkaline earth metals are among the most electropositive elements, while the chalcogens, halogens, and noble gases are among the most electronegative ones.<ref name="Greenwood25" />
=== Atomic and ionic radii ===
{{main|Atomic radius}}
]
Atomic radii vary in a predictable and explicable 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 ] theory of the atom; they provided important evidence for the development and confirmation of ].<ref name="Greenwood27">{{Greenwood&Earnshaw1st|page=27}}</ref>


Electronegativity is generally measured on the Pauling scale, on which the most electronegative reactive atom (]) is given electronegativity 4.0, and the least electronegative atom (]) is given electronegativity 0.79.<ref name="cartoon" /> In fact ] is the most electronegative element, but the Pauling scale cannot measure its electronegativity because it does not form covalent bonds with most elements.<ref>{{cite journal |doi=10.1021/ja00207a003 |title=Electronegativity is the average one-electron energy of the valence-shell electrons in ground-state free atoms|year=1989|author=Allen, Leland C.|journal=Journal of the American Chemical Society |volume=111|pages=9003–9014 |issue=25|bibcode=1989JAChS.111.9003A }}</ref>
The way the atomic radius varies with increasing atomic number (''Z'') can be explained by the arrangement of electrons in shells of fixed capacity. The shells are generally filled in order of increasing radius, since the negatively ] electrons are attracted by the positively charged protons in the nucleus. As the atomic number increases along each row of the periodic table, the additional electrons go into the same outermost shell; whose radius gradually contracts, due to the increasing nuclear charge. In a noble gas, the outermost shell is completely filled; therefore, the additional electron of next alkali metal will go into the next outer shell, accounting for the sudden increase in the atomic radius.<ref name="Greenwood27"/>


An element's electronegativity varies with the identity and number of the atoms it is bonded to, as well as how many electrons it has already lost: an atom becomes more electronegative when it has lost more electrons.<ref name="Greenwood25">Greenwood and Earnshaw, pp. 25–6</ref> This sometimes makes a large difference: lead in the +2 oxidation state has electronegativity 1.87 on the Pauling scale, while lead in the +4 oxidation state has electronegativity 2.33.<ref>{{cite book |last1=Dieter |first1=R. K. |last2=Watson |first2=R. T. |chapter=Transmetalation reactions producing organocopper compounds |pages=443–526 |editor-last1=Rappoport |editor-first1=Z. |editor-last2=Marek |editor-first2=I. |title=The Chemistry of Organocopper Compounds |volume=1 |year=2009 |publisher=John Wiley & Sons |isbn=978-0-470-77296-6 |chapter-url=https://books.google.com/books?id=263AXB0Q6tAC |access-date=6 April 2022 |archive-date=17 October 2022 |archive-url=https://web.archive.org/web/20221017193845/https://books.google.com/books?id=263AXB0Q6tAC |url-status=live }}<!--specifically page 509--></ref>
The electrons in the 4f-subshell, which is progressively filled from ] (''Z''&nbsp;= 58) to ] (''Z''&nbsp;= 71), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the ]s have atomic radii which are smaller than would be expected and which are almost identical to the atomic radii of the elements immediately above them.<ref name="Jolly_contract">
{{cite book
|last1=Jolly |first=W. L.
|year=1991
|title=Modern Inorganic Chemistry
|page=22 |edition=2nd
|publisher=]
|isbn=978-0-07-112651-9
}}</ref> Hence ] has virtually the same atomic radius (and chemistry) as ], and ] has an atomic radius similar to ], and so forth. The effect of the lanthanide contraction is noticeable up to ] (''Z''&nbsp;= 78), after which it is masked by a ] known as the ].<ref name="Greenwood28">{{Greenwood&Earnshaw1st|page=28}}</ref>


=== Metallicity ===
The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it is the poor shielding capacity of the 3d-electrons which affects the atomic radii and chemistries of the elements immediately following the first row of the ]s, from gallium (''Z''&nbsp;= 31) to ] (''Z''&nbsp;= 35).<ref name="Jolly_contract"/>
]
A simple substance is a substance formed from atoms of one chemical element. The simple substances of the more electronegative atoms tend to share electrons (form covalent bonds) with each other. They form either small molecules (like hydrogen or oxygen, whose atoms bond in pairs) or giant structures stretching indefinitely (like carbon or silicon). The noble gases simply stay as single atoms, as they already have a full shell.<ref name="cartoon" /> Substances composed of discrete molecules or single atoms are held together by weaker attractive forces between the molecules, such as the ]: as electrons move within the molecules, they create momentary imbalances of electrical charge, which induce similar imbalances on nearby molecules and create synchronized movements of electrons across many neighbouring molecules.<ref>{{cite web|url=https://www.chemguide.co.uk/atoms/bonding/vdw.html|title=Intermolecular bonding – van der Waals forces|access-date=17 November 2021|archive-date=22 January 2022|archive-url=https://web.archive.org/web/20220122154740/https://www.chemguide.co.uk/atoms/bonding/vdw.html|url-status=live}}</ref>


]
=== Ionization energy and reactivity ===
The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations.<ref name="cartoon" /> The outer orbitals of one atom overlap to share electrons with all its neighbours, creating a giant structure of molecular orbitals extending over all the atoms.<ref name="chemguidemetal">{{cite web |url=https://www.chemguide.co.uk/atoms/bonding/metallic.html |title=Metallic Bonding |last=Clark |first=Jim |date=2019 |website=Chemguide |access-date=30 March 2021 |archive-date=21 April 2021 |archive-url=https://web.archive.org/web/20210421105423/https://www.chemguide.co.uk/atoms/bonding/metallic.html |url-status=live }}</ref> This negatively charged "sea" pulls on all the ions and keeps them together in a ]. Elements forming such bonds are often called ]s; those which do not are often called ]s.<ref name="cartoon" /> Some elements can form multiple simple substances with different structures: these are called ]s. For example, ] and ] are two allotropes of carbon.<ref name="Scerri14" />{{efn|The boundary between dispersion forces and metallic bonding is gradual, like that between ionic and covalent bonding. Characteristic metallic properties do not appear in small mercury clusters, but do appear in large ones.<ref>{{cite journal |last1=Pastor |first1=G. M. |last2=Stampfli |first2=P. |last3=Bennemann |first3=K. |date=1988 |title=On the transition from Van der Waals- to metallic bonding in Hg-clusters as a function of cluster size |url= |journal=Physica Scripta |volume=38 |issue=4 |pages=623–626 |doi=10.1088/0031-8949/38/4/022 |bibcode=1988PhyS...38..623P |s2cid=250842014 }}</ref>}}
{{main|Ionization energy|Reactivity (chemistry)}}
]


The metallicity of an element can be predicted from electronic properties. When atomic orbitals overlap during metallic or covalent bonding, they create both bonding and antibonding ]s of equal capacity, with the antibonding orbitals of higher energy. Net bonding character occurs when there are more electrons in the bonding orbitals than there are in the antibonding orbitals. Metallic bonding is thus possible when the number of electrons delocalized by each atom is less than twice the number of orbitals contributing to the overlap. This is the situation for elements in groups 1 through 13; they also have too few valence electrons to form giant covalent structures where all atoms take equivalent positions, and so almost all of them metallise. The exceptions are hydrogen and boron, which have too high an ionisation energy. Hydrogen thus forms a covalent H<sub>2</sub> molecule, and boron forms a giant covalent structure based on icosahedral B<sub>12</sub> clusters. In a metal, the bonding and antibonding orbitals have overlapping energies, creating a single band that electrons can freely flow through, allowing for electrical conduction.<ref name=Siekierski>Siekierski and Burgess, pp. 60–66</ref>
The first ionization energy is the energy it takes to remove an electron from an atom, and the second ionization energy is the energy it takes to remove a second electron from an atom. Future ionization energies follow this same pattern. Higher ionization energies tend to be larger than lower ionization energies. 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.<ref name="Greenwood28"/>


] prohibits them from having the same energy, so the orbitals hybridize into ''N'' molecular orbitals each with a different energy, where ''N'' is the number of atoms in the crystal. Since ''N'' is such a large number, adjacent orbitals are extremely close together in energy so the orbitals can be considered a continuous energy band. At the actual diamond crystal cell size (denoted by ''a''), two bands are formed, called the valence and conduction bands, separated by a 5.5&nbsp;] band gap. (Here only the valence 2s and 2p electrons have been illustrated; the 1s orbitals do not significantly overlap, so the bands formed from them are much narrower.)]]
Large jumps in the successive molar ionization energies occur when passing noble gas configurations. For example, as can be seen in the table above, the first two molar ionization energies of magnesium (stripping the two 3s electrons from a magnesium atom) are much smaller than the third, which requires stripping off a 2p electron from the very stable ] configuration of Mg<sup>2+</sup>.<ref name="Greenwood28"/>
In group 14, both metallic and covalent bonding become possible. In a diamond crystal, covalent bonds between carbon atoms are strong, because they have a small atomic radius and thus the nucleus has more of a hold on the electrons. Therefore, the bonding orbitals that result are much lower in energy than the antibonding orbitals, and there is no overlap, so electrical conduction becomes impossible: carbon is a nonmetal. However, covalent bonding becomes weaker for larger atoms and the energy gap between the bonding and antibonding orbitals decreases. Therefore, silicon and germanium have smaller ]s and are ]s at ambient conditions: electrons can cross the gap when thermally excited. (Boron is also a semiconductor at ambient conditions.) The band gap disappears in tin, so that tin and lead become metals.<ref name=Siekierski/> As the temperature rises, all nonmetals develop some semiconducting properties, to a greater or lesser extent depending on the size of the band gap. Thus metals and nonmetals may be distinguished by the temperature dependence of their electrical conductivity: a metal's conductivity lowers as temperature rises (because thermal motion makes it more difficult for the electrons to flow freely), whereas a nonmetal's conductivity rises (as more electrons may be excited to cross the gap).<ref name=steudel/>


Elements in groups 15 through 17 have too many electrons to form giant covalent molecules that stretch in all three dimensions. For the lighter elements, the bonds in small diatomic molecules are so strong that a condensed phase is disfavoured: thus nitrogen (N<sub>2</sub>), oxygen (O<sub>2</sub>), white phosphorus and yellow arsenic (P<sub>4</sub> and As<sub>4</sub>), sulfur and red selenium (S<sub>8</sub> and Se<sub>8</sub>), and the stable halogens (F<sub>2</sub>, Cl<sub>2</sub>, Br<sub>2</sub>, and I<sub>2</sub>) readily form covalent molecules with few atoms. The heavier ones tend to form long chains (e.g. red phosphorus, grey selenium, tellurium) or layered structures (e.g. carbon as graphite, black phosphorus, grey arsenic, antimony, bismuth) that only extend in one or two rather than three dimensions. Both kinds of structures can be found as allotropes of phosphorus, arsenic, and selenium, although the long-chained allotropes are more stable in all three. As these structures do not use all their orbitals for bonding, they end up with bonding, nonbonding, and antibonding bands in order of increasing energy. Similarly to group 14, the band gaps shrink for the heavier elements and free movement of electrons between the chains or layers becomes possible. Thus for example black phosphorus, black arsenic, grey selenium, tellurium, and iodine are semiconductors; grey arsenic, antimony, and bismuth are ]s (exhibiting quasi-metallic conduction, with a very small band overlap); and polonium and probably astatine are true metals.<ref name=Siekierski/> Finally, the natural group 18 elements all stay as individual atoms.<ref name=Siekierski/>{{efn|All this describes the situation at standard pressure. Under sufficiently high pressure, the band gaps of any solid drop to zero and metallization occurs. Thus for example at about 170&nbsp;] iodine becomes a metal,<ref name=Siekierski/> and ] should form at pressures of about four million atmospheres.<ref>{{cite journal |last1=McMinis |first1=J. |last2=Clay |first2=R.C. |last3=Lee |first3=D. |last4=Morales |first4=M.A. |year=2015 |title=Molecular to Atomic Phase Transition in Hydrogen under High Pressure |journal=] |volume=114 |issue=10 |page=105305 |doi=10.1103/PhysRevLett.114.105305 |pmid=25815944 |bibcode=2015PhRvL.114j5305M|doi-access=free }}</ref> See ] for values for all nonmetals.}}
Ionization energy is also a periodic trend within the periodic table organization. Moving left to right within a ] or upward within a ], the first ionization energy generally increases. As the ] decreases, it becomes harder to remove an electron that is closer to a more positively charged nucleus.<ref name="Greenwood28"/>


The dividing line between metals and nonmetals is roughly diagonal from top left to bottom right, with the transition series appearing to the left of this diagonal (as they have many available orbitals for overlap). This is expected, as metallicity tends to be correlated with electropositivity and the willingness to lose electrons, which increases right to left and up to down. Thus the metals greatly outnumber the nonmetals. Elements near the borderline are difficult to classify: they tend to have properties that are intermediate between those of metals and nonmetals, and may have some properties characteristic of both. They are often termed semimetals or ]s.<ref name="cartoon" /> The term "semimetal" used in this sense should not be confused with its strict physical sense having to do with band structure: bismuth is physically a semimetal, but is generally considered a metal by chemists.<ref>{{cite journal |last1=Hawkes |first1=Stephen J. |date=2001 |title=Semimetallicity? |journal=Journal of Chemical Education |volume=78 |issue=12 |page=1686 |doi=10.1021/ed078p1686|bibcode=2001JChEd..78.1686H }}</ref>
=== Electronegativity ===
{{main|Electronegativity}}


The following table considers the most stable allotropes at standard conditions. The elements coloured yellow form simple substances that are well-characterised by metallic bonding. Elements coloured light blue form giant network covalent structures, whereas those coloured dark blue form small covalently bonded molecules that are held together by weaker ]s. The noble gases are coloured in violet: their molecules are single atoms and no covalent bonding occurs. Greyed-out cells are for elements which have not been prepared in sufficient quantities for their most stable allotropes to have been characterized in this way. Theoretical considerations and current experimental evidence suggest that all of those elements would metallise if they could form condensed phases,<ref name=Siekierski/> except perhaps for oganesson.<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 |journal=Angewandte Chemie |volume=58 |issue=40 |pages=14260–14264|doi=10.1002/anie.201908327|pmid=31343819|pmc=6790653}}</ref>{{efn|Descriptions of the structures formed by the elements can be found throughout Greenwood and Earnshaw. There are two borderline cases. Arsenic's most stable form conducts electricity like a metal, but the bonding is significantly more localized to the nearest neighbours than it is for the similar structures of antimony and bismuth,<ref>{{cite book |last=Smith |first=J. D. |date=1973 |title=The Chemistry of Arsenic, Antimony and Bismuth |publisher=Pergamon Press |page=556 |isbn=}}</ref> and unlike normal metals it does not have a long liquid range, but rather sublimes instead. Hence its structure is better treated as network covalent.<ref>{{cite book |last1=Rayner-Canham |first1=Geoff |last2=Overton |first2=Tina |author-link= |date=2008 |title=Descriptive Inorganic Chemistry |edition=5th |url= |location=New York |publisher=W. H. Freeman and Company |page=194 |isbn=978-1-4292-2434-5}}</ref> Carbon as ] shows metallic conduction parallel to its planes, but is a semiconductor perpendicular to them. Some computations predict copernicium and flerovium to be nonmetallic,<ref name=CRNL/><ref name=Florez/> but the most recent experiments on them suggest that they are metallic.<ref name=superheavy/><ref name=Ingo/><ref name=Yakushev/> Astatine is calculated to metallise at standard conditions,<ref name="Hermann">{{cite journal
]
|doi=10.1103/PhysRevLett.111.116404|title=Condensed Astatine: Monatomic and Metallic|year=2013|last1=Hermann|first1=A.|last2=Hoffmann|first2=R.|last3=Ashcroft|first3=N. W.|journal=Physical Review Letters|volume=111|issue=11|pages=116404-1–116404-5|bibcode=2013PhRvL.111k6404H|pmid=24074111}}</ref> so presumably tennessine should as well.<ref>{{cite news |last=Ball |first=Philip |date=13 September 2013 |title=
In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, ] is undoubtedly the most electronegative of the elements (not counting noble gases) while ]<!-- not francium; please don't change unless you supply a citation for published experimental results --> is the least electronegative, at least of those elements for which substantial data is available.<ref name="Greenwood30">{{Greenwood&Earnshaw1st|page=30}}</ref>
Metallic properties predicted for astatine |url=https://www.chemistryworld.com/news/metallic-properties-predicted-for-astatine/6582.article |work=Chemistry World |location= |access-date=7 April 2023}}</ref>}}


{{Periodic table (simple substance bonding)}}
There are some exceptions to this general rule. Gallium and ] have higher electronegativities than ] and ] respectively because of the ]. Elements of the ] 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 (see ], ] above).<ref name="Greenwood30"/> The anomalously high electronegativity of ], particularly when compared to ] and ], appears to be an artifact of data selection (and data availability)—methods of calculation other than the Pauling method show the normal periodic trends for these elements.<ref>{{cite journal |last=Allred |first=A. L. |year=1960 |title=Electronegativity values from thermochemical data |journal=Journal of Inorganic and Nuclear Chemistry |volume=17 |issue=3-4 |pages=215-221 |publisher=Northwestern University |doi=10.1016/0022-1902(61)80142-5 |url=http://www.sciencedirect.com/science/article/pii/0022190261801425 |accessdate=11 June 2012 }}</ref>
<gallery mode="packed">
<!--
File:Iron electrolytic and 1cm3 cube.jpg|Iron, a metal
=== Melting and boiling points ===
Sulfur - El Desierto mine, San Pablo de Napa, Daniel Campos Province, Potosí, Bolivia.jpg|Sulfur, a nonmetal
{{main|Melting point|Boiling point}}
Arsen 1a.jpg|Arsenic, an element often called a semi-metal or metalloid
{{empty section|date=May 2012}}
</gallery>


Generally, metals are shiny and dense.<ref name="cartoon" /> They usually have high melting and boiling points due to the strength of the metallic bond, and are often malleable and ductile (easily stretched and shaped) because the atoms can move relative to each other without breaking the metallic bond.<ref name="chemguidem">{{cite web |url=https://www.chemguide.co.uk/atoms/structures/metals.html |title=Metallic Structures |last=Clark |first=Jim |date=2012 |website=Chemguide |access-date=30 March 2021 |archive-date=24 April 2021 |archive-url=https://web.archive.org/web/20210424070514/https://www.chemguide.co.uk/atoms/structures/metals.html |url-status=live }}</ref> They conduct electricity because their electrons are free to move in all three dimensions. Similarly, they conduct heat, which is transferred by the electrons as extra ]: they move faster. These properties persist in the liquid state, as although the crystal structure is destroyed on melting, the atoms still touch and the metallic bond persists, though it is weakened.<ref name="chemguidem" /> Metals tend to be reactive towards nonmetals.<ref name="cartoon" /> Some exceptions can be found to these generalizations: for example, beryllium, chromium,<ref name=raynercanham/> manganese,<ref name="Holl">{{cite book|publisher=Walter de Gruyter|date=1985|edition=91–100 |pages=1110–1117|isbn=978-3-11-007511-3|title=Lehrbuch der Anorganischen Chemie|first=Arnold F.|last=Holleman|author2=Wiberg, Egon|author3=Wiberg, Nils|language=de|chapter=Mangan}}</ref> antimony,<ref name="wiberg_holleman">{{cite book|title=Inorganic chemistry|author=Wiberg, Egon|author2=Wiberg, Nils|author3=Holleman, Arnold Frederick|name-list-style=amp|publisher=Academic Press|date=2001|isbn=978-0-12-352651-9|page=758}}</ref> bismuth,<ref name="CRC">{{cite book| first = C. R.| last = Hammond| pages = | title = The Elements, in Handbook of Chemistry and Physics| edition = 81st| location = Boca Raton (FL, US)| publisher = CRC press| isbn = 978-0-8493-0485-9| date = 2004| url-access = registration| url = https://archive.org/details/crchandbookofche81lide/page/4}}</ref> and uranium are brittle (not an exhaustive list);<ref name=raynercanham/> chromium is extremely hard;<ref name=r1>{{cite book|editor=G.V. Samsonov|chapter=Mechanical Properties of the Elements|doi=10.1007/978-1-4684-6066-7_7|isbn=978-1-4684-6066-7|url=https://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|publisher=IFI-Plenum|place=New York, USA|year=1968|pages=387–446|archive-url=https://web.archive.org/web/20150402123344/https://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|archive-date=2 April 2015 |title=Handbook of the Physicochemical Properties of the Elements }}</ref> gallium, rubidium, caesium, and mercury are liquid at or close to room temperature;{{efn|See ]. The same is probably true of francium, but due to its extreme instability, this has never been experimentally confirmed. Copernicium and flerovium are expected to be liquids,<ref name="CRNL">{{cite journal |last1=Mewes |first1=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 |date=2019 |volume=58|issue=50|pages=17964–17968|doi=10.1002/anie.201906966 |pmid=31596013 |pmc=6916354 |url=}}</ref><ref name=Florez>{{cite journal |last1=Florez |first1=Edison |last2=Smits |first2=Odile R. |last3=Mewes |first3=Jan-Michael |last4=Jerabek |first4=Paul |last5=Schwerdtfeger |first5=Peter |date=2022 |title=From the gas phase to the solid state: The chemical bonding in the superheavy element flerovium |journal=The Journal of Chemical Physics |volume=157 |issue=6 |page=064304 |doi=10.1063/5.0097642|pmid=35963734 |bibcode=2022JChPh.157f4304F |s2cid=250539378 }}</ref> similar to mercury, and experimental evidence suggests that they are metals.<ref name="superheavy">
=== Density ===
{{Cite web
{{main|Density}}
|last1=Gäggeler
{{empty section|date=May 2012}}-->
|first1=H. W.
|year=2007
|title=Gas Phase Chemistry of Superheavy Elements
|url=https://lch.web.psi.ch/files/lectures/TexasA&M/TexasA&M.pdf
|pages=26–28
|publisher=]
|archive-url=https://web.archive.org/web/20120220090755/https://lch.web.psi.ch/files/lectures/TexasA%26M/TexasA%26M.pdf
|archive-date=20 February 2012
}}</ref><ref name=Ingo>{{cite news |last=Ingo |first=Peter |date=15 September 2022 |title=Study shows flerovium is the most volatile metal in the periodic table |url=https://phys.org/news/2022-09-flerovium-volatile-metal-periodic-table.html |work=phys.org<!--but provided by GSI Helmholtz--> |location= |access-date=22 November 2022}}</ref><ref name=Yakushev>{{cite journal |last1=Yakushev |first1=A. |last2=Lens |first2=L. |first3=Ch. E. |last3=Düllmann |first4=J. |last4=Khuyagbaatar |first5=E. |last5=Jäger |first6=J. |last6=Krier |first7=J. |last7=Runke |first8=H. M. |last8=Albers |first9=M. |last9=Asai |first10=M. |last10=Block |first11=J. |last11=Despotopulos |first12=A. |last12=Di Nitto |first13=K. |last13=Eberhardt |first14=U. |last14=Forsberg |first15=P. |last15=Golubev |first16=M. |last16=Götz |first17=S. |last17=Götz |first18=H. |last18=Haba |first19=L. |last19=Harkness-Brennan |first20=R.-D. |last20=Herzberg |first21=F. P. |last21=Heßberger |first22=D. |last22=Hinde |first23=A. |last23=Hübner |first24=D. |last24=Judson |first25=B. |last25=Kindler |first26=Y. |last26=Komori |first27=J. |last27=Konki |first28=J. V. |last28=Kratz |first29=N. |last29=Kurz |first30=M. |last30=Laatiaoui |first31=S. |last31=Lahiri |first32=B. |last32=Lommel |first33=M. |last33=Maiti |first34=A. K. |last34=Mistry |first35=Ch. |last35=Mokry |first36=K. J. |last36=Moody |first37=Y. |last37=Nagame |first38=J. P. |last38=Omtvedt |first39=P. |last39=Papadakis |first40=V. |last40=Pershina |first41=D. |last41=Rudolph |first42=L. G. |last42=Samiento |first43=T. K. |last43=Sato |first44=M. |last44=Schädel |first45=P. |last45=Scharrer |first46=B. |last46=Schausten |first47=D. A. |last47=Shaughnessy |first48=J. |last48=Steiner |first49=P. |last49=Thörle-Pospiech |first50=A. |last50=Toyoshima |first51=N. |last51=Trautmann |first52=K. |last52=Tsukada |first53=J. |last53=Uusitalo |first54=K.-O. |last54=Voss |first55=A. |last55=Ward |first56=M. |last56=Wegrzecki |first57=N. |last57=Wiehl |first58=E. |last58=Williams |first59=V. |last59=Yakusheva |display-authors=3 |date=25 August 2022 |title=On the adsorption and reactivity of element 114, flerovium |journal=Frontiers in Chemistry |volume=10 |issue=976635 |page=976635 |doi=10.3389/fchem.2022.976635 |pmid=36092655 |pmc=9453156 |bibcode=2022FrCh...10.6635Y |doi-access=free }}</ref>}} and ]s such as gold are chemically very inert.<ref>{{cite journal |doi=10.1038/376238a0 |title=Why gold is the noblest of all the metals |date=1995 |last1=Hammer |first1=B. |last2=Norskov |first2=J. K. |journal=Nature |volume=376 |issue=6537 |pages=238–240 |bibcode=1995Natur.376..238H|s2cid=4334587 }}</ref><ref>{{cite journal |doi=10.1103/PhysRevB.6.4370 |title=Optical Constants of the Noble Metals |date=1972 |last1=Johnson |first1=P. B. |last2=Christy |first2=R. W. |journal=Physical Review B |volume=6 |issue=12 |pages=4370–4379 |bibcode=1972PhRvB...6.4370J}}</ref>


Nonmetals exhibit different properties. Those forming giant covalent crystals exhibit high melting and boiling points, as it takes considerable energy to overcome the strong covalent bonds. Those forming discrete molecules are held together mostly by dispersion forces, which are more easily overcome; thus they tend to have lower melting and boiling points,<ref>{{cite web |url=https://www.chemguide.co.uk/inorganic/period3/elementsphys.html |title=Atomic and Physical Properties of the Period 3 Elements |last=Clark |first=Jim |date=2018 |website=Chemguide |access-date=30 March 2021 |archive-date=22 April 2021 |archive-url=https://web.archive.org/web/20210422142013/https://www.chemguide.co.uk/inorganic/period3/elementsphys.html |url-status=live }}</ref> and many are liquids or gases at room temperature.<ref name="cartoon" /> Nonmetals are often dull-looking. They tend to be reactive towards metals, except for the noble gases, which are inert towards most substances.<ref name="cartoon" /> They are brittle when solid as their atoms are held tightly in place. They are less dense and conduct electricity poorly,<ref name="cartoon" /> because there are no mobile electrons.<ref name="group4">{{cite web |url=https://www.chemguide.co.uk/inorganic/group4/properties.html |title=The Trend From Non-Metal to Metal In the Group 4 Elements |last=Clark |first=Jim |date=2015 |website=Chemguide |access-date=30 March 2021 |archive-date=27 April 2021 |archive-url=https://web.archive.org/web/20210427234147/https://www.chemguide.co.uk/inorganic/group4/properties.html |url-status=live }}</ref> Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium,<ref name="group4" /> and tellurium.<ref name=Siekierski/> Selenium has both a semiconducting grey allotrope and an insulating red allotrope; arsenic has a metallic grey allotrope, a semiconducting black allotrope, and an insulating yellow allotrope (though the last is unstable at ambient conditions).<ref name=steudel/> Again there are exceptions; for example, diamond has the highest thermal conductivity of all known materials, greater than any metal.<ref name=PNU>{{cite journal |doi=10.1103/PhysRevLett.70.3764 |title=Thermal conductivity of isotopically modified single crystal diamond |year=1993 |last1=Wei |first1=Lanhua |last2=Kuo |first2=P. K. |last3=Thomas |first3=R. L. |last4=Anthony |first4=T. R. |last5=Banholzer |first5=W. F. |journal=Physical Review Letters |volume=70 |issue=24 |pages=3764–3767 |pmid=10053956 |bibcode=1993PhRvL..70.3764W}}</ref>
== Alternatives ==
{{main|Alternative periodic tables}}


It is common to designate a class of metalloids straddling the boundary between metals and nonmetals, as elements in that region are intermediate in both physical and chemical properties.<ref name="cartoon" /> However, no consensus exists in the literature for precisely which elements should be so designated. When such a category is used, silicon, germanium, arsenic, and tellurium are almost always included, and boron and antimony usually are; but most sources include other elements as well, without agreement on which extra elements should be added, and some others subtract from this list instead.{{efn|See ]. For example, a periodic table used by the American Chemical Society includes polonium as a metalloid,<ref name="ACS" /> but one used by the Royal Society of Chemistry does not,<ref>{{cite web |url=https://www.rsc.org/periodic-table |title=Periodic Table |date=2021 |website=www.rsc.org |publisher=] |access-date=27 March 2021 |archive-date=21 March 2021 |archive-url=https://web.archive.org/web/20210321033913/https://www.rsc.org/periodic-table |url-status=live }}</ref> and that included in the '']'' does not refer to metalloids or semi-metals at all.<ref name="EB" /> Classification can change even within a single work. For example, Sherwin and Weston's ''Chemistry of the Non-Metallic Elements'' (1966) has a periodic table on p. 7 classifying antimony as a nonmetal, but on p. 115 it is called a metal.<ref>{{cite book |last1=Sherwin |first1=E. |last2=Weston |first2=G. J. |editor=Spice, J. E. |date=1966 |title=Chemistry of the Non-Metallic Elements |publisher=Pergamon Press |isbn=978-1-4831-3905-0}}</ref>|name=metalloids}} For example, unlike all the other elements generally considered metalloids or nonmetals, antimony's only stable form has metallic conductivity. Moreover, the element resembles bismuth and, more generally, the other p-block metals in its physical and chemical behaviour. On this basis some authors have argued that it is better classified as a metal than as a metalloid.<ref name=raynercanham/><ref name=hawkes>{{cite journal |last1=Hawkes |first1=Stephen J. |date=2001 |title=Semimetallicity? |url= |journal=Journal of Chemical Education |volume=78 |issue=12 |pages=1686–1687 |doi=10.1021/ed078p1686 |bibcode=2001JChEd..78.1686H |access-date=}}</ref><ref name=steudel>{{cite book |last1=Steudel |first1=Ralf |first2=David |last2=Scheschkewitz |author-link= |date=2020 |title=Chemistry of the Non-Metals |url= |location= |publisher=Walter de Gruyter |pages=154–155, 425, 436 |isbn=978-3-11-057805-8 |quote=In Group 15 of the Periodic Table, as in both neighboring groups, the metallic character increases when going down. More specifically, there is a transition from a purely non-metallic element (N) via elements with nonmetallic and metallic modifications to purely metallic elements (Sb, Bi). This chapter addresses the two elements besides nitrogen, which are clearly nonmetallic under standard conditions: phosphorus and arsenic. The chemistry of arsenic, however, is only briefly described as many of the arsenic compounds resemble the corresponding phosphorus species.}}</ref> On the other hand, selenium has some semiconducting properties in its most stable form (though it also has insulating allotropes) and it has been argued that it should be considered a metalloid<ref name=hawkes/> – though this situation also holds for phosphorus,<ref name=steudel/> which is a much rarer inclusion among the metalloids.{{efn|name=metalloids}}
]


=== Further manifestations of periodicity ===
While the iconic format presented above is widely used,<ref name="Gray12"/> other ] exist, including not only various rectangular formats, but also circular or cylindrical versions in which the rows (periods) flow from one into another, without the arbitrary breaks required at the margins of the usual printed or screen-formatted versions. Alternative periodic tables are developed often to highlight or emphasize different chemical or physical properties of the elements which are not as apparent in traditional periodic tables. Some tables aim to emphasize both the nucleon and electronic structure of atoms. This can be done by changing the spatial relationship or representation each element has with respect to another element in the table. <!--Other tables aim to emphasize the chemical element isolations by humans over time.{{Citation needed|date=February 2012}}-->
There are some other relationships throughout the periodic table between elements that are not in the same group, such as the ]s between elements that are diagonally adjacent (e.g. lithium and magnesium).<ref name="PTSS2">Scerri, pp. 407–420</ref> Some similarities can also be found between the main groups and the transition metal groups, or between the early actinides and early transition metals, when the elements have the same number of valence electrons. Thus uranium somewhat resembles chromium and tungsten in group 6,<ref name="PTSS2" /> as all three have six valence electrons.<ref name="Jensen" /> Relationships between elements with the same number of valence electrons but different types of valence orbital have been called secondary or isodonor relationships: they usually have the same maximum oxidation states, but not the same minimum oxidation states. For example, chlorine and manganese both have +7 as their maximum oxidation state (e.g. ] and ]), but their respective minimum oxidation states are −1 (e.g. ]) and −3 (K<sub>2</sub>). Elements with the same number of valence vacancies but different numbers of valence electrons are related by a tertiary or isoacceptor relationship: they usually have similar minimum but not maximum oxidation states. For example, hydrogen and chlorine both have −1 as their minimum oxidation state (in ]s and ]s), but hydrogen's maximum oxidation state is +1 (e.g. ]) while chlorine's is +7.<ref name=jensenlaw/>


Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as ]s, ]s, ], ], ], and so on. Similar periodic variations appear for the compounds of the elements, which can be observed by comparing hydrides, oxides, sulfides, halides, and so on.<ref name="Greenwood25" /> Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include the variation in the ]ic and ] properties of the elements and their compounds, the stabilities of compounds, and methods of isolating the elements.<ref name="Greenwood27" /> Periodicity is and has been used very widely to predict the properties of unknown new elements and new compounds, and is central to modern chemistry.<ref name="Greenwood29bis">Greenwood and Earnshaw, pp. 29–31</ref>
A common alternate layout is ]'s Left Step Periodic Table, which organizes elements according to orbital filling. The modern version, known as the ADOMAH Periodic Table named after the biblical ], Adomah,<ref>{{cite web|title=Periodic Law can be understood in terms of the Tetrahedral Sphere Packing! |author=Tsimmerman, Valery|date=19 March 2009 |accessdate=5 July 2012 |url=http://perfectperiodictable.com/|publisher=]}}</ref>, helps with writing electron configurations; the table is oriented 90˚ from the traditional periodic table, with the s-block moved to the end, after the noble gases.<ref>{{cite journal |title=Charles Janet: unrecognized genius of the periodic system |last=Stewart |first=Philip J. |year=2009 |journal=Foundations of Chemistry |volume=12 |issue=1 |accessdate=April 2010|doi=10.1007/s10698-008-9062-5}}</ref>


== Classification of elements ==
Another alternative layout is Theodor Benfey's periodic table, where elements are arranged in a spiral with hydrogen at the center and spiraling outward, with the transition metals, lanthanides, and actinides as peninsulas.<ref>{{cite journal |last=Seaborg |first=Glenn |year=1964 |title=Plutonium: The Ornery Element |journal=Chemistry |volume=37 |issue=6 |pages=14 }}</ref>
] |access-date=27 March 2021 |archive-date=3 February 2021 |archive-url=https://web.archive.org/web/20210203123434/https://www.acs.org/content/acs/en/education/whatischemistry/periodictable.html |url-status=live }}</ref> Lutetium and lawrencium in group 3 are also transition metals.<ref name=jensenlaw/>{{col-begin}}{{col-break}}{{Legend inline|ff9d9d}} ]<br />{{Legend inline|ffdead}} ]<br />{{Legend inline|ffbfff}} ]<br />{{Legend inline|ff99cc}} ]<br />{{Legend inline|ffc0c0}} ]{{col-break}}{{Legend inline|cccccc}} ]<br />{{Legend inline|cccc99}} ]<br />{{Legend inline|a0ffa0}} ]s<br />{{Legend inline|ffff99}} ]<br />{{Legend inline|c0ffff}} ]{{col-end}}]]
Many terms have been used in the literature to describe sets of elements that behave similarly. The group names ''alkali metal'', ''alkaline earth metal'', ''triel'', ''tetrel'', ''pnictogen'', ''chalcogen'', ''halogen'', and ''noble gas'' are acknowledged by IUPAC; the other groups can be referred to by their number, or by their first element (e.g., group 6 is the chromium group).<ref name="IUPAC" /><ref>{{cite journal |last1=Fernelius |first1=W. C. |last2=Loening |first2=Kurt |last3=Adams |first3=Roy M. |date=1971 |title=Names of groups and elements |url= |journal=Journal of Chemical Education |volume=48 |issue=11 |pages=730–731 |doi=10.1021/ed048p730 |bibcode=1971JChEd..48..730F }}</ref> Some divide the p-block elements from groups 13 to 16 by metallicity,<ref name="EB" /><ref name="ACS" /> although there is neither an IUPAC definition nor a precise consensus on exactly which elements should be considered metals, nonmetals, or semi-metals (sometimes called metalloids).<ref name="EB" /><ref name="ACS" /><ref name="IUPAC" /> Neither is there a consensus on what the metals succeeding the transition metals ought to be called, with '']'' and ''poor metal'' being among the possibilities having been used. Some advanced monographs exclude the elements of group 12 from the transition metals on the grounds of their sometimes quite different chemical properties, but this is not a universal practice<ref>{{cite journal|last1=Jensen |first1=William B. |year=2003 |title=The Place of Zinc, Cadmium, and Mercury in the Periodic Table |journal=Journal of Chemical Education |volume=80 |issue=8 |pages=952–961 |doi=10.1021/ed080p952 |bibcode=2003JChEd..80..952J |url=https://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |access-date=6 May 2012 |archive-url=https://web.archive.org/web/20100611152417/https://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/091.%20Zn-Cd-Hg.pdf |archive-date=11 June 2010 }}</ref> and IUPAC does not presently mention it as allowable in its ''Principles of Chemical Nomenclature''.<ref>{{cite book |editor-last=Leigh |editor-first=G. J. |date=2011 |title=Principles of Chemical Nomenclature |url=https://iupac.org/wp-content/uploads/2021/12/Principles_Leigh2011-compressed.pdf |location= |publisher=The Royal Society of Chemistry |page=9 |isbn=978-1-84973-007-5}}</ref>


The ''lanthanides'' are considered to be the elements La–Lu, which are all very similar to each other: historically they included only Ce–Lu, but lanthanum became included by common usage.<ref name=IUPAC/> The '']s'' (or rare earth metals) add scandium and yttrium to the lanthanides.<ref name="IUPAC" /> Analogously, the ''actinides'' are considered to be the elements Ac–Lr (historically Th–Lr),<ref name="IUPAC" /> although variation of properties in this set is much greater than within the lanthanides.<ref name="Jorgensen" /> IUPAC recommends the names ''lanthanoids'' and ''actinoids'' to avoid ambiguity, as the -ide suffix typically denotes a negative ion; however ''lanthanides'' and ''actinides'' remain common.<ref name="IUPAC" /> With the increasing recognition of lutetium and lawrencium as d-block elements, some authors began to define the lanthanides as La–Yb and the actinides as Ac–No, matching the f-block.<ref name=KW/><ref name=Jensen1982/><ref>{{cite web |url=https://www.webelements.com/ |title=WebElements |last=Winter |first=Mark |date=1993–2022 |website= |publisher=The University of Sheffield and WebElements Ltd, UK |access-date=5 December 2022 |quote=}}</ref><ref>{{cite book |last=Cowan |first=Robert D. |date=1981 |title=The Theory of Atomic Structure and Spectra |url= |location= |publisher=University of California Press |page=598 |isbn=978-0-520-90615-0}}</ref><ref>{{cite journal |last1=Villar |first1=G. E. |date=1966 |title=A suggested modification to the periodic chart |url= |journal=Journal of Inorganic and Nuclear Chemistry |volume=28 |issue=1 |pages=25–29 |doi=10.1016/0022-1902(66)80224-5 }}</ref><ref name=sacotton>{{cite journal |last1=Cotton |first1=S. A. |date=1996 |title=After the actinides, then what? |url= |journal=Chemical Society Reviews |volume=25 |issue=3 |pages=219–227 |doi=10.1039/CS9962500219 |access-date=}}</ref> The ''transactinides'' or '']s'' are the short-lived elements beyond the actinides, starting at lawrencium or rutherfordium (depending on where the actinides are taken to end).<ref name=sacotton/><ref name=Neve>{{cite journal |last1=Neve |first1=Francesco |date=2022 |title=Chemistry of superheavy transition metals |url= |journal=Journal of Coordination Chemistry |volume=75 |issue=17–18 |pages=2287–2307 |doi=10.1080/00958972.2022.2084394 |s2cid=254097024 }}</ref><ref name=Mingos>{{cite book |last=Mingos |first=Michael |author-link=Michael Mingos |date=1998 |title=Essential Trends in Inorganic Chemistry |url= |location= |publisher=Oxford University Press |page=387 |isbn=978-0-19-850109-1}}<!--uses "transactinide" rather than "superheavy--></ref><ref>{{cite web |url=https://science.osti.gov/-/media/np/nsac/pdf/202310/October-4-LRP-Report.pdf |archive-url=https://web.archive.org/web/20231005134013/https://science.osti.gov/-/media/np/nsac/pdf/202310/October-4-LRP-Report.pdf |archive-date=5 October 2023 |title=A New Era of Discovery: the 2023 Long Range Plan for Nuclear Science |publisher=U.S. Department of Energy |date=October 2023 |via=OSTI |access-date=20 October 2023 |quote=Superheavy elements (''Z'' > 102) are teetering at the limits of mass and charge.}}</ref><ref>{{cite arXiv|last=Kragh|first=Helge|author-link=|date=2017|title=The search for superheavy elements: Historical and philosophical perspectives|eprint=1708.04064|class=physics.hist-ph}}</ref>
Three-dimensional periodic tables exist as well, such as ]'s periodic table, which has four billboards, each representing a block, with elements on the front and back. Hydrogen and helium are omitted.<ref>{{cite book |title=Graphical Representations of the Periodic System During One Hundred Years |last=Mazurs |first= E.G. |year= 1974 |publisher=University of Alabama Press |location=Alabama |accessdate=8 April 2012 |isbn=9780817332006}}</ref>


Many more categorizations exist and are used according to certain disciplines. In astrophysics, a metal is defined as any element with atomic number greater than 2, i.e. anything except hydrogen and helium.<ref>{{cite web |url=https://icc.dur.ac.uk/~tt/Lectures/Galaxies/TeX/lec/node27.html |title=Metallicity of stars |last=Theuns |first=Tom |website=icc.dur.ac.uk |publisher=Durham University |access-date=27 March 2021 |archive-date=27 September 2021 |archive-url=https://web.archive.org/web/20210927160927/https://icc.dur.ac.uk/~tt/Lectures/Galaxies/TeX/lec/node27.html |url-status=live }}</ref> The term "semimetal" has a different definition in physics than it does in chemistry: bismuth is a semimetal by physical definitions, but chemists generally consider it a metal.<ref>{{cite book
== Future developments and the end of the periodic table ==
|first=Gerald
{{main|Extended periodic table}}
|last=Burns
|title=Solid State Physics
|date=1985
|publisher=Academic Press, Inc.
|isbn=978-0-12-146070-9
|pages=339–40}}</ref> A few terms are widely used, but without any very formal definition, such as "]", which has been given such a wide range of definitions that it has been criticized as "effectively meaningless".<ref>{{cite journal |last1=Duffus |first1=John H. |date=2002 |title="Heavy Metals"–A Meaningless Term? |url=https://publications.iupac.org/pac/2002/pdf/7405x0793.pdf |journal=Pure and Applied Chemistry |volume=74 |issue=5 |pages=793–807 |doi=10.1351/pac200274050793 |s2cid=46602106 |access-date=27 March 2021 |archive-date=11 April 2021 |archive-url=https://web.archive.org/web/20210411012337/https://publications.iupac.org/pac/2002/pdf/7405x0793.pdf |url-status=live }}</ref>


The scope of terms varies significantly between authors. For example, according to IUPAC, the noble gases extend to include the whole group, including the very radioactive superheavy element oganesson.<ref>{{cite journal|last=Koppenol|first=W.|date=2016|title=How to name new chemical elements|journal=Pure and Applied Chemistry|publisher=DeGruyter|doi=10.1515/pac-2015-0802 |hdl=10045/55935|s2cid=102245448|url=https://rua.ua.es/dspace/bitstream/10045/55935/1/2016_Koppenol_etal_PureApplChem.pdf|hdl-access=free|access-date=15 August 2021|archive-date=11 May 2020|archive-url=https://web.archive.org/web/20200511193435/https://rua.ua.es/dspace/bitstream/10045/55935/1/2016_Koppenol_etal_PureApplChem.pdf|url-status=live}}</ref> However, among those who specialize in the superheavy elements, this is not often done: in this case "noble gas" is typically taken to imply the unreactive behaviour of the lighter elements of the group. Since calculations generally predict that oganesson should not be particularly inert due to relativistic effects, and may not even be a gas at room temperature if it could be produced in bulk, its status as a noble gas is often questioned in this context.<ref>{{cite journal |last1=Roth |first1=Klaus |date=3 April 2018 |title=Is Element 118 a Noble Gas? |url=https://www.chemistryviews.org/details/ezine/10907570/New_Kids_on_the_Table_Is_Element_118_a_Noble_Gas__Part_1.html |journal=Chemie in unserer Zeit |doi=10.1002/chemv.201800029 |access-date=27 March 2021 |archive-date=2 March 2021 |archive-url=https://web.archive.org/web/20210302084128/https://www.chemistryviews.org/details/ezine/10907570/New_Kids_on_the_Table_Is_Element_118_a_Noble_Gas__Part_1.html |url-status=live }}</ref> Furthermore, national variations are sometimes encountered: in Japan, alkaline earth metals often do not include beryllium and magnesium as their behaviour is different from the heavier group 2 metals.<ref>{{cite web |url=https://www.chemistry.or.jp/news/information/1-2.html |title=【お知らせ】高等学校化学で用いる用語に関する提案（1）への反応 |author=The Chemical Society of Japan |date=25 January 2018 |website=www.chemistry.or.jp |publisher=The Chemical Society of Japan |access-date=3 April 2021 |quote=「12．アルカリ土類金属」の範囲についても，△を含めれば，すべての教科書で提案が考慮されている。歴史的には第4 周期のカルシウム以下を指していた用語だったが，「周期表の2 族に対応する用語とする」というIUPAC の勧告1）に従うのは現在では自然な流れだろう。 |archive-date=16 May 2021 |archive-url=https://web.archive.org/web/20210516062728/https://www.chemistry.or.jp/news/information/1-2.html |url-status=live }}</ref>
Although all elements up to ununoctium have been discovered, only the first ] and ] have known chemical properties. The other elements may behave differently from what would be predicted by extrapolation, due to relativistic effects; for example, ] has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the ].<ref>{{cite book|title=The Chemistry of Superheavy Elements|last=Schändel|first=Matthias|year=2003|publisher=Kluwer Academic Publishers|location=Dordrecht|isbn=1-4020-1250-0|page=277|accessdate=29 January 2012}}</ref>


== History ==
It is unclear whether new elements will continue the pattern of the current periodic table as ], or require further adaptations or adjustments. ] expected the eighth period, which includes a two-element s-block for elements ] and ], a ] for the next 18 elements, and 30 additional elements continuing the current f-, d-, and p-blocks.<ref>{{cite journal|doi=10.2307/3963006|last=Frazier|first=K.|title=Superheavy Elements|journal=Science News|volume=113|issue=15|pages=236–238|year=1978|jstor=3963006}}</ref> On the other side, some physicists like Pekka Pyykkö have theorized that these additional elements do not follow the ], which predicts how electron shells are filled, and thus affect the appearance of the present periodic table.<ref>{{Cite journal|last1=Pyykkö|first1=Pekka|title=A suggested periodic table up to Z ≤ 172, based on Dirac–Fock calculations on atoms and ions|journal=Physical Chemistry Chemical Physics|volume=13|issue=1|pages=161–8|year=2011|pmid=20967377|doi=10.1039/c0cp01575j|bibcode = 2011PCCP...13..161P}}</ref>
{{Main|History of the periodic table}}
{{See also|Timeline of chemical element discoveries}}
=== Early history ===
In 1817, German physicist ] began to formulate one of the earliest attempts to classify the elements.<ref>{{cite journal|last1=Wurzer|first1=Ferdinand|title=Auszug eines Briefes vom Hofrath Wurzer, Prof. der Chemie zu Marburg|journal=Annalen der Physik|date=1817|volume=56|issue=7|pages=331–334|url=https://babel.hathitrust.org/cgi/pt?id=chi.096071138;view=1up;seq=351|trans-title=Excerpt of a letter from Court Advisor Wurzer, Professor of Chemistry at Marburg|language=de|doi=10.1002/andp.18170560709|bibcode=1817AnP....56..331.|access-date=15 August 2021|archive-date=8 October 2021|archive-url=https://web.archive.org/web/20211008024621/https://babel.hathitrust.org/cgi/pt?id=chi.096071138;view=1up;seq=351|url-status=live}} Here, Döbereiner found that strontium's properties were intermediate to those of calcium and barium.</ref> In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups ].<ref>{{cite journal|last1=Döbereiner|first1=J. W.|title=Versuch zu einer Gruppirung der elementaren Stoffe nach ihrer Analogie|journal=Annalen der Physik und Chemie|date=1829|volume=15|issue=2|pages=301–307|url=https://babel.hathitrust.org/cgi/pt?id=mdp.39015065410634;view=1up;seq=315|series=2nd series|trans-title=An attempt to group elementary substances according to their analogies|language=de|bibcode=1829AnP....91..301D|doi=10.1002/andp.18290910217|access-date=15 August 2021|archive-date=8 October 2021|archive-url=https://web.archive.org/web/20211008024625/https://babel.hathitrust.org/cgi/pt?id=mdp.39015065410634;view=1up;seq=315|url-status=live}} For an English translation of this article, see: {{Webarchive|url=https://web.archive.org/web/20190309161429/https://web.lemoyne.edu/~GIUNTA/dobereiner.html |date=9 March 2019 }}</ref><ref>{{cite book |last=Horvitz |first=L.|title=Eureka!: Scientific Breakthroughs That Changed The World |year=2002 |publisher=John Wiley |location=New York|isbn=978-0-471-23341-1 |oclc=50766822 |page=43|bibcode=2001esbt.book.....H}}</ref> Chlorine, bromine, and iodine formed a triad; as did calcium, strontium, and barium; lithium, sodium, and potassium; and sulfur, selenium, and tellurium. Today, all these triads form part of modern-day groups: the halogens, alkaline earth metals, alkali metals, and chalcogens.<ref>Scerri, p. 47</ref> Various chemists continued his work and were able to identify more and more relationships between small groups of elements. However, they could not build one scheme that encompassed them all.<ref>{{cite book|last=Ball|first=P.|author-link=Philip Ball|title=The Ingredients: A Guided Tour of the Elements |location=Oxford|publisher=Oxford University Press |year=2002 |isbn=978-0-19-284100-1|page=100}}</ref>


]
] noted<ref>{{cite web|author=Elert, G.|title=Atomic Models|url=http://physics.info/atomic-models/|work=The Physics Hypertextbook|accessdate=2009-10-09}}</ref> that literally interpreting the ] ] has problems with electron orbitals at ''Z'' > 137, suggesting that neutral atoms cannot exist beyond ], and that a periodic table based on electron orbitals breaks down at this point. A more rigorous analysis calculates the limit to be ''Z'' ≈ 173.<ref name=AJP>{{cite journal|last1=Greiner|first1=W.|last2=Schramm|first2=S.|year=2008|title=]|volume=76|page=509}}</ref>
] published a letter in the ''Chemical News'' in February 1863 on the periodicity among the chemical elements.<ref name=EB1911>{{cite EB1911 |wstitle=Newlands, John Alexander Reina |volume=19 |page=515}}</ref> In 1864 Newlands published an article in the ''Chemical News'' showing that if the elements are arranged in the order of their atomic weights, those having consecutive numbers frequently either belong to the same group or occupy similar positions in different groups, and he pointed out that each eighth element starting from a given one is in this arrangement a kind of repetition of the first, like the eighth note of an octave in music (The Law of Octaves).<ref name=EB1911/> However, Newlands's formulation only worked well for the main-group elements, and encountered serious problems with the others.<ref name=jensenlaw/>


German chemist ] noted the sequences of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates (i.e. vertically) and the atomic volumes as abscissas (i.e. horizontally)—the curve obtained a series of maximums and minimums—the most ] elements would appear at the peaks of the curve in the order of their atomic weights. In 1864, a book of his was published; it contained an early version of the periodic table containing 28 elements, and classified elements into six families by their ]—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight had until then been stymied by inaccurate measurements of the atomic weights.<ref name="Meyer table">Meyer, Julius Lothar; Die modernen Theorien der Chemie (1864); {{Webarchive|url=https://web.archive.org/web/20190102050414/https://reader.digitale-sammlungen.de/de/fs1/object/goToPage/bsb10073411.html?pageNo=147 |date=2 January 2019 }}</ref> In 1868, he revised his table, but this revision was published as a draft only after his death.<ref>Scerri, pp. 106–108</ref>
=== Calculations from the Bohr model{{Anchor|Bohr model}} ===
The ] has problems for atoms with atomic number greater than 137, because the speed of an electron in a ], ''v'', is given by


=== Mendeleev ===
:<math>v = Z \alpha c \approx \frac{Z c}{137.036}</math>
{{multiple image
| total_width = 700
| align = right
| image1 = 1869-periodic-table.jpg
| caption1 = ] 1869 periodic table
| image2 = Mendelejevs periodiska system 1871.png
| caption2 = ] 1871 periodic table
}}
The definitive breakthrough came from the Russian chemist ]. Although other chemists (including Meyer) had found some other versions of the periodic system at about the same time, Mendeleev was the most dedicated to developing and defending his system, and it was his system that most affected the scientific community.<ref>Scerri, p. 113</ref> On 17 February 1869 (1 March 1869 in the Gregorian calendar), Mendeleev began arranging the elements and comparing them by their atomic weights. He began with a few elements, and over the course of the day his system grew until it encompassed most of the known elements. After he found a consistent arrangement, his printed table appeared in May 1869 in the journal of the Russian Chemical Society.<ref name="Scerri117">Scerri, pp. 117–123</ref> When elements did not appear to fit in the system, he boldly predicted that either valencies or atomic weights had been measured incorrectly, or that there was a missing element yet to be discovered.<ref name=jensenlaw/> In 1871, Mendeleev published a long article, including an updated form of his table, that made his predictions for unknown elements explicit. Mendeleev predicted the properties of three of these unknown elements in detail: as they would be missing heavier homologues of boron, aluminium, and silicon, he named them eka-boron, eka-aluminium, and eka-silicon ("eka" being Sanskrit for "one").<ref name="Scerri117" /><ref name="mendeleev1871">{{cite journal
|last1=Mendeleev
|first1=D.
|title=The natural system of elements and its application to the indication of the properties of undiscovered elements
|journal=Journal of the Russian Chemical Society
|date=1871
|volume=3
|pages=25–56
|url=https://www.knigafund.ru/books/56718/read#page31
|access-date=23 August 2017
|language=ru
|archive-date=13 August 2017
|archive-url=https://web.archive.org/web/20170813142644/https://www.knigafund.ru/books/56718/read#page31
}}</ref>{{rp|45}}
In 1875, the French chemist ], working without knowledge of Mendeleev's prediction, discovered a new element in a sample of the mineral ], and named it gallium. He isolated the element and began determining its properties. Mendeleev, reading de Boisbaudran's publication, sent a letter claiming that gallium was his predicted eka-aluminium. Although Lecoq de Boisbaudran was initially sceptical, and suspected that Mendeleev was trying to take credit for his discovery, he later admitted that Mendeleev was correct.<ref>Scerri, p. 149</ref> In 1879, the Swedish chemist ] discovered a new element, which he named scandium: it turned out to be eka-boron. Eka-silicon was found in 1886 by German chemist ], who named it germanium. The properties of gallium, scandium, and germanium matched what Mendeleev had predicted.<ref>Scerri, p. 151–2</ref> In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law".<ref>{{cite web|last=Rouvray|first=R.|url=https://www.newscientist.com/people/dmitri-mendeleev/|title=Dmitri Mendeleev|website=New Scientist|language=en-US|access-date=19 April 2020|archive-date=15 August 2021|archive-url=https://web.archive.org/web/20210815074119/https://www.newscientist.com/people/dmitri-mendeleev/|url-status=live}}</ref> Even the discovery of the noble gases at the close of the 19th century, which Mendeleev had not predicted, fitted neatly into his scheme as an eighth main group.<ref name="Scerri164">Scerri, pp. 164–169</ref>


Mendeleev nevertheless had some trouble fitting the known lanthanides into his scheme, as they did not exhibit the periodic change in valencies that the other elements did. After much investigation, the Czech chemist ] suggested in 1902 that the lanthanides could all be placed together in one group on the periodic table. He named this the "asteroid hypothesis" as an astronomical analogy: just as there is an ] instead of a single planet between Mars and Jupiter, so the place below yttrium was thought to be occupied by all the lanthanides instead of just one element.<ref name=Thyssen />
where ''Z'' is the atomic number, ''c'' is the ], and ''α'' is the ].<ref>{{cite book|last1=Eisberg|first1=R.|last2=Resnick|first2=R.|year=1974|title=Quantum Physics of Atoms, Molecules, Solids, Nuclei and Particles|publisher=]|isbn=9780471873730}}</ref> Under this model, any element with ''Z'' greater than 137 would require 1s electrons to be traveling faster than the speed of light. Thus, relativistic models must be used for Z > 137.


=== Calculations from the Dirac equation === === Atomic number ===
]]]
The ] ] also has problems for ''Z''&nbsp;>&nbsp;137, for the ground state energy is
After the internal structure of the atom was probed, amateur Dutch physicist ] proposed in 1913 that the nuclear charge determined the placement of elements in the periodic table.<ref name="moseley2010">{{cite magazine |last1=Marshall |first1=J.L. |last2=Marshall |first2=V.R. |date=2010 |title=Rediscovery of the Elements: Moseley and Atomic Numbers |pages=42–47 |magazine=The Hexagon |volume=101 |issue=3 |publisher=] |s2cid=94398490 |url=https://pdfs.semanticscholar.org/afe4/8822cd0871e65dc5401166e7df68dc0ecb7f.pdf |access-date=15 August 2021 |archive-date=16 July 2019 |archive-url=https://web.archive.org/web/20190716215907/https://pdfs.semanticscholar.org/afe4/8822cd0871e65dc5401166e7df68dc0ecb7f.pdf }}</ref><ref>A. van den Broek, '']'', 14, (1913), 32–41</ref> The New Zealand physicist ] coined the word "atomic number" for this nuclear charge.<ref>Scerri, p. 185</ref> In van den Broek's published article he illustrated the first electronic periodic table showing the elements arranged according to the number of their electrons.<ref>A. van den Broek, Die Radioelemente, das periodische System und die Konstitution der Atom, Physik. Zeitsch., 14, 32, (1913).</ref> Rutherford confirmed in his 1914 paper that Bohr had accepted the view of van den Broek.<ref>E. Rutherford, Phil. Mag., 27, 488–499 (Mar. 1914). "This has led to an interesting suggestion by van Broek that the number of units of charge on the nucleus, and consequently the number of external electrons, may be equal to the number of the elements when arranged in order of increasing atomic weight. On this view, the nucleus charges of hydrogen, helium, and carbon are 1, 2, 6 respectively, and so on for the other elements, provided there is no gap due to a missing element. This view has been taken by Bohr in his theory of the constitution of simple atoms and molecules."</ref>


The same year, English physicist ] using ] confirmed van den Broek's proposal experimentally. Moseley determined the value of the nuclear charge of each element from ] to ] and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge.<ref>{{cite book |title=The Periodic Kingdom |author=Atkins, P. W. |author-link=P. W. Atkins |publisher=HarperCollins Publishers, Inc. |year=1995 |page= |isbn=978-0-465-07265-1 |url=https://archive.org/details/periodickingdomj00atki/page/87 }}</ref> Nuclear charge is identical to ] count and determines the value of the ] (''Z'') of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley's research immediately resolved discrepancies between atomic weight and chemical properties; these were cases such as tellurium and iodine, where atomic number increases but atomic weight decreases.<ref name="moseley2010" /> Although Moseley was soon killed in World War I, the Swedish physicist ] continued his work up to ], and established that it was the element with the highest atomic number then known (92).<ref>{{cite journal |last1=Egdell |first1=Russell G. |last2=Bruton |first2=Elizabeth |date=2020 |title=Henry Moseley, X-ray spectroscopy and the periodic table |journal=Philosophical Transactions of the Royal Society A: Mathematical, Physical and Engineering Sciences |volume=378 |issue=2180 |doi=10.1002/chem.202004775|pmid=32811359 |doi-access=free }}</ref> Based on Moseley and Siegbahn's research, it was also known which atomic numbers corresponded to missing elements yet to be found: 43, 61, 72, 75, 85, and 87.<ref name="moseley2010" /> (Element 75 had in fact already been found by Japanese chemist ] in 1908 and named ''nipponium'', but he mistakenly assigned it as element 43 instead of 75 and so his discovery was not generally recognized until later. The contemporarily accepted discovery of element 75 came in 1925, when ], ], and ] independently rediscovered it and gave it its present name, ].)<ref name=nipponium2022>{{cite journal |last1=Hisamatsu |first1=Yoji |last2=Egashira |first2=Kazuhiro |first3=Yoshiteru |last3=Maeno |date=2022 |title=Ogawa's nipponium and its re-assignment to rhenium |journal=Foundations of Chemistry |volume=24 |issue= |pages=15–57 |doi=10.1007/s10698-021-09410-x |doi-access=free }}</ref>
:<math>E=m c^2 \sqrt{1-Z^2 \alpha^2}</math>


The dawn of atomic physics also clarified the situation of ]s. In the ]s of the primordial radioactive elements thorium and uranium, it soon became evident that there were many apparent new elements that had different atomic weights but exactly the same chemical properties. In 1913, ] coined the term "isotope" to describe this situation, and considered isotopes to merely be different forms of the same chemical element. This furthermore clarified discrepancies such as tellurium and iodine: tellurium's natural isotopic composition is weighted towards heavier isotopes than iodine's, but tellurium has a lower atomic number.<ref name=7elements>{{cite book |last=Scerri |first=Eric |author-link= |date=2013 |title=A Tale of Seven Elements |url= |location= |publisher=Oxford University Press |pages=47–53, 115 |isbn=978-0-19-539131-2}}</ref>
where ''m'' is the rest mass of the electron. Although for ''Z''&nbsp;>&nbsp;137, the ] 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 ].<ref>

{{cite book
=== Electron shells ===
|first1=J. D.|last1= Bjorken|first2=S. D.|last2= Drell

|year=1965
The Danish physicist ] applied ]'s idea of quantization to the atom. He concluded that the energy levels of electrons were quantised: only a discrete set of stable energy states were allowed. Bohr then attempted to understand periodicity through electron configurations, surmising in 1913 that the inner electrons should be responsible for the chemical properties of the element.<ref>See Bohr table from 1913 paper below.</ref><ref>Helge Kragh, Aarhus, Lars Vegard, Atomic Structure, and the Periodic System, Bull. Hist. Chem., VOLUME 37, Number 1 (2012), p.43.</ref> In 1913, he produced the first electronic periodic table based on a quantum atom.<ref name="Scerri208">Scerri, pp. 208–218</ref>
|title=Relativistic Quantum Mechanics

|publisher=]
Bohr called his electron shells "rings" in 1913: atomic orbitals within shells did not exist at the time of his planetary model. Bohr explains in Part 3 of his famous 1913 paper that the maximum electrons in a shell is eight, writing, "We see, further, that a ring of {{Var|n}} electrons cannot rotate in a single ring round a nucleus of charge ne unless {{Var|n}} < 8." For smaller atoms, the electron shells would be filled as follows: "rings of electrons will only join if they contain equal numbers of electrons; and that accordingly the numbers of electrons on inner rings will only be 2, 4, 8." However, in larger atoms the innermost shell would contain eight electrons: "on the other hand, the periodic system of the elements strongly suggests that already in neon {{Var|N}} = 10 an inner ring of eight electrons will occur." His proposed electron configurations for the atoms (shown to the right) mostly do not accord with those now known.<ref>Niels Bohr, "On the Constitution of Atoms and Molecules, Part III, Systems containing several nuclei" Philosophical Magazine 26:857--875 (1913)</ref><ref>{{Cite journal|last=Kragh|first=Helge|date=1 January 1979|title=Niels Bohr's Second Atomic Theory|url=https://online.ucpress.edu/hsns/article/doi/10.2307/27757389/47571/Niels-Bohr-s-Second-Atomic-Theory|journal=Historical Studies in the Physical Sciences|language=en|volume=10|pages=123–186|doi=10.2307/27757389|jstor=27757389 |issn=0073-2672}}</ref> They were improved further after the work of ] and ] discovered more quantum numbers.<ref name=7elements/>
|isbn=0031-9228

}}</ref> More accurate calculations taking into account the effects of the finite size of the nucleus indicate that the binding energy first exceeds 2''mc''<sup>2</sup> for''Z''&nbsp;>&nbsp;''Z''<sub>cr</sub>&nbsp;≈&nbsp;173. For ''Z''&nbsp;>&nbsp;''Z''<sub>cr</sub>, if the innermost orbital is not filled, the ] of the nucleus will pull an electron out of the vacuum, resulting in the ].<ref>
{| class="wikitable" style="float:right; font-size:95%; margin:0.5em;"
{{cite doi|10.1016/j.physrep.2009.10.004}}</ref>
|+ Bohr's electron configurations for light elements
|-
! Element !! Electrons per shell
|-
| 4 || 2,2
|-
| 6 || 2,4
|-
| 7 || 4,3
|-
| 8 || 4,2,2
|-
| 9 || 4,4,1
|-
| 10 || 8,2
|-
| 11 || 8,2,1
|-
| 16 || 8,4,2,2
|-
| 18 || 8,8,2
|}

The first one to systematically expand and correct the chemical potentials of Bohr's atomic theory was ] in 1914 and in 1916. Kossel explained that in the periodic table new elements would be created as electrons were added to the outer shell. In Kossel's paper, he writes: <blockquote>This leads to the conclusion that the electrons, which are added further, should be put into concentric rings or shells, on each of which ... only a certain number of electrons—namely, eight in our case—should be arranged. As soon as one ring or shell is completed, a new one has to be started for the next element; the number of electrons, which are most easily accessible, and lie at the outermost periphery, increases again from element to element and, therefore, in the formation of each new shell the chemical periodicity is repeated.<ref>W. Kossel, "Über Molekülbildung als Folge des Atom- baues", Ann. Phys., 1916, 49, 229–362 (237).</ref><ref>Translated in Helge Kragh, Aarhus, Lars Vegard, Atomic Structure, and the Periodic System, Bull. Hist. Chem., VOLUME 37, Number 1 (2012), p.43.</ref></blockquote>

In a 1919 paper, ] postulated the existence of "cells" which we now call orbitals, which could each only contain two electrons each, and these were arranged in "equidistant layers" which we now call shells. He made an exception for the first shell to only contain two electrons.<ref>{{Cite journal |last=Langmuir |first=Irving |author-link=Irving Langmuir |date=June 1919 |title=The Arrangement of Electrons in Atoms and Molecules |url=https://pubs.acs.org/doi/abs/10.1021/ja02227a002 |url-status=live |journal=] |language=en |volume=41 |issue=6 |pages=868–934 |doi=10.1021/ja02227a002 |bibcode=1919JAChS..41..868L |issn=0002-7863 |archive-url=https://web.archive.org/web/20210126003324/https://zenodo.org/record/1429026 |archive-date=26 January 2021 |access-date=22 October 2021}}</ref> The chemist ] suggested in 1921 that eight and eighteen electrons in a shell form stable configurations. Bury proposed that the electron configurations in transitional elements depended upon the valence electrons in their outer shell.<ref name="Bury">{{Cite journal |last=Bury |first=Charles R. |author-link=Charles Rugeley Bury |date=July 1921 |title=Langmuir's Theory of the Arrangement of Electrons in Atoms and Molecules |url=https://pubs.acs.org/doi/abs/10.1021/ja01440a023 |url-status=live |journal=] |language=en |volume=43 |issue=7 |pages=1602–1609 |doi=10.1021/ja01440a023 |bibcode=1921JAChS..43.1602B |issn=0002-7863 |archive-url=https://web.archive.org/web/20211030145903/https://zenodo.org/record/1428812 |archive-date=30 October 2021 |access-date=22 October 2021}}</ref> He introduced the word ''transition'' to describe the elements now known as ]s or transition elements.<ref name="Jensen2003">{{cite journal|last=Jensen|first=William B.|year=2003|title=The Place of Zinc, Cadmium, and Mercury in the Periodic Table|url=https://www.uv.es/~borrasj/ingenieria_web/temas/tema_1/lecturas_comp/p952.pdf|journal=Journal of Chemical Education|volume=80|issue=8|pages=952–961|bibcode=2003JChEd..80..952J|doi=10.1021/ed080p952|quote=The first use of the term "transition" in its modern electronic sense appears to be due to the British chemist C. R.Bury, who first used the term in his 1921 paper on the electronic structure of atoms and the periodic table|access-date=18 September 2021|archive-date=19 April 2012|archive-url=https://web.archive.org/web/20120419082806/https://www.uv.es/~borrasj/ingenieria_web/temas/tema_1/lecturas_comp/p952.pdf|url-status=live}}</ref> Bohr's theory was vindicated by the discovery of element 72: ] claimed to have discovered it as the ] ''celtium'', but Bury and Bohr had predicted that element 72 could not be a rare earth element and had to be a homologue of ]. ] and ] searched for the element in zirconium ores and found element 72, which they named ] after Bohr's hometown of ] (''Hafnia'' in Latin).<ref name="CosterHevesy1923">{{cite journal|journal = Nature|volume = 111|page=79|date=1923|doi = 10.1038/111079a0|title = On the Missing Element of Atomic Number 72|first = D.|last = Coster|author2=Hevesy, G.|issue=2777|bibcode=1923Natur.111...79C|doi-access = free}}</ref><ref>{{cite journal|title = Hafnium|url = http://www.jce.divched.org/Journal/Issues/1982/Mar/jceSubscriber/JCE1982p0242.pdf|journal = Journal of Chemical Education|last = Fernelius|first = W. C.|date = 1982|page = 242|doi = 10.1021/ed059p242|volume = 59|issue = 3|bibcode = 1982JChEd..59..242F|access-date = 3 September 2009|archive-date = 15 March 2020|archive-url = https://web.archive.org/web/20200315031648/http://www.jce.divched.org/Journal/Issues/1982/Mar/jceSubscriber/JCE1982p0242.pdf|url-status = dead}}</ref> Urbain's celtium proved to be simply purified ] (element 71).<ref>{{cite journal |last1=Burdette |first1=Shawn C. |last2=Thornton |first2=Brett F. |date=2018 |title=Hafnium the lutécium I used to be |url=https://www.nature.com/articles/s41557-018-0140-6 |journal=Nature Chemistry |volume=10 |issue= 10|pages=1074 |doi=10.1038/s41557-018-0140-6 |pmid=30237529 |bibcode=2018NatCh..10.1074B |access-date=8 February 2024}}</ref> Hafnium and rhenium thus became the last stable elements to be discovered.<ref name=7elements/>

Prompted by Bohr, ] took up the problem of electron configurations in 1923. Pauli extended Bohr's scheme to use four ]s, and formulated his ] which stated that no two electrons could have the same four quantum numbers. This explained the lengths of the periods in the periodic table (2, 8, 18, and 32), which corresponded to the number of electrons that each shell could occupy.<ref name="Scerri218">Scerri, pp. 218–23</ref> In 1925, ] arrived at configurations close to the modern ones.<ref>{{cite journal |last1=Jensen |first1=William B. |date=2007 |title=The Origin of the s, p, d, f Orbital Labels |url=https://www.che.uc.edu/jensen/w.%20b.%20jensen/reprints/137.%20s,%20p,%20d,%20f.pdf |journal=Journal of Chemical Education |volume=84 |issue=5 |pages=757–8 |doi=10.1021/ed084p757 |bibcode=2007JChEd..84..757J |archive-url=https://web.archive.org/web/20181123140649/https://www.che.uc.edu/jensen/w.%20b.%20jensen/reprints/137.%20s,%20p,%20d,%20f.pdf |access-date=15 August 2021|archive-date=23 November 2018 }}</ref> As a result of these advances, periodicity became based on the number of chemically active or valence electrons rather than by the valences of the elements.<ref name=jensenlaw/> The ] that describes the electron configurations of the elements was first empirically observed by ] in 1926,<ref name="Goudsmit">{{cite journal |title=The Order of Electron Shells in Ionized Atoms |last1=Goudsmit |first1=S. A. |last2=Richards |first2=Paul I. |journal=] |pages=664–671 (with correction on p&nbsp;906) |volume=51 |issue=4 |date=1964 |url=https://www.pnas.org/content/51/4/664.full.pdf |bibcode=1964PNAS...51..664G |doi=10.1073/pnas.51.4.664 |pmid=16591167 |doi-access=free |pmc=300183 |access-date=15 August 2021 |archive-date=10 October 2017 |archive-url=https://web.archive.org/web/20171010113455/https://www.pnas.org/content/51/4/664.full.pdf |url-status=live }}</ref> though the first to publish it was ] in 1930.<ref>{{cite journal |last1=Karapetoff |first1=Vladimir |date=1930 |title=A chart of consecutive sets of electronic orbits within atoms of chemical elements |url= |journal=Journal of the Franklin Institute |volume=210 |issue=5 |pages=609–624 |doi=10.1016/S0016-0032(30)91131-3 }}</ref><ref name=Ostro>{{cite journal |last1=Ostrovsky |first1=Valentin N. |date=2003 |title=Physical Explanation of the Periodic Table |url= |journal=Annals of the New York Academy of Sciences |volume=988 |issue=1 |pages=182–192 |doi=10.1111/j.1749-6632.2003.tb06097.x |pmid=12796101 |bibcode=2003NYASA.988..182O |s2cid=21629328 }}</ref> In 1961, ] derived the first part of the Madelung rule (that orbitals fill in order of increasing ''n'' + ℓ) from the ];<ref>{{cite journal |last1=Klechkovskii |first1=V.M. |title=Justification of the Rule for Successive Filling of (n+l) Groups |journal=Journal of Experimental and Theoretical Physics |date=1962 |volume=14 |issue=2 |page=334 |url=http://jetp.ras.ru/cgi-bin/e/index?t=&au=+Klechkovskii&yf=2022&yt=2022&se=1&a=s |access-date=23 June 2022}}</ref> the complete rule was derived from a similar potential in 1971 by Yury N. Demkov and Valentin N. Ostrovsky.<ref name=DO>{{cite journal |last1=Demkov |first1=Yury N. |last2=Ostrovsky |first2=Valentin N. |date=1972 |title=n+l Filling Rule in the Periodic System and Focusing Potentials |url=http://jetp.ras.ru/cgi-bin/e/index/e/35/1/p66?a=list |journal=Journal of Experimental and Theoretical Physics |volume=35 |issue=1 |pages=66–69 |doi= |bibcode=1972JETP...35...66D |access-date=25 November 2022}}</ref>{{efn|Demkov and Ostrovsky consider the potential <math>U_{1/2}(r) = -\frac{2v}{rR(r+R)^2}</math> where <math>R</math> and <math>v</math> are constant parameters; this approaches a ] for small <math>r</math>. When <math>v</math> satisfies the condition <math>v=v_N=\frac{1}{4}R^2 N(N+1)</math>, where <math>N=n+l</math>, the zero-energy solutions to the ] for this potential can be described analytically with ]. As <math>v</math> passes through each of these values, a manifold containing all states with that value of <math>N</math> arises at zero energy and then becomes bound, recovering the Madelung order. Perturbation-theory considerations show that states with smaller <math>n</math> have lower energy, and that the s-orbitals (with <math>l=0</math>) have their energies approaching the next <math>n+l</math> group.<ref name=DO/><ref name=shattered/>}}

]
The quantum theory clarified the transition metals and lanthanides as forming their own separate groups, transitional between the main groups, although some chemists had already proposed tables showing them this way before then: the English chemist Henry Bassett did so in 1892, the Danish chemist ] in 1895, and the Swiss chemist ] in 1905. Bohr used Thomsen's form in his 1922 Nobel Lecture; Werner's form is very similar to the modern 32-column form. In particular, this supplanted Brauner's asteroidal hypothesis.<ref name="Thyssen">{{cite book|last1=Thyssen|first1=P.|last2=Binnemans|first2=K.|editor1-last=Gschneidner|editor1-first= K. A. Jr.|editor2-last=Bünzli|editor2-first=J-C.G|editor3-last=Vecharsky|editor3-first=Bünzli|date=2011|chapter=Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis|title=Handbook on the Physics and Chemistry of Rare Earths|publisher=Elsevier|location=Amsterdam|volume=41|pages=1–93|isbn=978-0-444-53590-0|doi=10.1016/B978-0-444-53590-0.00001-7}}</ref>

The exact position of the lanthanides, and thus the composition of ], remained under dispute for decades longer because their electron configurations were initially measured incorrectly.<ref name=Jensen1982/><ref name="PTSS">Scerri, pp. 392−401</ref> On chemical grounds Bassett, Werner, and Bury grouped scandium and yttrium with lutetium rather than lanthanum (the former two left an empty space below yttrium as lutetium had not yet been discovered).<ref name=Thyssen/><ref name=Bury/> Hund assumed in 1927 that all the lanthanide atoms had configuration 4f<sup>0−14</sup>5d<sup>1</sup>6s<sup>2</sup>, on account of their prevailing trivalency. It is now known that the relationship between chemistry and electron configuration is more complicated than that.{{efn|For example, the early actinides continue to behave more like the d-block transition metals in their propensity towards high oxidation states all the way from actinium to uranium, even though it is actually only actinium and thorium that have d-block-like configurations in the gas phase; f-electrons appear already at protactinium.<ref name=johnson/> Uranium's actual configuration of 5f<sup>3</sup>6d<sup>1</sup>7s<sup>2</sup> is in fact analogous to that Hund assumed for the lanthanides, but uranium does not favour the trivalent state, preferring to be tetravalent or hexavalent.<ref name=rareearths/> On the other hand, lanthanide-like configurations for the actinides begin at plutonium, but the shift towards lanthanide-like behaviour is only clear at curium: the elements between uranium and curium form a transition from transition-metal-like behaviour to lanthanide-like behaviour.<ref name=johnson/> Thus chemical behaviour and electron configuration do not exactly match each other.<ref name=johnson/>}}<ref name=rareearths>{{cite book |last1=Jørgensen |first1=Christian Klixbüll |editor1-last=Gschneidner Jr. |editor1-first=Karl A. |editor2-last=Eyring |editor2-first=Leroy |date=1988 |title=Handbook on the Physics and Chemistry of Rare Earths |publisher=Elsevier |volume=11 |pages=197–292 |chapter=Influence of Rare Earths on Chemical Understanding and Classification |isbn=978-0-444-87080-3}}</ref> Early spectroscopic evidence seemed to confirm these configurations, and thus the periodic table was structured to have group 3 as scandium, yttrium, lanthanum, and actinium, with fourteen f-elements breaking up the d-block between lanthanum and hafnium.<ref name=Jensen1982/> But it was later discovered that this is only true for four of the fifteen lanthanides (lanthanum, cerium, gadolinium, and lutetium), and that the other lanthanide atoms do not have a d-electron. In particular, ytterbium completes the 4f shell and thus Soviet physicists Lev Landau and Evgeny Lifshitz noted in 1948 that lutetium is correctly regarded as a d-block rather than an f-block element;<ref name=Landau/> that bulk lanthanum is an f-metal was first suggested by ] in 1963, on the grounds of its low-temperature ].<ref name=Kondo>{{cite journal |last1=Kondō |first1=Jun |date=January 1963 |title=Superconductivity in Transition Metals |url= |journal=Progress of Theoretical Physics |volume=29 |issue=1 |pages=1–9 |doi=10.1143/PTP.29.1 |bibcode=1963PThPh..29....1K |doi-access=free }}</ref> This clarified the importance of looking at low-lying excited states of atoms that can play a role in chemical environments when classifying elements by block and positioning them on the table.<ref name=Hamilton/><ref name=JensenLr/><ref name=Jensen1982/> Many authors subsequently rediscovered this correction based on physical, chemical, and electronic concerns and applied it to all the relevant elements, thus making group 3 contain scandium, yttrium, lutetium, and lawrencium<ref name=Hamilton/><ref name=Fluck/><ref name=PTSS/> and having lanthanum through ytterbium and actinium through nobelium as the f-block rows:<ref name=Hamilton/><ref name=Fluck/> this corrected version achieves consistency with the Madelung rule and vindicates Bassett, Werner, and Bury's initial chemical placement.<ref name=Thyssen/>

In 1988, IUPAC released a report supporting this composition of group 3,<ref name=Fluck/> a decision that was reaffirmed in 2021.<ref name="2021IUPAC">{{cite journal |last1=Scerri |first1=Eric |date=18 January 2021 |title=Provisional Report on Discussions on Group 3 of the Periodic Table |url=https://iupac.org/wp-content/uploads/2021/04/ChemInt_Jan2021_PP.pdf |journal=Chemistry International |volume=43 |issue=1 |pages=31–34 |doi=10.1515/ci-2021-0115 |s2cid=231694898 |access-date=9 April 2021 |archive-date=13 April 2021 |archive-url=https://web.archive.org/web/20210413150110/https://iupac.org/wp-content/uploads/2021/04/ChemInt_Jan2021_PP.pdf |url-status=live }}</ref> Variation can still be found in textbooks on the composition of group 3,<ref name=2015IUPAC/> and some argumentation against this format is still published today,<ref name=Jensen2015/> but chemists and physicists who have considered the matter largely agree on group 3 containing scandium, yttrium, lutetium, and lawrencium and challenge the counterarguments as being inconsistent.<ref name=Jensen2015/>

=== Synthetic elements ===
]
By 1936, the pool of missing elements from hydrogen to uranium had shrunk to four: elements 43, 61, 85, and 87 remained missing. Element 43 eventually became the first element to be synthesized artificially via nuclear reactions rather than discovered in nature. It was discovered in 1937 by Italian chemists ] and ], who named their discovery ], after the Greek word for "artificial".<ref>Scerri, pp. 313–321</ref> Elements 61 (]) and 85 (]) were likewise produced artificially in 1945 and 1940 respectively; element 87 (]) became the last element to be discovered in nature, by French chemist ] in 1939.<ref>Scerri, pp. 322–340</ref>{{efn|Technetium, promethium, astatine, neptunium, and plutonium were eventually discovered to occur in nature as well, albeit in tiny traces. See ].}} The elements beyond uranium were likewise discovered artificially, starting with ] and ]'s 1940 discovery of ] (via bombardment of uranium with neutrons).<ref name="Scerri354">Scerri, p. 354–6</ref> ] and his team at the ] (LBNL) continued discovering transuranium elements, starting with ] in 1941, and discovered that contrary to previous thinking, the elements from actinium onwards were congeners of the lanthanides rather than transition metals.<ref name=Seaborg /> Bassett (1892), Werner (1905), and the French engineer ] (1928) had previously suggested this, but their ideas did not then receive general acceptance.<ref name=Thyssen /> Seaborg thus called them the actinides.<ref name="Seaborg">{{cite web |url=https://fas.org/sgp/othergov/doe/lanl/orgs/nmt/97summer.pdf |title=Source of the Actinide Concept |last=Seaborg |first=Glenn T. |date=1997 |website=fas.org |publisher=Los Alamos National Laboratory |access-date=28 March 2021 |archive-date=15 August 2021 |archive-url=https://web.archive.org/web/20210815074120/https://fas.org/sgp/othergov/doe/lanl/orgs/nmt/97summer.pdf |url-status=live }}</ref> Elements up to 101 (named mendelevium in honour of Mendeleev) were synthesized up to 1955, either through neutron or alpha-particle irradiation, or in nuclear explosions in the cases of 99 (einsteinium) and 100 (fermium).<ref name=Scerri354/>

A significant controversy arose with elements 102 through 106 in the 1960s and 1970s, as competition arose between the LBNL team (now led by ]) and a team of Soviet scientists at the ] (JINR) led by ]. Each team claimed discovery, and in some cases each proposed their own name for the element, creating an ] that lasted decades. These elements were made by bombardment of actinides with light ions.<ref>Scerri, pp. 356–9</ref> IUPAC at first adopted a hands-off approach, preferring to wait and see if a consensus would be forthcoming. But as it was also the height of the ], it became clear that this would not happen. As such, IUPAC and the ] (IUPAP) created a ] (TWG, fermium being element 100) in 1985 to set out criteria for discovery,<ref>{{cite journal |last1=Öhrström |first1=Lars |last2=Holden |first2=Norman E. |date=2016 |title=The Three-letter Element Symbols |journal=Chemistry International |volume=38 |issue=2 |pages=4–8 |doi=10.1515/ci-2016-0204 |s2cid=124737708 |doi-access=free }}</ref> which were published in 1991.<ref>{{cite journal |last1=Wapstra |first1=A. H. |date=1991 |title=Criteria that must be satisfied for the discovery of a new chemical element to be recognized |url=https://old.iupac.org/reports/1991/6306wapstra/index.html |journal=Pure and Applied Chemistry |volume=63 |issue=6 |pages=879–886 |doi=10.1351/pac199163060879 |s2cid=95737691 |access-date=18 October 2022}}</ref> After some further controversy, these elements received their final names in 1997, including seaborgium (106) in honour of Seaborg.<ref>{{cite journal | doi=10.1351/pac199769122471|title=Names and symbols of transfermium elements (IUPAC Recommendations 1997) | year=1997 | journal=Pure and Applied Chemistry | volume=69 | pages=2471–2474 | issue=12| doi-access=free }}</ref>

]
The TWG's criteria were used to arbitrate later element discovery claims from LBNL and JINR, as well as from research institutes in Germany (]) and Japan (]).<ref>{{cite journal |last1=Hofmann |first1=Sigurd |date=2019 |title=Criteria for New Element Discovery |journal=Chemistry International |volume=41 |issue=1 |pages=10–15 |doi=10.1515/ci-2019-0103|doi-access=free }}</ref> Currently, consideration of discovery claims is performed by a ]. After priority was assigned, the elements were officially added to the periodic table, and the discoverers were invited to propose their names.<ref name="IUPAC-redbook" /> By 2016, this had occurred for all elements up to 118, therefore completing the periodic table's first seven rows.<ref name="IUPAC-redbook">{{cite web |url=https://iupac.org/what-we-do/periodic-table-of-elements/ |title=Periodic Table of Elements |author=<!--Not stated--> |date=2021 |website=iupac.org |publisher=IUPAC |access-date=3 April 2021 |archive-date=10 April 2016 |archive-url=https://web.archive.org/web/20160410043726/https://iupac.org/what-we-do/periodic-table-of-elements/ |url-status=live }}</ref><ref name="finally">{{cite journal|last=Scerri|first=E.|author-link=Eric Scerri|year=2012|journal=Chemistry International|volume=34|issue=4|url=https://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|archive-url=https://web.archive.org/web/20170705051357/https://www.iupac.org/publications/ci/2012/3404/ud.html|archive-date=5 July 2017|doi=10.1515/ci.2012.34.4.28|doi-access=free}}</ref> The discoveries of elements beyond 106 were made possible by techniques devised by ] at the JINR: cold fusion (bombardment of lead and bismuth by heavy ions) made possible the 1981–2004 discoveries of elements 107 through 112 at GSI and 113 at Riken, and he led the JINR team (in collaboration with American scientists) to discover elements 114 through 118 using hot fusion (bombardment of actinides by calcium ions) in 1998–2010.<ref>Scerri, pp. 356–363</ref><ref name="Chapman">{{cite journal|last1=Chapman|first1=Kit|title=What it takes to make a new element|journal=]|date=30 November 2016|url=https://www.chemistryworld.com/what-it-takes-to-make-a-new-element/1017677.article|publisher=]|access-date=22 March 2022|archive-date=28 October 2017|archive-url=https://web.archive.org/web/20171028122035/https://www.chemistryworld.com/what-it-takes-to-make-a-new-element/1017677.article|url-status=live }}</ref> The heaviest known element, oganesson (118), is named in Oganessian's honour. Element 114 is named flerovium in honour of his predecessor and mentor Flyorov.<ref name=Chapman/>

In celebration of the periodic table's 150th anniversary, the ] declared the year 2019 as the International Year of the Periodic Table, celebrating "one of the most significant achievements in science".<ref name=":1">{{Cite news|url=https://www.bbc.com/news/science-environment-47008289|title=150 years of the periodic table: Test your knowledge |last=Briggs|first=Helen|date=29 January 2019|access-date=8 February 2019|language=en-GB|archive-url=https://web.archive.org/web/20190209210210/https://www.bbc.com/news/science-environment-47008289|archive-date=9 February 2019|url-status=live}}</ref> The discovery criteria set down by the TWG were updated in 2020 in response to experimental and theoretical progress that had not been foreseen in 1991.<ref>{{cite journal |last1=Hofmann |first1=Sigurd |last2=Dmitriev |first2=Sergey N. |last3=Fahlander |first3=Claes |last4=Gates |first4=Jacklyn M. |last5=Roberto |first5=James B. |last6=Sakai |first6=Hideyuki |date=4 August 2020 |title=On the discovery of new elements (IUPAC/IUPAP Report) |s2cid-access=free |journal=Pure and Applied Chemistry |volume=92 |issue=9 |pages=1387–1446 |doi=10.1515/pac-2020-2926 |s2cid=225377737 |doi-access=free }}</ref> Today, the periodic table is among the most recognisable icons of chemistry.<ref name="Lemonick" /> IUPAC is involved today with many processes relating to the periodic table: the recognition and naming of new elements, recommending group numbers and collective names, and the updating of atomic weights.<ref name="IUPAC-redbook" />

== Future extension beyond the seventh period ==
{{Main|Extended periodic table}}
{{See also|Island of stability}}
]
The most recently named elements – nihonium (113), moscovium (115), tennessine (117), and oganesson (118) – completed the seventh row of the periodic table.<ref name="IUPAC-redbook" /> Future elements would have to begin an ]. These elements may be referred to either by their atomic numbers (e.g. "]"), or by the IUPAC ]s adopted in 1978, which directly relate to the atomic numbers (e.g. "unhexquadium" for element 164, derived from Latin ''unus'' "one", Greek ''hexa '' "six", Latin ''quadra'' "four", and the traditional ''-ium'' suffix for metallic elements).<ref name="IUPAC-redbook" /> All attempts to synthesize such elements have failed so far. An attempt to make element 119 has been ongoing since 2018 at the Riken research institute in Japan. The LBNL in the United States, the JINR in Russia, and the Heavy Ion Research Facility in ] (HIRFL) in China also plan to make their own attempts at synthesizing the first few period 8 elements.<ref name="nature2019">{{cite journal |last=Ball |first=P. |title=Extreme chemistry: experiments at the edge of the periodic table |date=2019 |journal=Nature |volume=565 |issue=7741 |pages=552–555 |issn=1476-4687 |doi=10.1038/d41586-019-00285-9|pmid=30700884 |bibcode=2019Natur.565..552B |doi-access=free }}</ref><ref name="SHEfactory">{{cite conference |url=https://www.epj-conferences.org/articles/epjconf/pdf/2016/26/epjconf-NS160-08001.pdf |title=Status and perspectives of the Dubna superheavy element factory |last1=Dmitriev |first1=Sergey |last2=Itkis |first2=Mikhail |last3=Oganessian |first3=Yuri |date=2016 |conference=Nobel Symposium NS160 – Chemistry and Physics of Heavy and Superheavy Elements |doi=10.1051/epjconf/201613108001 |access-date=15 August 2021 |archive-date=28 August 2021 |archive-url=https://web.archive.org/web/20210828071031/https://www.epj-conferences.org/articles/epjconf/pdf/2016/26/epjconf-NS160-08001.pdf |url-status=live }}</ref><ref>{{cite web |url=https://www.jinr.ru/posts/how-are-new-chemical-elements-born/ |title=How are new chemical elements born? |last1=Sokolova |first1=Svetlana |last2=Popeko |first2=Andrei |date=24 May 2021 |website=jinr.ru |publisher=JINR |access-date=4 November 2021 |quote= |archive-date=4 November 2021 |archive-url=https://web.archive.org/web/20211104173902/https://www.jinr.ru/posts/how-are-new-chemical-elements-born/ |url-status=live }}</ref><ref>{{cite news |last=Chapman |first=Kit |date=10 October 2023 |title=Berkeley Lab to lead US hunt for element 120 after breakdown of collaboration with Russia |url=https://www.chemistryworld.com/news/berkeley-lab-to-lead-us-hunt-for-element-120-after-breakdown-of-collaboration-with-russia/4018207.article |work=Chemistry World |location= |access-date=20 October 2023}}</ref><ref>{{cite web |url=https://physicalsciences.lbl.gov/2023/10/16/berkeley-lab-to-test-new-approach-to-making-superheavy-elements/ |title=Berkeley Lab to Test New Approach to Making Superheavy Elements |last=Biron |first=Lauren |date=16 October 2023 |website=lbl.gov |publisher=] |access-date=20 October 2023 |quote=}}</ref><ref>{{cite journal |last1=Gan |first1=Z. G. |last2=Huang |first2=W. X. |last3=Zhang |first3=Z. Y. |last4=Zhou |first4=X. H. |last5=Xu |first5=H. S. |date=2022 |title=Results and perspectives for study of heavy and super-heavy nuclei and elements at IMP/CAS |url= |journal=The European Physical Journal A |volume=58 |issue=158 |pages= |doi=10.1140/epja/s10050-022-00811-w |bibcode=2022EPJA...58..158G |access-date=}}</ref>

If the eighth period followed the pattern set by the earlier periods, then it would contain fifty elements, filling the 8s, {{Not a typo|5g}}, 6f, 7d, and finally 8p subshells in that order. But by this point, relativistic effects should result in significant deviations from the Madelung rule. Various different models have been suggested for the configurations of eighth-period elements, as well as how to show the results in a periodic table. All agree that the eighth period should begin like the previous ones with two 8s elements, 119 and ]. However, after that the massive energetic overlaps between the {{Not a typo|5g}}, 6f, 7d, and 8p subshells means that they all begin to fill together, and it is not clear how to separate out specific {{not a typo|5g}} and 6f series.<ref name="nefedov">{{cite journal |last1=Nefedov |first1=V.I. |last2=Trzhaskovskaya |first2=M.B. |last3=Yarzhemskii |first3=V.G. |title=Electronic Configurations and the Periodic Table for Superheavy Elements |journal=Doklady Physical Chemistry |date=2006 |volume=408 |issue=2 |pages=149–151 |doi=10.1134/S0012501606060029 |s2cid=95738861 |issn=0012-5016 |url=https://www.primefan.ru/stuff/chem/nefedov.pdf |access-date=15 August 2021 |archive-date=13 October 2016 |archive-url=https://web.archive.org/web/20161013113837/https://www.primefan.ru/stuff/chem/nefedov.pdf |url-status=live }}</ref><ref name=recentattempts>{{cite journal |last1=Scerri |first1=Eric |date=2020 |title=Recent attempts to change the periodic table |journal=Philosophical Transactions of the Royal Society A |volume=378 |issue=2180 |doi=10.1098/rsta.2019.0300|pmid=32811365 |bibcode=2020RSPTA.37890300S |s2cid=221136189 |doi-access=free }}</ref><ref>{{cite journal|doi=10.2307/3963006|last=Frazier|first=K.|title=Superheavy Elements|journal=Science News|volume=113|issue=15|pages=236–38|year=1978|jstor=3963006}}</ref><ref name="Fricke">{{cite journal |last1=Fricke |first1=B. |last2=Greiner |first2=W. |last3=Waber |first3=J. T. |year=1971 |title=The continuation of the periodic table up to Z = 172. The chemistry of superheavy elements |journal=Theoretica Chimica Acta |volume=21 |issue=3 |pages=235–60 |doi=10.1007/BF01172015 |s2cid=117157377 }}</ref><ref name="PT172">{{Cite journal|last1=Pyykkö|first1=P.|author-link=Pekka Pyykkö|title=A suggested periodic table up to Z ≤ 172, based on Dirac–Fock calculations on atoms and ions|journal=Physical Chemistry Chemical Physics|volume=13|issue=1|pages=161–68|year=2011|pmid=20967377|doi=10.1039/c0cp01575j|bibcode=2011PCCP...13..161P|s2cid=31590563}}</ref> Elements ] through 156 thus do not fit well as chemical analogues of any previous group in the earlier parts of the table,<ref name=actrev/> although they have sometimes been placed as {{not a typo|5g}}, 6f, and other series to formally reflect their electron configurations.<ref name=actrev/> Eric Scerri has raised the question of whether an extended periodic table should take into account the failure of the Madelung rule in this region, or if such exceptions should be ignored.<ref name=recentattempts /> The shell structure may also be fairly formal at this point: already the electron distribution in an oganesson atom is expected to be rather uniform, with no discernible shell structure.<ref name="oganesson-elf">{{cite journal| journal=Phys. Rev. Lett.| volume=120| issue=5| page=053001| date=2018| title=Electron and Nucleon Localization Functions of Oganesson: Approaching the Thomas-Fermi Limit| first1=Paul |last1=Jerabek |first2=Bastian |last2=Schuetrumpf |first3=Peter |last3=Schwerdtfeger |first4=Witold |last4=Nazarewicz| doi=10.1103/PhysRevLett.120.053001| pmid=29481184| arxiv = 1707.08710 | bibcode = 2018PhRvL.120e3001J| s2cid=3575243}}</ref>

The situation from elements 157 to 172 should return to normalcy and be more reminiscent of the earlier rows.<ref name=BFricke1977/> The heavy p-shells are split by the ]: one p-orbital (p<sub>1/2</sub>) is more stabilized, and the other two (p<sub>3/2</sub>) are destabilized. (Such shifts in the quantum numbers happen for all types of shells, but it makes the biggest difference to the order for the p-shells.) It is likely that by element 157, the filled 8s and 8p<sub>1/2</sub> shells with four electrons in total have sunk into the core. Beyond the core, the next orbitals are 7d and 9s at similar energies, followed by 9p<sub>1/2</sub> and 8p<sub>3/2</sub> at similar energies, and then a large gap.<ref name="BFricke1977">{{cite journal |last1=Fricke |first1=Burkhard |year=1977 |title=Dirac–Fock–Slater calculations for the elements Z = 100, fermium, to Z = 173 |journal=Recent Impact of Physics on Inorganic Chemistry |volume=19 |pages=83–192 |doi=10.1016/0092-640X(77)90010-9 |url=http://kobra.bibliothek.uni-kassel.de/bitstream/urn:nbn:de:hebis:34-2008071622807/1/Fricke_Dirac_1977.pdf |access-date=25 February 2016 |bibcode=1977ADNDT..19...83F |archive-date=22 March 2016 |archive-url=https://web.archive.org/web/20160322072636/http://kobra.bibliothek.uni-kassel.de/bitstream/urn:nbn:de:hebis:34-2008071622807/1/Fricke_Dirac_1977.pdf |url-status=dead }}</ref> Thus, the 9s and 9p<sub>1/2</sub> orbitals in essence replace the 8s and 8p<sub>1/2</sub> ones, making elements 157–172 probably chemically analogous to groups 3–18: for example, element 164 would appear two places below lead in group 14 under the usual pattern, but is calculated to be very analogous to palladium in group 10 instead.<ref name=rareearths/><ref name=Fricke/><ref name=nefedov/><ref name="BFricke">{{Cite journal |last1=Fricke |first1=Burkhard |year=1975 |title=Superheavy elements: a prediction of their chemical and physical properties |journal=Recent Impact of Physics on Inorganic Chemistry |volume=21 |pages= |doi=10.1007/BFb0116498 |url=https://archive.org/details/recentimpactofph0000unse/page/89 |access-date=4 October 2013 |series=Structure and Bonding |isbn=978-3-540-07109-9 }}</ref><ref name=actrev>{{cite journal |last1=Fricke |first1=Burkhard |last2=Waber |first2=J. T. |date=1971 |title=Theoretical Predictions of the Chemistry of Superheavy Elements: Continuation of the Periodic Table up to Z{{=}}184 |url=https://kobra.uni-kassel.de/bitstream/handle/123456789/2008100124269/Fricke_theoretical_1971.pdf |journal=Actinides Reviews |volume=1 |issue= |pages=433–485 |doi= |access-date=5 January 2024}}</ref> Thus, it takes fifty-four elements rather than fifty to reach the next noble element after 118.<ref name=wothers>{{cite book |last=Wothers |first=Peter |author-link= |date=2019 |title=Antimony, Gold, and Jupiter's Wolf |url= |location= |publisher=Oxford University Press |page=vii |isbn=978-0-19-965272-3}}</ref> However, while these conclusions about elements 157 through 172's chemistry are generally agreed by models,<ref name=actrev/><ref name=nefedov/> there is disagreement on whether the periodic table should be drawn to reflect chemical analogies, or if it should reflect likely formal electron configurations, which should be quite different from earlier periods and are not agreed between sources. Discussion about the format of the eighth row thus continues.<ref name=nefedov/><ref name=Fricke/><ref name=PT172/><ref name=smits>{{cite journal |last1=Smits |first1=Odile R. |last2=Düllmann |first2=Christoph E. |last3=Indelicato |first3=Paul |last4=Nazarewicz |first4=Witold |last5=Schwerdtfeger |first5=Peter |date=2023 |title=The quest for superheavy elements and the limit of the periodic table |url= |journal=Nature Reviews Physics |volume= 6|issue= 2|pages= 86–98|doi=10.1038/s42254-023-00668-y |s2cid=266276980 |access-date=}}</ref>

Beyond element 172, calculation is complicated by the 1s electron energy level becoming ]. Such a situation does have a physical interpretation and does not in itself pose an electronic limit to the periodic table, but the correct way to incorporate such states into multi-electron calculations is still an open question needing to be solved to calculate the periodic table's structure beyond this point.<ref name=gamowstates>{{cite journal |last1=Smits |first1=O. R. |last2=Indelicato |first2=P. |first3=W. |last3=Nazarewicz |first4=M. |last4=Piibeleht |first5=P. |last5=Schwerdtfeger |date=2023 |title=Pushing the limits of the periodic table—A review on atomic relativistic electronic structure theory and calculations for the superheavy elements |url= |journal=Physics Reports |volume=1035 |issue= |pages=1–57 |doi=10.1016/j.physrep.2023.09.004 |arxiv=2301.02553 |bibcode=2023PhR..1035....1S |access-date=}}</ref>

Nuclear stability will likely prove a decisive factor constraining the number of possible elements. It depends on the balance between the electric repulsion between protons and the strong force binding protons and neutrons together.<ref>{{cite journal |last1=Pershina |first1=Valeria |date=2020 |title=Relativistic effects on the electronic structure of the heaviest elements. Is the Periodic Table endless? |url=https://comptes-rendus.academie-sciences.fr/chimie/article/CRCHIM_2020__23_3_255_0.pdf |journal=Comptes Rendus Chimie |volume=23 |issue=3 |pages=255–265 |doi=10.5802/crchim.25 |s2cid=222225772 |access-date=28 March 2021 |archive-date=11 December 2020 |archive-url=https://web.archive.org/web/20201211103843/https://comptes-rendus.academie-sciences.fr/chimie/article/CRCHIM_2020__23_3_255_0.pdf |url-status=live }}</ref> Protons and neutrons are arranged in ], just like electrons, and so a closed shell can significantly increase stability: the known superheavy nuclei exist because of such a shell closure, probably at around 114–] protons and 184 neutrons.<ref name=gamowstates/> They are probably close to a predicted ], where superheavy nuclides should be more long-lived than expected: predictions for the longest-lived nuclides on the island range from microseconds to millions of years.<ref name=smits/><ref name="physorg">{{cite web |url=https://newscenter.lbl.gov/2009/09/24/114-confirmed/ |title=Superheavy Element 114 Confirmed: A Stepping Stone to the Island of Stability |date=2009 |access-date=23 October 2019 |publisher=] |archive-date=20 July 2019 |archive-url=https://web.archive.org/web/20190720200414/https://newscenter.lbl.gov/2009/09/24/114-confirmed/ |url-status=live }}</ref><ref name="nuclei">{{cite journal |last=Oganessian |first=Yu. Ts. |year=2012 |title=Nuclei in the "Island of Stability" of Superheavy Elements |journal=] |volume=337 |issue=1 |page=012005 |bibcode=2012JPhCS.337a2005O |doi=10.1088/1742-6596/337/1/012005|doi-access=free }}</ref> It should nonetheless be noted that these are essentially extrapolations into an unknown part of the chart of nuclides, and systematic model uncertainties need to be taken into account.<ref name=smits/>

As the closed shells are passed, the stabilizing effect should vanish.<ref name=relqed/> Thus, superheavy nuclides with more than 184 neutrons are expected to have much shorter lifetimes, spontaneously fissioning within 10<sup>−15</sup>&nbsp;seconds. If this is so, then it would not make sense to consider them chemical elements: IUPAC defines an element to exist only if the nucleus lives longer than 10<sup>−14</sup>&nbsp;seconds, the time needed for it to gather an electron cloud. Nonetheless, theoretical estimates of half-lives are very model-dependent, ranging over many orders of magnitude.<ref name=gamowstates/> The extreme repulsion between protons is predicted to result in exotic nuclear topologies, with bubbles, rings, and tori expected: this further complicates extrapolation.<ref name=smits/> It is not clear if any further-out shell closures exist, due to an expected smearing out of distinct nuclear shells (as is already expected for the electron shells at oganesson).<ref name=relqed>{{cite journal |last1=Schwerdtfeger |first1=Peter |last2=Pašteka |first2=Lukáš F. |last3=Punnett |first3=Andrew |last4=Bowman |first4=Patrick O. |date=2015 |title=Relativistic and quantum electrodynamic effects in superheavy elements |journal=Nuclear Physics A |volume=944 |issue=December 2015 |pages=551–577 |doi=10.1016/j.nuclphysa.2015.02.005|bibcode=2015NuPhA.944..551S }}</ref> Furthermore, even if later shell closures exist, it is not clear if they would allow such heavy elements to exist.<ref name="greinernuclei">{{cite journal|last=Greiner|first=W.|date=2013|title=Nuclei: superheavy-superneutronic-strange-and of antimatter|url=https://inspirehep.net/record/1221632/files/jpconf13_413_012002.pdf|journal=Journal of Physics: Conference Series|volume=413|issue=1|pages=012002-1–012002-9<!-- Deny Citation Bot-->|doi=10.1088/1742-6596/413/1/012002|bibcode=2013JPhCS.413a2002G|doi-access=free|access-date=15 August 2021|archive-date=30 March 2019|archive-url=https://web.archive.org/web/20190330183222/https://inspirehep.net/record/1221632/files/jpconf13_413_012002.pdf|url-status=live}}</ref><ref name="radiochimica">{{cite journal |last1=Hofmann |first1=Sigurd |date=2019 |title=Synthesis and properties of isotopes of the transactinides |journal=Radiochimica Acta |volume=107 |issue=9–11 |pages=879–915 |doi=10.1515/ract-2019-3104|s2cid=203848120 }}</ref><ref name="PTSS1">Scerri, p. 386</ref><ref name="EB">{{cite encyclopedia|last1=Seaborg|first1=G.|url=https://www.britannica.com/EBchecked/topic/603220/transuranium-element|title=transuranium element (chemical element)|encyclopedia=Encyclopædia Britannica|date=c. 2006|access-date=16 March 2010|url-status=live|archive-url=https://web.archive.org/web/20101130112151/https://www.britannica.com/EBchecked/topic/603220/transuranium-element|archive-date=30 November 2010}}</ref> As such, it may be that the periodic table practically ends around element 120, as elements become too short-lived to observe, and then too short-lived to have chemistry; the era of discovering new elements would thus be close to its end.<ref name="EB"/><ref>{{cite journal |author=Peter Möller |url=https://www.epj-conferences.org/articles/epjconf/abs/2016/26/epjconf-NS160-03002/epjconf-NS160-03002.html |title=The limits of the nuclear chart set by fission and alpha decay {{pipe}} EPJ Web of Conferences |journal=European Physical Journal Web of Conferences |year=2016 |doi=10.1051/epjconf/201613103002 |publisher=Epj-conferences.org |volume=131 |page=03002 |bibcode=2016EPJWC.13103002M |access-date=13 June 2022 |archive-date=20 June 2022 |archive-url=https://web.archive.org/web/20220620210806/https://www.epj-conferences.org/articles/epjconf/abs/2016/26/epjconf-NS160-03002/epjconf-NS160-03002.html |url-status=live |doi-access=free }}</ref> If another proton shell closure beyond 126 does exist, then it probably occurs around 164;<ref name=greinernuclei/> thus the region where periodicity fails more or less matches the region of instability between the shell closures.<ref name=actrev/>

Alternatively, ] may become stable at high mass numbers, in which the nucleus is composed of freely flowing ] and ]s instead of binding them into protons and neutrons; this would create a ] instead of an island.<ref name="udQM">{{cite journal |last1=Holdom |first1=B. |last2=Ren |first2=J. |last3=Zhang |first3=C. |title=Quark matter may not be strange |date=2018 |journal=Physical Review Letters |volume=120 |issue=1 |pages=222001-1–222001-6 <!-- Deny Citation Bot-->|doi=10.1103/PhysRevLett.120.222001|pmid=29906186 |arxiv=1707.06610 |bibcode=2018PhRvL.120v2001H |s2cid=49216916 }}</ref><ref name="udQMnew">{{cite journal |last1=Cheng-Jun |first1=Xia |last2=She-Sheng |first2=Xue |last3=Ren-Xin |first3=Xu |last4=Shan-Gui |first4=Zhou |title=Supercritically charged objects and electron-positron pair creation |doi=10.1103/PhysRevD.101.103031 |journal=Physical Review D |year=2020 |volume=101 |issue=10 |page=103031|arxiv=2001.03531 |bibcode=2020PhRvD.101j3031X |s2cid=210157134 }}</ref> Other effects may come into play: for example, in very heavy elements the 1s electrons are likely to spend a significant amount of time so close to the nucleus that they are actually inside it, which would make them vulnerable to ].<ref name=colloq>{{cite journal |title=Colloquium: Superheavy elements: Oganesson and beyond |first1=S. A. |last1=Giuliani |first2=Z. |last2=Matheson |first3=W. |last3=Nazarewicz |first4=E. |last4=Olsen |first5=P.-G. |last5=Reinhard |first6=J. |last6=Sadhukhan |first7=B. |last7=Schtruempf |first8=N. |last8=Schunck |first9=P. |last9=Schwerdtfeger |date=2019 |journal=Reviews of Modern Physics |volume=91 |issue=1 |pages=011001-1–011001-25 |doi=10.1103/RevModPhys.91.011001|bibcode=2019RvMP...91a1001G |s2cid=126906074 |doi-access=free }}</ref>

Even if eighth-row elements can exist, producing them is likely to be difficult, and it should become even more difficult as atomic number rises.<ref>{{cite journal|last1=Giardina|first1=G.|last2=Fazio|first2=G.|last3=Mandaglio|first3=G.|last4=Manganaro|first4=M.|last5=Nasirov|first5=A.K.|last6=Romaniuk|first6=M.V.|last7=Saccà|first7=C.|title=Expectations and limits to synthesize nuclei with Z ≥ 120|date=2010|journal=International Journal of Modern Physics E|volume=19|issue=5 & 6|pages=882–893|doi=10.1142/S0218301310015333|url=https://www.researchgate.net/publication/263915732|bibcode=2010IJMPE..19..882G|access-date=15 August 2021|archive-date=19 October 2021|archive-url=https://web.archive.org/web/20211019202251/https://www.researchgate.net/publication/263915732_EXPECTATIONS_AND_LIMITS_TO_SYNTHESIZE_NUCLEI_WITH_Z_120|url-status=live}}</ref> Although the 8s elements 119 and 120 are expected to be reachable with present means, the elements beyond that are expected to require new technology,<ref name="Zagrebaev">{{cite journal|title=Future of superheavy element research: Which nuclei could be synthesized within the next few years?|url=https://nrv.jinr.ru/pdf_file/J_phys_2013.pdf|first1=Valeriy|last1=Zagrebaev|first2=Alexander|last2=Karpov|first3=Walter|last3=Greiner|date=2013|journal=Journal of Physics|volume=420|issue=1|page=012001|publisher=IOP Publishing Ltd.|doi=10.1088/1742-6596/420/1/012001|arxiv=1207.5700|bibcode=2013JPhCS.420a2001Z|s2cid=55434734|access-date=1 December 2020|archive-date=3 October 2015|archive-url=https://web.archive.org/web/20151003154020/https://nrv.jinr.ru/pdf_file/J_phys_2013.pdf|url-status=live}}</ref> if they can be produced at all.<ref name="Bloomberg">{{cite web|last=Subramanian|first=S.|author-link=Samanth Subramanian|date=2019|url=https://www.bloomberg.com/news/features/2019-08-28/making-new-elements-doesn-t-pay-just-ask-this-berkeley-scientist|title=Making New Elements Doesn't Pay. Just Ask This Berkeley Scientist|website=]|access-date=18 January 2020|archive-date=11 December 2019|archive-url=https://web.archive.org/web/20191211191525/https://www.bloomberg.com/news/features/2019-08-28/making-new-elements-doesn-t-pay-just-ask-this-berkeley-scientist|url-status=live}}</ref> Experimentally characterizing these elements chemically would also pose a great challenge.<ref name="nature2019" />

== Alternative periodic tables ==
{{Main|Types of periodic tables}}
]'s spiral periodic table (1964)]]
The periodic law may be represented in multiple ways, of which the standard periodic table is only one.<ref>Scerri, p. 20</ref> Within 100 years of the appearance of Mendeleev's table in 1869, ] had collected an estimated 700 different published versions of the periodic table.<ref name="Jensen">{{cite journal|last1=Jensen|first1=William B.|title=Classification, symmetry and the periodic table|journal=Comp. & Maths. With Appls.|date=1986|volume=12B|issue=I/2|url=https://www.che.uc.edu/Jensen/W.%20B.%20Jensen/Reprints/028.%20Periodic%20Table.pdf|access-date=18 January 2017|archive-url=https://web.archive.org/web/20170131184706/https://www.che.uc.edu/Jensen/W.%20B.%20Jensen/Reprints/028.%20Periodic%20Table.pdf|archive-date=31 January 2017|url-status=live}}</ref><ref name="Papers">{{Cite book |url=https://othmerlib.sciencehistory.org/record=b1069103~S6 |title=Edward G. Mazurs Collection of Periodic Systems Images |publisher=] |type=Finding Aid |access-date=2 October 2018 |archive-url=https://web.archive.org/web/20190327082328/https://othmerlib.sciencehistory.org/record%3Db1069103~S6 |archive-date=27 March 2019 |url-status=live}} Click on 'Finding Aid' to go to full finding aid.</ref> Many forms retain the rectangular structure, including ]'s left-step periodic table (pictured below), and the modernised form of Mendeleev's original 8-column layout that is still common in Russia. Other periodic table formats have been shaped much more exotically, such as spirals (]'s pictured to the right), circles and triangles.<ref>{{cite journal |last=Francl |first=M. |title=Table manners |journal=Nature Chemistry |volume=1 |date=May 2009 |pages=97–98 |url=https://ericscerri.com/Michelle-Nat%20Chem.pdf |bibcode=2009NatCh...1...97F |doi=10.1038/nchem.183 |issue=2 |pmid=21378810 |url-status=live |archive-url=https://web.archive.org/web/20121025135145/https://ericscerri.com/Michelle-Nat%20Chem.pdf |archive-date=25 October 2012}}</ref>

Alternative periodic tables are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables, with different ones skewed more towards emphasizing chemistry or physics at either end.<ref name="Scerri402">Scerri, pp. 402–3</ref> The standard form, which remains by far the most common, is somewhere in the middle.<ref name="Scerri402" />

The many different forms of the periodic table have prompted the questions of whether there is an optimal or definitive form of the periodic table, and if so, what it might be. There are no current consensus answers to either question.<ref name="sesqui">{{cite web
|url=https://blog.oup.com/2019/01/happy-sesquicentennial-periodic-table-elements/
|title=Happy sesquicentennial to the periodic table of the elements
|last=Scerri
|first=Eric
|date=29 January 2019
|publisher=Oxford University Press
|access-date=12 April 2019
|archive-url=https://web.archive.org/web/20190327082337/https://blog.oup.com/2019/01/happy-sesquicentennial-periodic-table-elements/
|archive-date=27 March 2019
|url-status=live
}}</ref><ref name="Scerri402" /> Janet's left-step table is being increasingly discussed as a candidate for being the optimal or most fundamental form; Scerri has written in support of it, as it clarifies helium's nature as an s-block element, increases regularity by having all period lengths repeated, faithfully follows Madelung's rule by making each period correspond to one value of {{mvar|n}} + {{math|ℓ}},{{efn|name=lowdin|Authors differ on whether the {{mvar|n}} + {{math|ℓ}} rule has yet been derived from quantum mechanics. Scerri claims that it has not,<ref>Scerri, p. 255</ref><ref>{{cite book |last=Scerri |first=ER |date=2021 |editor-last1=Giunta |editor-first1=CJ |editor-last2=Mainz |editor-first2=VV |editor-last3=Girolami |editor-first3=GS |title=150 Years of the Periodic Table: Perspectives on the History of Chemistry|publisher=Book Publishers |pages=409–423(414) |chapter=The Impact of Twentieth-Century Physics on the Periodic Table and Some Remaining Questions in the Twenty-First Century |doi=10.1007/978-3-030-67910-1_16|isbn=978-3-030-67909-5 }}</ref> despite several attempts to do so.<ref>
{{cite journal |last1=Scerri |first1=ER |date=2009 |title=The dual sense of the term "element", attempts to derive the Madelung rule and the optimal form of the periodic table, if any|url= |journal=Int J Quantum Chem |volume=109 |issue= 5|pages=959–971 |doi=10.1002/qua.21914 |bibcode=2009IJQC..109..959S |access-date=}};
{{cite journal |last1=Bent |first1=HA |last2= Weinhold|first2=F |date=2007 |title=News from the periodic table: an introduction to periodicity symbols, tables and models for higher order valency and donor-acceptor kinships|url= |journal=J Chem Educ |volume=84 |issue= |pages=1145–1146 |doi=10.1021/ed084p1145 |access-date=}};
{{cite journal |last1=Allen |first1=LC |last2=Knight |first2=ET |date=2002 |title=The Löwdin challenge: origin of the (Madelung) rule for filling the orbital configurations of the periodic table|url= |journal=J Quantum Chem |volume=90 |issue= |pages=80–82 |doi= 10.1002/qua.965|access-date=}};
{{cite journal |last1=Wong |first1=DP |date=1979 |title=Theoretical justification of Madelung's rule|url= |journal= J Chem Educ |volume=56 |issue= 11|pages=714–717 |doi=10.1021/ed056p714 |bibcode=1979JChEd..56..714W |access-date=}};
{{cite journal |last1=Demkov |first1=YN |last2=Ostrovsky |first2=V |date= 1972|title=n + ' filling rule in the periodic system and focusing potentials.|url= http://www.jetp.ras.ru/cgi-bin/dn/e_035_01_0066.pdf|journal=Soviet Physics JETP |volume=35 |issue= |pages=66–69 |doi= |access-date=8 February 2024}}</ref> On the other hand, Ostrovsky, who has claimed such justification from 1971, wrote 'Some authors insist that "still nobody has deduced the n+l rule from the principles of quantum
mechanics", while others present quantum justification of the rule that was not ever disputed.'<ref>{{cite journal |last1=Ostrovsky |first1=V. N. |date=2005 |title=On Recent Discussion Concerning Quantum Justification of the Periodic Table of the Elements |url= |journal=Foundations of Chemistry |volume=7 |issue=3 |pages=235–239 |doi=10.1007/s10698-005-2141-y |s2cid=93589189 |access-date=}}</ref> Other authors argue that such a derivation is not necessary, because it admits exceptions.<ref>{{cite journal |last1=Scerri |first1=Eric |date=2012 |title=What is an element? What is the periodic table? And what does quantum mechanics contribute to the question? |url=https://philpapers.org/archive/SCEWIA.pdf |journal=Foundations of Chemistry |volume=14 |issue= |pages=69–81 |doi=10.1007/s10698-011-9124-y |s2cid=254503469}}</ref>}} and regularises atomic number triads and the first-row anomaly trend. While he notes that its placement of helium atop the alkaline earth metals can be seen a disadvantage from a chemical perspective, he counters this by appealing to the first-row anomaly, pointing out that the periodic table "fundamentally reduces to quantum mechanics", and that it is concerned with "abstract elements" and hence atomic properties rather than macroscopic properties.<ref>{{cite journal |last1=Scerri |first1=Eric |date=2021 |title=Various forms of the periodic table including the left-step table, the regularization of atomic number triads and first-member anomalies |url= |journal=ChemTexts |volume=8 |issue=6 |pages= |doi=10.1007/s40828-021-00157-8 |s2cid=245540088 }}</ref>

{{Periodic table (left step)}}


== See also == == See also ==
* ]
{{Misplaced Pages books|Periodic table}}

* ]
== Notes ==
* ]
{{notelist}}
* ]
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* ]
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* ]


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


== Bibliography == == Bibliography ==
{{refbegin}} {{Refbegin}}
* {{Greenwood&Earnshaw2nd}}
* {{cite book |last=Ball |first=Philip |title=The Ingredients: A Guided Tour of the Elements |publisher=Oxford University Press |year=2002 |isbn=0-19-284100-9}}
* {{cite book |last1 = Petrucci |first1 = Ralph H. |last2 = Harwood |first2 = William S. |last3 = Herring |first3 = F. Geoffrey |date=2002 |title = General chemistry: principles and modern applications |url = https://archive.org/details/generalchemistry00hill |url-access = registration |edition=8th |location=Upper Saddle River, N.J. |publisher=Prentice Hall |isbn = 978-0-13-014329-7 |lccn=2001032331 |oclc=46872308 }}
* {{cite journal |author=Bouma, J. |title=An Application-Oriented Periodic Table of the Elements |journal=J. Chem. Ed. |volume=66 |page=741 |year=1989 |doi=10.1021/ed066p741 |bibcode = 1989JChEd..66..741B |issue=9}}
* {{cite book|last=Gray|first=Theodore|authorlink=Theodore Gray|title=The Elements: A Visual Exploration of Every Known Atom in the Universe|year=2009|publisher=Black Dog & Leventhal Publishers|location=New York|isbn=978-1-57912-814-2}} * {{cite book|title=Concise Chemistry of the Elements|year=2002|publisher=Horwood|isbn=978-1-898563-71-6|last1=Siekierski|first1=S.|last2=Burgess|first2 =J.}}
* ] (2020). ''The Periodic Table, Its Story and Its Significance'' (2nd ed.). Oxford University Press, New York, {{ISBN|978-0-19-091436-3}}.
* {{cite journal |author= Hjørland, Birger |year= 2011 |title= The periodic table and the philosophy of classification |journal= ] |volume= 38 |issue= 1 |pages= 9–21 |accessdate= 2011-03-13 |url= http://ucla.academia.edu/EricScerri/Papers/432740/Forum%5FThe%5FPhilosophy%5Fof%5FClassification }}
* {{cite book |author=Kean, Sam |title=The Disappearing Spoon - and other true tales from the Periodic Table|publisher=Black Swan |location=London |year=2010 |isbn=978-0-552-77750-6}} * {{cite book |last=Wulfsberg |first=Gary |author-link= |date=2000 |title=Inorganic Chemistry |url= |location= |publisher=University Science Books |page= |isbn=9781891389016}}
{{Refend}}
* {{cite book |author=], Primo |title=The Periodic Table |publisher=Penguin Books|location=London |year=1984|isbn=978-0-14-139944-7}}

* {{cite book |author=Mazurs, E.G |title=Graphical Representations of the Periodic System During One Hundred Years |publisher=University of Alabama Press |location=Alabama |year=1974}}
== Further reading ==
* {{cite book |last=Moore |first=John |title=Chemistry ] |year=2003 |publisher=Wiley Publications |location=New York |isbn=978-0-7645-5430-8 |oclc=51168057 |page=111}}
{{Refbegin}}
* {{cite book |author=Scerri, Eric |title=The periodic table: its story and its significance |publisher=Oxford University Press |location=Oxford |year=2007 |isbn=0-19-530573-6}}
* {{cite book |last=Calvo |first=Miguel |year=2019 |title= Construyendo la Tabla Periódica|page=407|location=Zaragoza, Spain |publisher=Prames |isbn=978-84-8321-908-9}}
{{refend}}
* {{cite book |last=Emsley |first=J. |year=2011 |title=Nature's Building Blocks: An A–Z Guide to the Elements |chapter=The Periodic Table|pages=634–651|location=Oxford|publisher=Oxford University Press |edition=New|isbn=978-0-19-960563-7 |author-link=John Emsley }}
* {{cite book |first1=Marco |last1=Fontani |first2=Mariagrazia |last2=Costa |first3=Mary Virginia |last3=Orna |year=2007 |title=The Lost Elements: The Periodic Table's Shadow Side |page=508|location=Oxford|publisher= Oxford University Press |isbn=978-0-19-938334-4}}
* {{cite book |last=Mazurs |first=E. G.|year=1974 |title=Graphic Representations of the Periodic System During One Hundred Years |location=Alabama |publisher=University of Alabama Press |isbn=978-0-19-960563-7 |author-link=John Emsley }}
* {{cite conference |editor1-first=D.H. |editor1-last=Rouvray |editor2-first=R. B. |editor2-last=King |title=The Periodic Table: Into the 21st Century |publisher= Research Studies Press |location=Baldock, Hertfordshire|conference=Proceedings of the 2nd International Conference on the Periodic Table, part 1, Kananaskis Guest Ranch, Alberta, 14–20 July 2003 |isbn=978-0-86380-292-8 |year=2004}}
* {{cite conference |editor1-first=D.H. |editor1-last=Rouvray |editor2-first=R. B. |editor2-last=King |title=The Mathematics of the Periodic Table |publisher= Nova Science |location=New York|conference=Proceedings of the 2nd International Conference on the Periodic Table, part 2, Kananaskis Guest Ranch, Alberta, 14–20 July 2003 |isbn=978-1-59454-259-6 |year=2006}}
* {{cite web |url=https://www.ericscerri.com/books_elements.pdf |title=Books on the Elements and the Periodic Table |last=Scerri |first=E |date=n.d. |access-date=9 July 2018 |archive-date=11 August 2020 |archive-url=https://web.archive.org/web/20200811052254/https://www.ericscerri.com/books_elements.pdf |url-status=live }}
* {{cite conference |editor1-first=E. |editor1-last=Scerri |editor2-first=G |editor2-last=Restrepo |title=Mendeleev to Oganesson: A Multidisciplinary Perspective on the Periodic Table |publisher= Oxford University Press |location=Oxford|conference=Proceedings of the 3rd International Conference on the Periodic Table, Cuzco, Peru 14–16 August 2012 |isbn=978-0-86380-292-8 |year=2018}}
* {{cite book|last=van Spronsen|first=J. W.|title=The Periodic System of Chemical Elements: A History of the First Hundred Years |location=Amsterdam|publisher=Elsevier |year=1969 |isbn=978-0-444-40776-4}}
* {{cite conference |editor-first=M. |editor-last=Verde |title= Atti del convegno Mendeleeviano: Periodicità e simmetrie nella struttura elementare della materia |publisher=Accademia delle Scienze di Torino |location=Torino |trans-title=Proceedings of the Mendeleevian conference: Periodicity and symmetry in the elementary structure of matter|conference=1st International Conference on the Periodic Table, Torino-Roma, 15–21 September 1969|year=1971}}
{{Refend}}


== External links == == External links ==
{{Sister project links|Periodic table}} {{Sister project links|Periodic table}}
* featured topic page on ] 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.
*
* , with interactive layouts
* , leading philosopher of science specializing in the history and philosophy of the periodic table
*
*
*
* {{Webarchive|url=https://web.archive.org/web/20230703042907/https://www.periodicvideos.com/ |date=3 July 2023 }}
*
* {{Webarchive|url=https://web.archive.org/web/20221230200920/https://www.periodicgraphicsofelements.com/ |date=30 December 2022 }}


{{Periodic table (navbox)}}
* {{cite web |url=http://www.ptable.com |title=Dynamic Periodic Table |author=M. Dayah|accessdate=14 May 2012}}
{{Navbox periodic table}}

* {{cite web |url=http://www.periodicvideos.com |title=The Periodic Table of Videos |author=Brady Haran |publisher=University of Nottingham |accessdate=14 May 2012}}

* {{cite web |url=http://www.webelements.com |title=WebElements: the periodic table on the web |author=Mark Winter |publisher=University of Sheffield |accessdate=14 May 2012}}

* {{cite web |url=http://www.meta-synthesis.com/webbook//35_pt/pt_database.php |title=The INTERNET Database of Periodic Tables |author=Mark R. Leach |accessdate=14 May 2012}}

* {{cite web |url=http://www.jergym.hiedu.cz/~canovm/vyhledav/chemici2.html |title=Periodic Table of the Elements in Four Hundred Languages |accessdate=14 May 2012}}
{{Compact periodic table}}
{{PeriodicTablesFooter}}
{{BranchesofChemistry}} {{BranchesofChemistry}}
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Latest revision as of 18:30, 24 December 2024

Tabular arrangement of the chemical elements ordered by atomic number This article is about the table used in chemistry and physics. For other uses, see Periodic table (disambiguation).

Periodic table of the chemical elements showing the most or more commonly named sets of elements (in periodic tables), and a traditional dividing line between metals and nonmetals. The f-block actually fits between groups 2 and 3; it is usually shown at the foot of the table to save horizontal space.
Part of a series on the
Periodic table
Periodic table forms
Periodic table history
Sets of elements
By periodic table structure
By metallic classification
By other characteristics
Elements
List of chemical elements
Properties of elements
Data pages for elements

The periodic table, also known as the periodic table of the elements, is an ordered arrangement of the chemical elements into rows ("periods") and columns ("groups"). It is an icon of chemistry and is widely used in physics and other sciences. It is a depiction of the periodic law, which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks. Elements in the same group tend to show similar chemical characteristics.

Vertical, horizontal and diagonal trends characterize the periodic table. Metallic character increases going down a group and from right to left across a period. Nonmetallic character increases going from the bottom left of the periodic table to the top right.

The first periodic table to become generally accepted was that of the Russian chemist Dmitri Mendeleev in 1869; he formulated the periodic law as a dependence of chemical properties on atomic mass. As not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to predict some properties of some of the missing elements. The periodic law was recognized as a fundamental discovery in the late 19th century. It was explained early in the 20th century, with the discovery of atomic numbers and associated pioneering work in quantum mechanics, both ideas serving to illuminate the internal structure of the atom. A recognisably modern form of the table was reached in 1945 with Glenn T. Seaborg's discovery that the actinides were in fact f-block rather than d-block elements. The periodic table and law are now a central and indispensable part of modern chemistry.

The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 exist; to go further, it was necessary to synthesize new elements in the laboratory. By 2010, the first 118 elements were known, thereby completing the first seven rows of the table; however, chemical characterization is still needed for the heaviest elements to confirm that their properties match their positions. New discoveries will extend the table beyond these seven rows, though it is not yet known how many more elements are possible; moreover, theoretical calculations suggest that this unknown region will not follow the patterns of the known part of the table. Some scientific discussion also continues regarding whether some elements are correctly positioned in today's table. Many alternative representations of the periodic law exist, and there is some discussion as to whether there is an optimal form of the periodic table.

Structure

Periodic table
Group 1 2   3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Hydrogen &
alkali metals 		Alkaline earth metals Triels Tetrels Pnicto­gens Chal­co­gens Halo­gens Noble
gases
Period

1

Hydro­gen1H​1.0080 He­lium2He​4.0026
2 Lith­ium3Li​6.94 Beryl­lium4Be​9.0122 Boron5B​10.81 Carbon6C​12.011 Nitro­gen7N​14.007 Oxy­gen8O​15.999 Fluor­ine9F​18.998 Neon10Ne​20.180
3 So­dium11Na​22.990 Magne­sium12Mg​24.305 Alumin­ium13Al​26.982 Sili­con14Si​28.085 Phos­phorus15P​30.974 Sulfur16S​32.06 Chlor­ine17Cl​35.45 Argon18Ar​39.95
4 Potas­sium19K​39.098 Cal­cium20Ca​40.078 Scan­dium21Sc​44.956 Tita­nium22Ti​47.867 Vana­dium23V​50.942 Chrom­ium24Cr​51.996 Manga­nese25Mn​54.938 Iron26Fe​55.845 Cobalt27Co​58.933 Nickel28Ni​58.693 Copper29Cu​63.546 Zinc30Zn​65.38 Gallium31Ga​69.723 Germa­nium32Ge​72.630 Arsenic33As​74.922 Sele­nium34Se​78.971 Bromine35Br​79.904 Kryp­ton36Kr​83.798
5 Rubid­ium37Rb​85.468 Stront­ium38Sr​87.62 Yttrium39Y​88.906 Zirco­nium40Zr​91.224 Nio­bium41Nb​92.906 Molyb­denum42Mo​95.95 Tech­netium43Tc​ Ruthe­nium44Ru​101.07 Rho­dium45Rh​102.91 Pallad­ium46Pd​106.42 Silver47Ag​107.87 Cad­mium48Cd​112.41 Indium49In​114.82 Tin50Sn​118.71 Anti­mony51Sb​121.76 Tellur­ium52Te​127.60 Iodine53I​126.90 Xenon54Xe​131.29
6 Cae­sium55Cs​132.91 Ba­rium56Ba​137.33 1 asterisk Lute­tium71Lu​174.97 Haf­nium72Hf​178.49 Tanta­lum73Ta​180.95 Tung­sten74W​183.84 Rhe­nium75Re​186.21 Os­mium76Os​190.23 Iridium77Ir​192.22 Plat­inum78Pt​195.08 Gold79Au​196.97 Mer­cury80Hg​200.59 Thallium81Tl​204.38 Lead82Pb​207.2 Bis­muth83Bi​208.98 Polo­nium84Po​ Asta­tine85At​ Radon86Rn​
7 Fran­cium87Fr​ Ra­dium88Ra​ 1 asterisk Lawren­cium103Lr​ Ruther­fordium104Rf​ Dub­nium105Db​ Sea­borgium106Sg​ Bohr­ium107Bh​ Has­sium108Hs​ Meit­nerium109Mt​ Darm­stadtium110Ds​ Roent­genium111Rg​ Coper­nicium112Cn​ Nihon­ium113Nh​ Flerov­ium114Fl​ Moscov­ium115Mc​ Liver­morium116Lv​ Tenness­ine117Ts​ Oga­nesson118Og​
1 asterisk Lan­thanum57La​138.91 Cerium58Ce​140.12 Praseo­dymium59Pr​140.91 Neo­dymium60Nd​144.24 Prome­thium61Pm​ Sama­rium62Sm​150.36 Europ­ium63Eu​151.96 Gadolin­ium64Gd​157.25 Ter­bium65Tb​158.93 Dyspro­sium66Dy​162.50 Hol­mium67Ho​164.93 Erbium68Er​167.26 Thulium69Tm​168.93 Ytter­bium70Yb​173.05  
1 asterisk Actin­ium89Ac​ Thor­ium90Th​232.04 Protac­tinium91Pa​231.04 Ura­nium92U​238.03 Neptu­nium93Np​ Pluto­nium94Pu​ Ameri­cium95Am​ Curium96Cm​ Berkel­ium97Bk​ Califor­nium98Cf​ Einstei­nium99Es​ Fer­mium100Fm​ Mende­levium101Md​ Nobel­ium102No​
Primordial  From decay  Synthetic Border shows natural occurrence of the element Standard atomic weight Ar, std(E)
  • Ca: 40.078 — Abridged value (uncertainty omitted here)
  • Po:  — mass number of the most stable isotope
s-block f-block d-block p-block
3D views of some hydrogen-like atomic orbitals showing probability density and phase (g orbitals and higher are not shown)

Each chemical element has a unique atomic number (Z— for "Zahl", German for "number") representing the number of protons in its nucleus. Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements. The chemical elements are what the periodic table classifies and organizes. Hydrogen is the element with atomic number 1; helium, atomic number 2; lithium, atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li. Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element. Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined atomic weights, defined as the average mass of a naturally occurring atom of that element. All elements have multiple isotopes, variants with the same number of protons but different numbers of neutrons. 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. When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses.

In the standard periodic table, the elements are listed in order of increasing atomic number. 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, sulfur, and selenium are in the same column because they all 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.

Today, 118 elements are known, the first 94 of which are known to occur naturally on Earth at present. 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. A few of the latter are so rare that they were not discovered in nature, but were synthesized in the laboratory before it was determined that they do exist in nature after all: technetium (element 43), promethium (element 61), astatine (element 85), neptunium (element 93), and plutonium (element 94). No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form, nor has astatine; francium (element 87) has been only photographed in the form of light emitted from microscopic quantities (300,000 atoms). Of the 94 natural elements, eighty have a stable isotope and one more (bismuth) has an almost-stable isotope (with a half-life of 2.01×10 years, over a billion times the age of the universe). Two more, thorium and uranium, have isotopes undergoing radioactive decay with a half-life comparable to the age of the Earth. The stable elements plus bismuth, thorium, and uranium make up the 83 primordial elements that survived from the Earth's formation. The remaining eleven natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of thorium and uranium. All 24 known artificial elements are radioactive.

Group names and numbers

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). The f-block groups are ignored in this numbering. Groups can also be named by their first element, e.g. the "scandium group" for group 3. Previously, groups were known by Roman numerals. In the United States, 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 for groups 1 through 7, and "B" was used for groups 11 through 17. 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 (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated.

Groups in the periodic table
IUPAC group 1 2 — 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Mendeleev (I–VIII) IA IIA IIIB IVB VB VIB VIIB VIIIB IB IIB IIIB IVB VB VIB VIIB
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 VIIIB IB IIB IIIB IVB VB VIB VIIB 0
Trivial name H and alkali metals alkaline earth metals triels tetrels pnicto­gens chal­co­gens halo­gens noble gases
Name by element 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
Group 1 is composed of hydrogen (H) and the alkali metals. Elements of the group have one s-electron in the outer electron shell. Hydrogen is not considered to be an alkali metal as it is not a metal, though it is more analogous to them than any other group. This makes the group somewhat exceptional.
The 14 f-block groups (columns) do not have a group number.
The correct composition of group 3 is scandium (Sc), yttrium (Y), lutetium (Lu), and lawrencium (Lr), as shown here: this is endorsed by 1988 and 2021 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, and Lev Landau and Evgeny Lifshitz already considered it incorrect in 1948. Arguments can still occasionally be encountered in the contemporary literature purporting to defend it, but most authors consider them logically inconsistent. 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. See also Group 3 element#Composition.
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.
Group name as recommended by IUPAC.

Presentation forms

Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson

32 columns

Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
​
Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium
Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium

18 columns

For reasons of space, the periodic table is commonly presented with the f-block elements cut out and positioned as a distinct part below the main body. This reduces the number of element columns from 32 to 18.

Both forms represent the same periodic table. The form with the f-block included in the main body is sometimes called the 32-column or long form; the form with the f-block cut out the 18-column or medium-long form. The 32-column form has the advantage of showing all elements in their correct sequence, but it has the disadvantage of requiring more space. The form chosen is an editorial choice, and does not imply any change of scientific claim or statement. For example, when discussing the composition of group 3, the options can be shown equally (unprejudiced) in both forms.

Periodic tables usually at least show the elements' symbols; many also provide supplementary information about the elements, either via colour-coding or as data in the cells. The above table shows the names and atomic numbers of the elements, and also their blocks, natural occurrences and standard atomic weights. For the short-lived elements without standard atomic weights, the mass number of the most stable known isotope is used instead. Other tables may include properties such as state of matter, melting and boiling points, densities, as well as provide different classifications of the elements.

Electron configurations

Main article: Electron configuration

The periodic table is a graphic description of the periodic law, which states that the properties and atomic structures of the chemical elements are a periodic function of their atomic number. Elements are placed in the periodic table according to their electron configurations, the periodic recurrences of which explain the trends in properties across the periodic table.

An electron can be thought of as inhabiting an atomic orbital, which characterizes the probability it can be found in any particular region around the atom. Their energies are quantised, which is to say that they can only take discrete values. Furthermore, electrons obey the Pauli exclusion principle: different electrons must always be in different states. This allows classification of the possible states an electron can take in various energy levels known as shells, divided into individual subshells, which each contain one or more orbitals. Each orbital can contain up to two electrons: they are distinguished by a quantity known as spin, conventionally labelled "up" or "down". In a cold atom (one in its ground state), electrons arrange themselves in such a way that the total energy they have is minimized by occupying the lowest-energy orbitals available. Only the outermost electrons (so-called valence electrons) have enough energy to break free of the nucleus and participate in chemical reactions with other atoms. The others are called core electrons.

ℓ = 0 1 2 3 4 5 6 Shell capacity (2n)
Orbital s p d f g h i
n = 1 1s 2
n = 2 2s 2p 8
n = 3 3s 3p 3d 18
n = 4 4s 4p 4d 4f 32
n = 5 5s 5p 5d 5f 5g 50
n = 6 6s 6p 6d 6f 6g 6h 72
n = 7 7s 7p 7d 7f 7g 7h 7i 98
Subshell capacity (4ℓ+2) 2 6 10 14 18 22 26

Elements are known with up to the first seven shells occupied. The first shell contains only one orbital, a spherical s orbital. As it is in the first shell, this is called the 1s orbital. This can hold up to two electrons. The second shell similarly contains a 2s orbital, and it also contains three dumbbell-shaped 2p orbitals, and can thus fill up to eight electrons (2×1 + 2×3 = 8). The third shell contains one 3s orbital, three 3p orbitals, and five 3d orbitals, and thus has a capacity of 2×1 + 2×3 + 2×5 = 18. The fourth shell contains one 4s orbital, three 4p orbitals, five 4d orbitals, and seven 4f orbitals, thus leading to a capacity of 2×1 + 2×3 + 2×5 + 2×7 = 32. Higher shells contain more types of orbitals that continue the pattern, but such types of orbitals are not filled in the ground states of known elements. The subshell types are characterized by the quantum numbers. Four numbers describe an orbital in an atom completely: the principal quantum number n, the azimuthal quantum number ℓ (the orbital type), the orbital magnetic quantum number mℓ, and the spin magnetic quantum number ms.

Order of subshell filling

Idealized order of subshell filling according to the Madelung rule

The sequence in which the subshells are filled is given in most cases by the Aufbau principle, also known as the Madelung or Klechkovsky rule (after Erwin Madelung and Vsevolod Klechkovsky respectively). This rule was first observed empirically by Madelung, and Klechkovsky and later authors gave it theoretical justification. The shells overlap in energies, and the Madelung rule specifies the sequence of filling according to:

1s ≪ 2s < 2p ≪ 3s < 3p ≪ 4s < 3d < 4p ≪ 5s < 4d < 5p ≪ 6s < 4f < 5d < 6p ≪ 7s < 5f < 6d < 7p ≪ ...

Here the sign ≪ means "much less than" as opposed to < meaning just "less than". Phrased differently, electrons enter orbitals in order of increasing n + ℓ, and if two orbitals are available with the same value of n + ℓ, the one with lower n is occupied first. In general, orbitals with the same value of n + ℓ are similar in energy, but in the case of the s-orbitals (with ℓ = 0), quantum effects raise their energy to approach that of the next n + ℓ group. Hence the periodic table is usually drawn to begin each row (often called a period) with the filling of a new s-orbital, which corresponds to the beginning of a new shell. Thus, with the exception of the first row, each period length appears twice:

2, 8, 8, 18, 18, 32, 32, ...

The overlaps get quite close at the point where the d-orbitals enter the picture, and the order can shift slightly with atomic number and atomic charge.

Starting from the simplest atom, this lets us build up the periodic table one at a time in order of atomic number, by considering the cases of single atoms. In hydrogen, there is only one electron, which must go in the lowest-energy orbital 1s. This electron configuration is written 1s, where the superscript indicates the number of electrons in the subshell. Helium adds a second electron, which also goes into 1s, completely filling the first shell and giving the configuration 1s.

Starting from the third element, lithium, the first shell is full, so its third electron occupies a 2s orbital, giving a 1s 2s configuration. The 2s electron is lithium's only valence electron, as the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms: such a shell is called a "core shell". The 1s subshell is a core shell for all elements from lithium onward. The 2s subshell is completed by the next element beryllium (1s 2s). The following elements then proceed to fill the 2p subshell. Boron (1s 2s 2p) puts its new electron in a 2p orbital; carbon (1s 2s 2p) fills a second 2p orbital; and with nitrogen (1s 2s 2p) all three 2p orbitals become singly occupied. This is consistent with Hund's rule, which states that atoms usually prefer to singly occupy each orbital of the same type before filling them with the second electron. Oxygen (1s 2s 2p), fluorine (1s 2s 2p), and neon (1s 2s 2p) then complete the already singly filled 2p orbitals; the last of these fills the second shell completely.

Starting from element 11, sodium, the second shell is full, making the second shell a core shell for this and all heavier elements. The eleventh electron begins the filling of the third shell by occupying a 3s orbital, giving a configuration of 1s 2s 2p 3s for sodium. This configuration is abbreviated 3s, where represents neon's configuration. Magnesium ( 3s) finishes this 3s orbital, and the following six elements aluminium, silicon, phosphorus, sulfur, chlorine, and argon fill the three 3p orbitals ( 3s 3p through 3s 3p). This creates an analogous series in which the outer shell structures of sodium through argon are analogous to those of lithium through neon, and is the basis for the periodicity of chemical properties that the periodic table illustrates: at regular but changing intervals of atomic numbers, the properties of the chemical elements approximately repeat.

The first 18 elements can thus be arranged as the start of a periodic table. Elements in the same column have the same number of valence electrons and have analogous valence electron configurations: these columns are called groups. The single exception is helium, which has two valence electrons like beryllium and magnesium, but is typically placed in the column of neon and argon to emphasise that its outer shell is full. (Some contemporary authors question even this single exception, preferring to consistently follow the valence configurations and place helium over beryllium.) There are eight columns in this periodic table fragment, corresponding to at most eight outer-shell electrons. A period begins when a new shell starts filling. Finally, the colouring illustrates the blocks: the elements in the s-block (coloured red) are filling s-orbitals, while those in the p-block (coloured yellow) are filling p-orbitals.

1
H 		2
He 		2×1 = 2 elements
1s 0p
3
Li 		4
Be 		5
B 		6
C 		7
N 		8
O 		9
F 		10
Ne 		2×(1+3) = 8 elements
2s 2p
11
Na 		12
Mg 		13
Al 		14
Si 		15
P 		16
S 		17
Cl 		18
Ar 		2×(1+3) = 8 elements
3s 3p

Starting the next row, for potassium and calcium the 4s subshell is the lowest in energy, and therefore they fill it. Potassium adds one electron to the 4s shell ( 4s), and calcium then completes it ( 4s). However, starting from scandium ( 3d 4s) the 3d subshell becomes the next highest in energy. The 4s and 3d subshells have approximately the same energy and they compete for filling the electrons, and so the occupation is not quite consistently filling the 3d orbitals one at a time. The precise energy ordering of 3d and 4s changes along the row, and also changes depending on how many electrons are removed from the atom. For example, due to the repulsion between the 3d electrons and the 4s ones, at chromium the 4s energy level becomes slightly higher than 3d, and so it becomes more profitable for a chromium atom to have a 3d 4s configuration than an 3d 4s one. A similar anomaly occurs at copper, whose atom has a 3d 4s configuration rather than the expected 3d 4s. These are violations of the Madelung rule. Such anomalies, however, do not have any chemical significance: most chemistry is not about isolated gaseous atoms, and the various configurations are so close in energy to each other that the presence of a nearby atom can shift the balance. Therefore, the periodic table ignores them and considers only idealized configurations.

At zinc ( 3d 4s), the 3d orbitals are completely filled with a total of ten electrons. Next come the 4p orbitals, completing the row, which are filled progressively by gallium ( 3d 4s 4p) through krypton ( 3d 4s 4p), in a manner analogous to the previous p-block elements. From gallium onwards, the 3d orbitals form part of the electronic core, and no longer participate in chemistry. The s- and p-block elements, which fill their outer shells, are called main-group elements; the d-block elements (coloured blue below), which fill an inner shell, are called transition elements (or transition metals, since they are all metals).

The next 18 elements fill the 5s orbitals (rubidium and strontium), then 4d (yttrium through cadmium, again with a few anomalies along the way), and then 5p (indium through xenon). Again, from indium onward the 4d orbitals are in the core. Hence the fifth row has the same structure as the fourth.

1
H 		2
He 		2×1 = 2 elements
1s 0d 0p
3
Li 		4
Be 		5
B 		6
C 		7
N 		8
O 		9
F 		10
Ne 		2×(1+3) = 8 elements
2s 0d 2p
11
Na 		12
Mg 		13
Al 		14
Si 		15
P 		16
S 		17
Cl 		18
Ar 		2×(1+3) = 8 elements
3s 0d 3p
19
K 		20
Ca 		21
Sc 		22
Ti 		23
V 		24
Cr 		25
Mn 		26
Fe 		27
Co 		28
Ni 		29
Cu 		30
Zn 		31
Ga 		32
Ge 		33
As 		34
Se 		35
Br 		36
Kr 		2×(1+3+5) = 18 elements
4s 3d 4p
37
Rb 		38
Sr 		39
Y 		40
Zr 		41
Nb 		42
Mo 		43
Tc 		44
Ru 		45
Rh 		46
Pd 		47
Ag 		48
Cd 		49
In 		50
Sn 		51
Sb 		52
Te 		53
I 		54
Xe 		2×(1+3+5) = 18 elements
5s 4d 5p

The sixth row of the table likewise starts with two s-block elements: caesium and barium. After this, the first f-block elements (coloured green below) begin to appear, starting with lanthanum. These are sometimes termed inner transition elements. As there are now not only 4f but also 5d and 6s subshells at similar energies, competition occurs once again with many irregular configurations; this resulted in some dispute about where exactly the f-block is supposed to begin, but most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle. Even though lanthanum does not itself fill the 4f subshell as a single atom, because of repulsion between electrons, its 4f orbitals are low enough in energy to participate in chemistry. At ytterbium, the seven 4f orbitals are completely filled with fourteen electrons; thereafter, a series of ten transition elements (lutetium through mercury) follows, and finally six main-group elements (thallium through radon) complete the period. From lutetium onwards the 4f orbitals are in the core, and from thallium onwards so are the 5d orbitals.

The seventh row is analogous to the sixth row: 7s fills (francium and radium), then 5f (actinium to nobelium), then 6d (lawrencium to copernicium), and finally 7p (nihonium to oganesson). Starting from lawrencium the 5f orbitals are in the core, and probably the 6d orbitals join the core starting from nihonium. Again there are a few anomalies along the way: for example, as single atoms neither actinium nor thorium actually fills the 5f subshell, and lawrencium does not fill the 6d shell, but all these subshells can still become filled in chemical environments. For a very long time, the seventh row was incomplete as most of its elements do not occur in nature. The missing elements beyond uranium started to be synthesized in the laboratory in 1940, when neptunium was made. (However, the first element to be discovered by synthesis rather than in nature was technetium in 1937.) The row was completed with the synthesis of tennessine in 2010 (the last element oganesson had already been made in 2002), and the last elements in this seventh row were given names in 2016.

1
H 		2
He 		2×1 = 2 elements
1s 0f 0d 0p
3
Li 		4
Be 		5
B 		6
C 		7
N 		8
O 		9
F 		10
Ne 		2×(1+3) = 8 elements
2s 0f 0d 2p
11
Na 		12
Mg 		13
Al 		14
Si 		15
P 		16
S 		17
Cl 		18
Ar 		2×(1+3) = 8 elements
3s 0f 0d 3p
19
K 		20
Ca 		21
Sc 		22
Ti 		23
V 		24
Cr 		25
Mn 		26
Fe 		27
Co 		28
Ni 		29
Cu 		30
Zn 		31
Ga 		32
Ge 		33
As 		34
Se 		35
Br 		36
Kr 		2×(1+3+5) = 18 elements
4s 0f 3d 4p
37
Rb 		38
Sr 		39
Y 		40
Zr 		41
Nb 		42
Mo 		43
Tc 		44
Ru 		45
Rh 		46
Pd 		47
Ag 		48
Cd 		49
In 		50
Sn 		51
Sb 		52
Te 		53
I 		54
Xe 		2×(1+3+5) = 18 elements
5s 0f 4d 5p
55
Cs 		56
Ba 		57
La 		58
Ce 		59
Pr 		60
Nd 		61
Pm 		62
Sm 		63
Eu 		64
Gd 		65
Tb 		66
Dy 		67
Ho 		68
Er 		69
Tm 		70
Yb 		71
Lu 		72
Hf 		73
Ta 		74
W 		75
Re 		76
Os 		77
Ir 		78
Pt 		79
Au 		80
Hg 		81
Tl 		82
Pb 		83
Bi 		84
Po 		85
At 		86
Rn 		2×(1+3+5+7) = 32 elements
6s 4f 5d 6p
87
Fr 		88
Ra 		89
Ac 		90
Th 		91
Pa 		92
U 		93
Np 		94
Pu 		95
Am 		96
Cm 		97
Bk 		98
Cf 		99
Es 		100
Fm 		101
Md 		102
No 		103
Lr 		104
Rf 		105
Db 		106
Sg 		107
Bh 		108
Hs 		109
Mt 		110
Ds 		111
Rg 		112
Cn 		113
Nh 		114
Fl 		115
Mc 		116
Lv 		117
Ts 		118
Og 		2×(1+3+5+7) = 32 elements
7s 5f 6d 7p

This completes the modern periodic table, with all seven rows completely filled to capacity.

Electron configuration table

The following table shows the electron configuration of a neutral gas-phase atom of each element. Different configurations can be favoured in different chemical environments. The main-group elements have entirely regular electron configurations; the transition and inner transition elements show twenty irregularities due to the aforementioned competition between subshells close in energy level. For the last ten elements (109–118), experimental data is lacking and therefore calculated configurations have been shown instead. Completely filled subshells have been greyed out.

Electron configurations of the chemical elements (neutral gaseous atoms in the ground state; predictions for elements 109–118)
Group: 1 2 ​ 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1s: 1
H
1 		2
He
2
[He]
2s:
2p: 		3
Li
1
- 		4
Be
2
- 		5
B
2
1 		6
C
2
2 		7
N
2
3 		8
O
2
4 		9
F
2
5 		10
Ne
2
6
[Ne]
3s:
3p: 		11
Na
1
- 		12
Mg
2
- 		13
Al
2
1 		14
Si
2
2 		15
P
2
3 		16
S
2
4 		17
Cl
2
5 		18
Ar
2
6
[Ar]
4s:
3d:
4p: 		19
K
1
-
- 		20
Ca
2
-
- 		21
Sc
2
1
- 		22
Ti
2
2
- 		23
V
2
3
- 		24
Cr
1
5
- 		25
Mn
2
5
- 		26
Fe
2
6
- 		27
Co
2
7
- 		28
Ni
2
8
- 		29
Cu
1
10
- 		30
Zn
2
10
- 		31
Ga
2
10
1 		32
Ge
2
10
2 		33
As
2
10
3 		34
Se
2
10
4 		35
Br
2
10
5 		36
Kr
2
10
6
[Kr]
5s:
4d:
5p: 		37
Rb
1
-
- 		38
Sr
2
-
- 		39
Y
2
1
- 		40
Zr
2
2
- 		41
Nb
1
4
- 		42
Mo
1
5
- 		43
Tc
2
5
- 		44
Ru
1
7
- 		45
Rh
1
8
- 		46
Pd
-
10
- 		47
Ag
1
10
- 		48
Cd
2
10
- 		49
In
2
10
1 		50
Sn
2
10
2 		51
Sb
2
10
3 		52
Te
2
10
4 		53
I
2
10
5 		54
Xe
2
10
6
[Xe]
6s:
4f:
5d:
6p: 		55
Cs
1
-
-
- 		56
Ba
2
-
-
- 		57
La
2
-
1
- 		58
Ce
2
1
1
- 		59
Pr
2
3
-
- 		60
Nd
2
4
-
- 		61
Pm
2
5
-
- 		62
Sm
2
6
-
- 		63
Eu
2
7
-
- 		64
Gd
2
7
1
- 		65
Tb
2
9
-
- 		66
Dy
2
10
-
- 		67
Ho
2
11
-
- 		68
Er
2
12
-
- 		69
Tm
2
13
-
- 		70
Yb
2
14
-
- 		71
Lu
2
14
1
- 		72
Hf
2
14
2
- 		73
Ta
2
14
3
- 		74
W
2
14
4
- 		75
Re
2
14
5
- 		76
Os
2
14
6
- 		77
Ir
2
14
7
- 		78
Pt
1
14
9
- 		79
Au
1
14
10
- 		80
Hg
2
14
10
- 		81
Tl
2
14
10
1 		82
Pb
2
14
10
2 		83
Bi
2
14
10
3 		84
Po
2
14
10
4 		85
At
2
14
10
5 		86
Rn
2
14
10
6
[Rn]
7s:
5f:
6d:
7p: 		87
Fr
1
-
-
- 		88
Ra
2
-
-
- 		89
Ac
2
-
1
- 		90
Th
2
-
2
- 		91
Pa
2
2
1
- 		92
U
2
3
1
- 		93
Np
2
4
1
- 		94
Pu
2
6
-
- 		95
Am
2
7
-
- 		96
Cm
2
7
1
- 		97
Bk
2
9
-
- 		98
Cf
2
10
-
- 		99
Es
2
11
-
- 		100
Fm
2
12
-
- 		101
Md
2
13
-
- 		102
No
2
14
-
- 		103
Lr
2
14
-
1 		104
Rf
2
14
2
- 		105
Db
2
14
3
- 		106
Sg
2
14
4
- 		107
Bh
2
14
5
- 		108
Hs
2
14
6
- 		109
Mt
2
14
7
- 		110
Ds
2
14
8
- 		111
Rg
2
14
9
- 		112
Cn
2
14
10
- 		113
Nh
2
14
10
1 		114
Fl
2
14
10
2 		115
Mc
2
14
10
3 		116
Lv
2
14
10
4 		117
Ts
2
14
10
5 		118
Og
2
14
10
6
s-block f-block d-block p-block

Variations

Period 1

Main article: Period 1 element

Although the modern periodic table is standard today, the placement of the period 1 elements hydrogen and helium remains an open issue under discussion, and some variation can be found. Following their respective s and s electron configurations, hydrogen would be placed in group 1, and helium would be placed in group 2. The group 1 placement of hydrogen is common, but helium is almost always placed in group 18 with the other noble gases. The debate has to do with conflicting understandings of the extent to which chemical or electronic properties should decide periodic table placement.

Like the group 1 metals, hydrogen has one electron in its outermost shell and typically loses its only electron in chemical reactions. Hydrogen has some metal-like chemical properties, being able to displace some metals from their salts. But it forms a diatomic nonmetallic gas at standard conditions, unlike the alkali metals which are reactive solid metals. This and hydrogen's formation of hydrides, in which it gains an electron, brings it close to the properties of the halogens which do the same (though it is rarer for hydrogen to form H than H). Moreover, the lightest two halogens (fluorine and chlorine) are gaseous like hydrogen at standard conditions. Some properties of hydrogen are not a good fit for either group: hydrogen is neither highly oxidizing nor highly reducing and is not reactive with water. Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens, but matches neither group perfectly, and is thus difficult to place by its chemistry. Therefore, while the electronic placement of hydrogen in group 1 predominates, some rarer arrangements show either hydrogen in group 17, duplicate hydrogen in both groups 1 and 17, or float it separately from all groups. This last option has nonetheless been criticized by the chemist and philosopher of science Eric Scerri on the grounds that it appears to imply that hydrogen is above the periodic law altogether, unlike all the other elements.

Helium is the only element that routinely occupies a position in the periodic table that is not consistent with its electronic structure. It has two electrons in its outermost shell, whereas the other noble gases have eight; and it is an s-block element, whereas all other noble gases are p-block elements. However it is unreactive at standard conditions, and has a full outer shell: these properties are like the noble gases in group 18, but not at all like the reactive alkaline earth metals of group 2. For these reasons helium is nearly universally placed in group 18 which its properties best match; a proposal to move helium to group 2 was rejected by IUPAC in 1988 for these reasons. Nonetheless, helium is still occasionally placed in group 2 today, and some of its physical and chemical properties are closer to the group 2 elements and support the electronic placement. Solid helium crystallises in a hexagonal close-packed structure, which matches beryllium and magnesium in group 2, but not the other noble gases in group 18. Recent theoretical developments in noble gas chemistry, in which helium is expected to show slightly less inertness than neon and to form (HeO)(LiF)2 with a structure similar to the analogous beryllium compound (but with no expected neon analogue), have resulted in more chemists advocating a placement of helium in group 2. This relates to the electronic argument, as the reason for neon's greater inertness is repulsion from its filled p-shell that helium lacks, though realistically it is unlikely that helium-containing molecules will be stable outside extreme low-temperature conditions (around 10 K).

The first-row anomaly in the periodic table has additionally been cited to support moving helium to group 2. It arises because the first orbital of any type is unusually small, since unlike its higher analogues, it does not experience interelectronic repulsion from a smaller orbital of the same type. This makes the first row of elements in each block unusually small, and such elements tend to exhibit characteristic kinds of anomalies for their group. Some chemists arguing for the repositioning of helium have pointed out that helium exhibits these anomalies if it is placed in group 2, but not if it is placed in group 18: on the other hand, neon, which would be the first group 18 element if helium was removed from that spot, does exhibit those anomalies. The relationship between helium and beryllium is then argued to resemble that between hydrogen and lithium, a placement which is much more commonly accepted. For example, because of this trend in the sizes of orbitals, a large difference in atomic radii between the first and second members of each main group is seen in groups 1 and 13–17: it exists between neon and argon, and between helium and beryllium, but not between helium and neon. This similarly affects the noble gases' boiling points and solubilities in water, where helium is too close to neon, and the large difference characteristic between the first two elements of a group appears only between neon and argon. Moving helium to group 2 makes this trend consistent in groups 2 and 18 as well, by making helium the first group 2 element and neon the first group 18 element: both exhibit the characteristic properties of a kainosymmetric first element of a group. The group 18 placement of helium nonetheless remains near-universal due to its extreme inertness. Additionally, tables that float both hydrogen and helium outside all groups may rarely be encountered.

Group 3

Main article: Group 3 element § Composition Group 3: Sc, Y, Lu, Lr Correct
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Correct depiction of Group 3 Group 3: Sc, Y, La, Ac Incorrect
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
Incorrect depiction of Group 3

In many periodic tables, the f-block is shifted one element to the right, so that lanthanum and actinium become d-block elements in group 3, and Ce–Lu and Th–Lr form the f-block. Thus the d-block is split into two very uneven portions. This is a holdover from early mistaken measurements of electron configurations; modern measurements are more consistent with the form with lutetium and lawrencium in group 3, and with La–Yb and Ac–No as the f-block.

The 4f shell is completely filled at ytterbium, and for that reason Lev Landau and Evgeny Lifshitz in 1948 considered it incorrect to group lutetium as an f-block element. They did not yet take the step of removing lanthanum from the d-block as well, but Jun Kondō realized in 1963 that lanthanum's low-temperature superconductivity implied the activity of its 4f shell. In 1965, David C. Hamilton linked this observation to its position in the periodic table, and argued that the f-block should be composed of the elements La–Yb and Ac–No. Since then, physical, chemical, and electronic evidence has supported this assignment. The issue was brought to wide attention by William B. Jensen in 1982, and the reassignment of lutetium and lawrencium to group 3 was supported by IUPAC reports dating from 1988 (when the 1–18 group numbers were recommended) and 2021. The variation nonetheless still exists because most textbook writers are not aware of the issue.

A third form can sometimes be encountered in which the spaces below yttrium in group 3 are left empty, such as the table appearing on the IUPAC web site, but this creates an inconsistency with quantum mechanics by making the f-block 15 elements wide (La–Lu and Ac–Lr) even though only 14 electrons can fit in an f-subshell. There is moreover some confusion in the literature on which elements are then implied to be in group 3. While the 2021 IUPAC report noted that 15-element-wide f-blocks are supported by some practitioners of a specialized branch of relativistic quantum mechanics focusing on the properties of superheavy elements, the project's opinion was that such interest-dependent concerns should not have any bearing on how the periodic table is presented to "the general chemical and scientific community". Other authors focusing on superheavy elements since clarified that the "15th entry of the f-block represents the first slot of the d-block which is left vacant to indicate the place of the f-block inserts", which would imply that this form still has lutetium and lawrencium (the 15th entries in question) as d-block elements in group 3. Indeed, when IUPAC publications expand the table to 32 columns, they make this clear and place lutetium and lawrencium under yttrium in group 3.

Several arguments in favour of Sc-Y-La-Ac can be encountered in the literature, but they have been challenged as being logically inconsistent. For example, it has been argued that lanthanum and actinium cannot be f-block elements because as individual gas-phase atoms, they have not begun to fill the f-subshells. But the same is true of thorium which is never disputed as an f-block element, and this argument overlooks the problem on the other end: that the f-shells complete filling at ytterbium and nobelium, matching the Sc-Y-Lu-Lr form, and not at lutetium and lawrencium as the Sc-Y-La-Ac form would have it. Not only are such exceptional configurations in the minority, but they have also in any case never been considered as relevant for positioning any other elements on the periodic table: in gaseous atoms, the d-shells complete their filling at copper, palladium, and gold, but it is universally accepted by chemists that these configurations are exceptional and that the d-block really ends in accordance with the Madelung rule at zinc, cadmium, and mercury. The relevant fact for placement is that lanthanum and actinium (like thorium) have valence f-orbitals that can become occupied in chemical environments, whereas lutetium and lawrencium do not: their f-shells are in the core, and cannot be used for chemical reactions. Thus the relationship between yttrium and lanthanum is only a secondary relationship between elements with the same number of valence electrons but different kinds of valence orbitals, such as that between chromium and uranium; whereas the relationship between yttrium and lutetium is primary, sharing both valence electron count and valence orbital type.

Periodic trends

Main article: Periodic trends

As chemical reactions involve the valence electrons, elements with similar outer electron configurations may be expected to react similarly and form compounds with similar proportions of elements in them. Such elements are placed in the same group, and thus there tend to be clear similarities and trends in chemical behaviour as one proceeds down a group. As analogous configurations occur at regular intervals, the properties of the elements thus exhibit periodic recurrences, hence the name of the periodic table and the periodic law. These periodic recurrences were noticed well before the underlying theory that explains them was developed.

Atomic radius

Historically, the physical size of atoms was unknown until the early 20th century. The first calculated estimate of the atomic radius of hydrogen was published by physicist Arthur Haas in 1910 to within an order of magnitude (a factor of 10) of the accepted value, the Bohr radius (~0.529 Å). In his model, Haas used a single-electron configuration based on the classical atomic model proposed by J. J. Thomson in 1904, often called the plum-pudding model.

Atomic radii (the size of atoms) are dependent on the sizes of their outermost orbitals. They generally decrease going left to right along the main-group elements, because the nuclear charge increases but the outer electrons are still in the same shell. However, going down a column, the radii generally increase, because the outermost electrons are in higher shells that are thus further away from the nucleus. The first row of each block is abnormally small, due to an effect called kainosymmetry or primogenic repulsion: the 1s, 2p, 3d, and 4f subshells have no inner analogues. For example, the 2p orbitals do not experience strong repulsion from the 1s and 2s orbitals, which have quite different angular charge distributions, and hence are not very large; but the 3p orbitals experience strong repulsion from the 2p orbitals, which have similar angular charge distributions. Thus higher s-, p-, d-, and f-subshells experience strong repulsion from their inner analogues, which have approximately the same angular distribution of charge, and must expand to avoid this. This makes significant differences arise between the small 2p elements, which prefer multiple bonding, and the larger 3p and higher p-elements, which do not. Similar anomalies arise for the 1s, 2p, 3d, 4f, and the hypothetical 5g elements: the degree of this first-row anomaly is highest for the s-block, is moderate for the p-block, and is less pronounced for the d- and f-blocks.

In the transition elements, an inner shell is filling, but the size of the atom is still determined by the outer electrons. The increasing nuclear charge across the series and the increased number of inner electrons for shielding somewhat compensate each other, so the decrease in radius is smaller. The 4p and 5d atoms, coming immediately after new types of transition series are first introduced, are smaller than would have been expected, because the added core 3d and 4f subshells provide only incomplete shielding of the nuclear charge for the outer electrons. Hence for example gallium atoms are slightly smaller than aluminium atoms. Together with kainosymmetry, this results in an even-odd difference between the periods (except in the s-block) that is sometimes known as secondary periodicity: elements in even periods have smaller atomic radii and prefer to lose fewer electrons, while elements in odd periods (except the first) differ in the opposite direction. Thus for example many properties in the p-block show a zigzag rather than a smooth trend along the group. For example, phosphorus and antimony in odd periods of group 15 readily reach the +5 oxidation state, whereas nitrogen, arsenic, and bismuth in even periods prefer to stay at +3. A similar situation holds for the d-block, with lutetium through tungsten atoms being slightly smaller than yttrium through molybdenum atoms respectively.

Liquid mercury. Its liquid state at standard conditions is the result of relativistic effects.

Thallium and lead atoms are about the same size as indium and tin atoms respectively, but from bismuth to radon the 6p atoms are larger than the analogous 5p atoms. This happens because when atomic nuclei become highly charged, special relativity becomes needed to gauge the effect of the nucleus on the electron cloud. These relativistic effects result in heavy elements increasingly having differing properties compared to their lighter homologues in the periodic table. Spin–orbit interaction splits the p-subshell: one p-orbital is relativistically stabilized and shrunken (it fills in thallium and lead), but the other two (filling in bismuth through radon) are relativistically destabilized and expanded. Relativistic effects also explain why gold is golden and mercury is a liquid at room temperature. They are expected to become very strong in the late seventh period, potentially leading to a collapse of periodicity. Electron configurations are only clearly known until element 108 (hassium), and experimental chemistry beyond 108 has only been done for elements 112 (copernicium) through 115 (moscovium), so the chemical characterization of the heaviest elements remains a topic of current research.

The trend that atomic radii decrease from left to right is also present in ionic radii, though it is more difficult to examine because the most common ions of consecutive elements normally differ in charge. Ions with the same electron configuration decrease in size as their atomic number rises, due to increased attraction from the more positively charged nucleus: thus for example ionic radii decrease in the series Se, Br, Rb, Sr, Y, Zr, Nb, Mo, Tc. Ions of the same element get smaller as more electrons are removed, because the attraction from the nucleus begins to outweigh the repulsion between electrons that causes electron clouds to expand: thus for example ionic radii decrease in the series V, V, V, V.

Ionisation energy

Graph of first ionisation energies of the elements in electronvolts (predictions used for elements 109–118)

The first ionisation energy of an atom is the energy required to remove an electron from it. This varies with the atomic radius: ionisation energy increases left to right and down to up, because electrons that are closer to the nucleus are held more tightly and are more difficult to remove. Ionisation energy thus is minimized at the first element of each period – hydrogen and the alkali metals – and then generally rises until it reaches the noble gas at the right edge of the period. There are some exceptions to this trend, such as oxygen, where the electron being removed is paired and thus interelectronic repulsion makes it easier to remove than expected.

In the transition series, the outer electrons are preferentially lost even though the inner orbitals are filling. For example, in the 3d series, the 4s electrons are lost first even though the 3d orbitals are being filled. The shielding effect of adding an extra 3d electron approximately compensates the rise in nuclear charge, and therefore the ionisation energies stay mostly constant, though there is a small increase especially at the end of each transition series.

As metal atoms tend to lose electrons in chemical reactions, ionisation energy is generally correlated with chemical reactivity, although there are other factors involved as well.

Electron affinity

Trend in electron affinities

The opposite property to ionisation energy is the electron affinity, which is the energy released when adding an electron to the atom. A passing electron will be more readily attracted to an atom if it feels the pull of the nucleus more strongly, and especially if there is an available partially filled outer orbital that can accommodate it. Therefore, electron affinity tends to increase down to up and left to right. The exception is the last column, the noble gases, which have a full shell and have no room for another electron. This gives the halogens in the next-to-last column the highest electron affinities.

Some atoms, like the noble gases, have no electron affinity: they cannot form stable gas-phase anions. (They can form metastable resonances if the incoming electron arrives with enough kinetic energy, but these inevitably and rapidly autodetach: for example, the lifetime of the most long-lived He level is about 359 microseconds.) The noble gases, having high ionisation energies and no electron affinity, have little inclination towards gaining or losing electrons and are generally unreactive.

Some exceptions to the trends occur: oxygen and fluorine have lower electron affinities than their heavier homologues sulfur and chlorine, because they are small atoms and hence the newly added electron would experience significant repulsion from the already present ones. For the nonmetallic elements, electron affinity likewise somewhat correlates with reactivity, but not perfectly since other factors are involved. For example, fluorine has a lower electron affinity than chlorine (because of extreme interelectronic repulsion for the very small fluorine atom), but is more reactive.

Valence and oxidation states

Lead(II) oxide (PbO, left) and lead(IV) oxide (PbO2, right), the two stable oxides of lead

The valence of an element can be defined either as the number of hydrogen atoms that can combine with it to form a simple binary hydride, or as twice the number of oxygen atoms that can combine with it to form a simple binary oxide (that is, not a peroxide or a superoxide). The valences of the main-group elements are directly related to the group number: the hydrides in the main groups 1–2 and 13–17 follow the formulae MH, MH2, MH3, MH4, MH3, MH2, and finally MH. The highest oxides instead increase in valence, following the formulae M2O, MO, M2O3, MO2, M2O5, MO3, M2O7. Today the notion of valence has been extended by that of the oxidation state, which is the formal charge left on an element when all other elements in a compound have been removed as their ions.

The electron configuration suggests a ready explanation from the number of electrons available for bonding; indeed, the number of valence electrons starts at 1 in group 1, and then increases towards the right side of the periodic table, only resetting at 3 whenever each new block starts. Thus in period 6, Cs–Ba have 1–2 valence electrons; La–Yb have 3–16; Lu–Hg have 3–12; and Tl–Rn have 3–8. However, towards the right side of the d- and f-blocks, the theoretical maximum corresponding to using all valence electrons is not achievable at all; the same situation affects oxygen, fluorine, and the light noble gases up to krypton.

Number of valence electrons
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1 H
1 		He
2
2 Li
1 		Be
2 		B
3 		C
4 		N
5 		O
6 		F
7 		Ne
8
3 Na
1 		Mg
2 		Al
3 		Si
4 		P
5 		S
6 		Cl
7 		Ar
8
4 K
1 		Ca
2 		Sc
3 		Ti
4 		V
5 		Cr
6 		Mn
7 		Fe
8 		Co
9 		Ni
10 		Cu
11 		Zn
12 		Ga
3 		Ge
4 		As
5 		Se
6 		Br
7 		Kr
8
5 Rb
1 		Sr
2 		Y
3 		Zr
4 		Nb
5 		Mo
6 		Tc
7 		Ru
8 		Rh
9 		Pd
10 		Ag
11 		Cd
12 		In
3 		Sn
4 		Sb
5 		Te
6 		I
7 		Xe
8
6 Cs
1 		Ba
2 		La
3 		Ce
4 		Pr
5 		Nd
6 		Pm
7 		Sm
8 		Eu
9 		Gd
10 		Tb
11 		Dy
12 		Ho
13 		Er
14 		Tm
15 		Yb
16 		Lu
3 		Hf
4 		Ta
5 		W
6 		Re
7 		Os
8 		Ir
9 		Pt
10 		Au
11 		Hg
12 		Tl
3 		Pb
4 		Bi
5 		Po
6 		At
7 		Rn
8
7 Fr
1 		Ra
2 		Ac
3 		Th
4 		Pa
5 		U
6 		Np
7 		Pu
8 		Am
9 		Cm
10 		Bk
11 		Cf
12 		Es
13 		Fm
14 		Md
15 		No
16 		Lr
3 		Rf
4 		Db
5 		Sg
6 		Bh
7 		Hs
8 		Mt
9 		Ds
10 		Rg
11 		Cn
12 		Nh
3 		Fl
4 		Mc
5 		Lv
6 		Ts
7 		Og
8

A full explanation requires considering the energy that would be released in forming compounds with different valences rather than simply considering electron configurations alone. For example, magnesium forms Mg rather than Mg cations when dissolved in water, because the latter would spontaneously disproportionate into Mg and Mg cations. This is because the enthalpy of hydration (surrounding the cation with water molecules) increases in magnitude with the charge and radius of the ion. In Mg, the outermost orbital (which determines ionic radius) is still 3s, so the hydration enthalpy is small and insufficient to compensate the energy required to remove the electron; but ionizing again to Mg uncovers the core 2p subshell, making the hydration enthalpy large enough to allow magnesium(II) compounds to form. For similar reasons, the common oxidation states of the heavier p-block elements (where the ns electrons become lower in energy than the np) tend to vary by steps of 2, because that is necessary to uncover an inner subshell and decrease the ionic radius (e.g. Tl uncovers 6s, and Tl uncovers 5d, so once thallium loses two electrons it tends to lose the third one as well). Analogous arguments based on orbital hybridization can be used for the less electronegative p-block elements.

Oxidation states of the transition metals. The solid dots show common oxidation states, and the hollow dots show possible but unlikely states.

For transition metals, common oxidation states are nearly always at least +2 for similar reasons (uncovering the next subshell); this holds even for the metals with anomalous ds or ds configurations (except for silver), because repulsion between d-electrons means that the movement of the second electron from the s- to the d-subshell does not appreciably change its ionisation energy. Because ionizing the transition metals further does not uncover any new inner subshells, their oxidation states tend to vary by steps of 1 instead. The lanthanides and late actinides generally show a stable +3 oxidation state, removing the outer s-electrons and then (usually) one electron from the (n−2)f-orbitals, that are similar in energy to ns. The common and maximum oxidation states of the d- and f-block elements tend to depend on the ionisation energies. As the energy difference between the (n−1)d and ns orbitals rises along each transition series, it becomes less energetically favourable to ionize further electrons. Thus, the early transition metal groups tend to prefer higher oxidation states, but the +2 oxidation state becomes more stable for the late transition metal groups. The highest formal oxidation state thus increases from +3 at the beginning of each d-block row, to +7 or +8 in the middle (e.g. OsO4), and then decrease to +2 at the end. The lanthanides and late actinides usually have high fourth ionisation energies and hence rarely surpass the +3 oxidation state, whereas early actinides have low fourth ionisation energies and so for example neptunium and plutonium can reach +7. The very last actinides go further than the lanthanides towards low oxidation states: mendelevium is more easily reduced to the +2 state than thulium or even europium (the lanthanide with the most stable +2 state, on account of its half-filled f-shell), and nobelium outright favours +2 over +3, in contrast to ytterbium.

As elements in the same group share the same valence configurations, they usually exhibit similar chemical behaviour. For example, the alkali metals in the first group all have one valence electron, and form a very homogeneous class of elements: they are all soft and reactive metals. However, there are many factors involved, and groups can often be rather heterogeneous. For instance, hydrogen also has one valence electron and is in the same group as the alkali metals, but its chemical behaviour is quite different. The stable elements of group 14 comprise a nonmetal (carbon), two semiconductors (silicon and germanium), and two metals (tin and lead); they are nonetheless united by having four valence electrons. This often leads to similarities in maximum and minimum oxidation states (e.g. sulfur and selenium in group 16 both have maximum oxidation state +6, as in SO3 and SeO3, and minimum oxidation state −2, as in sulfides and selenides); but not always (e.g. oxygen is not known to form oxidation state +6, despite being in the same group as sulfur and selenium).

Electronegativity

A water molecule is put into a see-through egg shape, which is colour-coded by electrostatic potential. A concentration of red is near the top of the shape, where the oxygen atom is, and gradually shifts through yellow, green, and then to blue near the lower-right and lower-left corners of the shape where the hydrogen atoms are.
Electrostatic potential map of a water molecule, where the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms

Another important property of elements is their electronegativity. Atoms can form covalent bonds to each other by sharing electrons in pairs, creating an overlap of valence orbitals. The degree to which each atom attracts the shared electron pair depends on the atom's electronegativity – the tendency of an atom towards gaining or losing electrons. The more electronegative atom will tend to attract the electron pair more, and the less electronegative (or more electropositive) one will attract it less. In extreme cases, the electron can be thought of as having been passed completely from the more electropositive atom to the more electronegative one, though this is a simplification. The bond then binds two ions, one positive (having given up the electron) and one negative (having accepted it), and is termed an ionic bond.

Electronegativity depends on how strongly the nucleus can attract an electron pair, and so it exhibits a similar variation to the other properties already discussed: electronegativity tends to fall going up to down, and rise going left to right. The alkali and alkaline earth metals are among the most electropositive elements, while the chalcogens, halogens, and noble gases are among the most electronegative ones.

Electronegativity is generally measured on the Pauling scale, on which the most electronegative reactive atom (fluorine) is given electronegativity 4.0, and the least electronegative atom (caesium) is given electronegativity 0.79. In fact neon is the most electronegative element, but the Pauling scale cannot measure its electronegativity because it does not form covalent bonds with most elements.

An element's electronegativity varies with the identity and number of the atoms it is bonded to, as well as how many electrons it has already lost: an atom becomes more electronegative when it has lost more electrons. This sometimes makes a large difference: lead in the +2 oxidation state has electronegativity 1.87 on the Pauling scale, while lead in the +4 oxidation state has electronegativity 2.33.

Metallicity

The diamond-cubic structure, a giant covalent structure adopted by carbon (as diamond), as well as by silicon, germanium, and (grey) tin, all in group 14.
(In grey tin, the band gap vanishes and metallization occurs. Tin has another allotrope, white tin, whose structure is even more metallic.)

A simple substance is a substance formed from atoms of one chemical element. The simple substances of the more electronegative atoms tend to share electrons (form covalent bonds) with each other. They form either small molecules (like hydrogen or oxygen, whose atoms bond in pairs) or giant structures stretching indefinitely (like carbon or silicon). The noble gases simply stay as single atoms, as they already have a full shell. Substances composed of discrete molecules or single atoms are held together by weaker attractive forces between the molecules, such as the London dispersion force: as electrons move within the molecules, they create momentary imbalances of electrical charge, which induce similar imbalances on nearby molecules and create synchronized movements of electrons across many neighbouring molecules.

Graphite and diamond, two allotropes of carbon

The more electropositive atoms, however, tend to instead lose electrons, creating a "sea" of electrons engulfing cations. The outer orbitals of one atom overlap to share electrons with all its neighbours, creating a giant structure of molecular orbitals extending over all the atoms. This negatively charged "sea" pulls on all the ions and keeps them together in a metallic bond. Elements forming such bonds are often called metals; those which do not are often called nonmetals. Some elements can form multiple simple substances with different structures: these are called allotropes. For example, diamond and graphite are two allotropes of carbon.

The metallicity of an element can be predicted from electronic properties. When atomic orbitals overlap during metallic or covalent bonding, they create both bonding and antibonding molecular orbitals of equal capacity, with the antibonding orbitals of higher energy. Net bonding character occurs when there are more electrons in the bonding orbitals than there are in the antibonding orbitals. Metallic bonding is thus possible when the number of electrons delocalized by each atom is less than twice the number of orbitals contributing to the overlap. This is the situation for elements in groups 1 through 13; they also have too few valence electrons to form giant covalent structures where all atoms take equivalent positions, and so almost all of them metallise. The exceptions are hydrogen and boron, which have too high an ionisation energy. Hydrogen thus forms a covalent H2 molecule, and boron forms a giant covalent structure based on icosahedral B12 clusters. In a metal, the bonding and antibonding orbitals have overlapping energies, creating a single band that electrons can freely flow through, allowing for electrical conduction.

Graph of carbon atoms being brought together to form a diamond crystal, demonstrating formation of the electronic band structure and band gap. The right graph shows the energy levels as a function of the spacing between atoms. When far apart (right side of graph) all the atoms have discrete valence orbitals p and s with the same energies. However, when the atoms come closer (left side), their electron orbitals begin to spatially overlap. The Pauli exclusion principle prohibits them from having the same energy, so the orbitals hybridize into N molecular orbitals each with a different energy, where N is the number of atoms in the crystal. Since N is such a large number, adjacent orbitals are extremely close together in energy so the orbitals can be considered a continuous energy band. At the actual diamond crystal cell size (denoted by a), two bands are formed, called the valence and conduction bands, separated by a 5.5 eV band gap. (Here only the valence 2s and 2p electrons have been illustrated; the 1s orbitals do not significantly overlap, so the bands formed from them are much narrower.)

In group 14, both metallic and covalent bonding become possible. In a diamond crystal, covalent bonds between carbon atoms are strong, because they have a small atomic radius and thus the nucleus has more of a hold on the electrons. Therefore, the bonding orbitals that result are much lower in energy than the antibonding orbitals, and there is no overlap, so electrical conduction becomes impossible: carbon is a nonmetal. However, covalent bonding becomes weaker for larger atoms and the energy gap between the bonding and antibonding orbitals decreases. Therefore, silicon and germanium have smaller band gaps and are semiconductors at ambient conditions: electrons can cross the gap when thermally excited. (Boron is also a semiconductor at ambient conditions.) The band gap disappears in tin, so that tin and lead become metals. As the temperature rises, all nonmetals develop some semiconducting properties, to a greater or lesser extent depending on the size of the band gap. Thus metals and nonmetals may be distinguished by the temperature dependence of their electrical conductivity: a metal's conductivity lowers as temperature rises (because thermal motion makes it more difficult for the electrons to flow freely), whereas a nonmetal's conductivity rises (as more electrons may be excited to cross the gap).

Elements in groups 15 through 17 have too many electrons to form giant covalent molecules that stretch in all three dimensions. For the lighter elements, the bonds in small diatomic molecules are so strong that a condensed phase is disfavoured: thus nitrogen (N2), oxygen (O2), white phosphorus and yellow arsenic (P4 and As4), sulfur and red selenium (S8 and Se8), and the stable halogens (F2, Cl2, Br2, and I2) readily form covalent molecules with few atoms. The heavier ones tend to form long chains (e.g. red phosphorus, grey selenium, tellurium) or layered structures (e.g. carbon as graphite, black phosphorus, grey arsenic, antimony, bismuth) that only extend in one or two rather than three dimensions. Both kinds of structures can be found as allotropes of phosphorus, arsenic, and selenium, although the long-chained allotropes are more stable in all three. As these structures do not use all their orbitals for bonding, they end up with bonding, nonbonding, and antibonding bands in order of increasing energy. Similarly to group 14, the band gaps shrink for the heavier elements and free movement of electrons between the chains or layers becomes possible. Thus for example black phosphorus, black arsenic, grey selenium, tellurium, and iodine are semiconductors; grey arsenic, antimony, and bismuth are semimetals (exhibiting quasi-metallic conduction, with a very small band overlap); and polonium and probably astatine are true metals. Finally, the natural group 18 elements all stay as individual atoms.

The dividing line between metals and nonmetals is roughly diagonal from top left to bottom right, with the transition series appearing to the left of this diagonal (as they have many available orbitals for overlap). This is expected, as metallicity tends to be correlated with electropositivity and the willingness to lose electrons, which increases right to left and up to down. Thus the metals greatly outnumber the nonmetals. Elements near the borderline are difficult to classify: they tend to have properties that are intermediate between those of metals and nonmetals, and may have some properties characteristic of both. They are often termed semimetals or metalloids. The term "semimetal" used in this sense should not be confused with its strict physical sense having to do with band structure: bismuth is physically a semimetal, but is generally considered a metal by chemists.

The following table considers the most stable allotropes at standard conditions. The elements coloured yellow form simple substances that are well-characterised by metallic bonding. Elements coloured light blue form giant network covalent structures, whereas those coloured dark blue form small covalently bonded molecules that are held together by weaker van der Waals forces. The noble gases are coloured in violet: their molecules are single atoms and no covalent bonding occurs. Greyed-out cells are for elements which have not been prepared in sufficient quantities for their most stable allotropes to have been characterized in this way. Theoretical considerations and current experimental evidence suggest that all of those elements would metallise if they could form condensed phases, except perhaps for oganesson.

Bonding of simple substances in the periodic table
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
Group →
↓ Period
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

Metallic Network covalent Molecular covalent Single atoms Unknown Background color shows bonding of simple substances in the periodic table. If there are several, the most stable allotrope is considered.

  • Iron, a metal Iron, a metal
  • Sulfur, a nonmetal Sulfur, a nonmetal
  • Arsenic, an element often called a semi-metal or metalloid Arsenic, an element often called a semi-metal or metalloid

Generally, metals are shiny and dense. They usually have high melting and boiling points due to the strength of the metallic bond, and are often malleable and ductile (easily stretched and shaped) because the atoms can move relative to each other without breaking the metallic bond. They conduct electricity because their electrons are free to move in all three dimensions. Similarly, they conduct heat, which is transferred by the electrons as extra kinetic energy: they move faster. These properties persist in the liquid state, as although the crystal structure is destroyed on melting, the atoms still touch and the metallic bond persists, though it is weakened. Metals tend to be reactive towards nonmetals. Some exceptions can be found to these generalizations: for example, beryllium, chromium, manganese, antimony, bismuth, and uranium are brittle (not an exhaustive list); chromium is extremely hard; gallium, rubidium, caesium, and mercury are liquid at or close to room temperature; and noble metals such as gold are chemically very inert.

Nonmetals exhibit different properties. Those forming giant covalent crystals exhibit high melting and boiling points, as it takes considerable energy to overcome the strong covalent bonds. Those forming discrete molecules are held together mostly by dispersion forces, which are more easily overcome; thus they tend to have lower melting and boiling points, and many are liquids or gases at room temperature. Nonmetals are often dull-looking. They tend to be reactive towards metals, except for the noble gases, which are inert towards most substances. They are brittle when solid as their atoms are held tightly in place. They are less dense and conduct electricity poorly, because there are no mobile electrons. Near the borderline, band gaps are small and thus many elements in that region are semiconductors, such as silicon, germanium, and tellurium. Selenium has both a semiconducting grey allotrope and an insulating red allotrope; arsenic has a metallic grey allotrope, a semiconducting black allotrope, and an insulating yellow allotrope (though the last is unstable at ambient conditions). Again there are exceptions; for example, diamond has the highest thermal conductivity of all known materials, greater than any metal.

It is common to designate a class of metalloids straddling the boundary between metals and nonmetals, as elements in that region are intermediate in both physical and chemical properties. However, no consensus exists in the literature for precisely which elements should be so designated. When such a category is used, silicon, germanium, arsenic, and tellurium are almost always included, and boron and antimony usually are; but most sources include other elements as well, without agreement on which extra elements should be added, and some others subtract from this list instead. For example, unlike all the other elements generally considered metalloids or nonmetals, antimony's only stable form has metallic conductivity. Moreover, the element resembles bismuth and, more generally, the other p-block metals in its physical and chemical behaviour. On this basis some authors have argued that it is better classified as a metal than as a metalloid. On the other hand, selenium has some semiconducting properties in its most stable form (though it also has insulating allotropes) and it has been argued that it should be considered a metalloid – though this situation also holds for phosphorus, which is a much rarer inclusion among the metalloids.

Further manifestations of periodicity

There are some other relationships throughout the periodic table between elements that are not in the same group, such as the diagonal relationships between elements that are diagonally adjacent (e.g. lithium and magnesium). Some similarities can also be found between the main groups and the transition metal groups, or between the early actinides and early transition metals, when the elements have the same number of valence electrons. Thus uranium somewhat resembles chromium and tungsten in group 6, as all three have six valence electrons. Relationships between elements with the same number of valence electrons but different types of valence orbital have been called secondary or isodonor relationships: they usually have the same maximum oxidation states, but not the same minimum oxidation states. For example, chlorine and manganese both have +7 as their maximum oxidation state (e.g. Cl2O7 and Mn2O7), but their respective minimum oxidation states are −1 (e.g. HCl) and −3 (K2). Elements with the same number of valence vacancies but different numbers of valence electrons are related by a tertiary or isoacceptor relationship: they usually have similar minimum but not maximum oxidation states. For example, hydrogen and chlorine both have −1 as their minimum oxidation state (in hydrides and chlorides), but hydrogen's maximum oxidation state is +1 (e.g. H2O) while chlorine's is +7.

Many other physical properties of the elements exhibit periodic variation in accordance with the periodic law, such as melting points, boiling points, heats of fusion, heats of vaporization, atomisation energy, and so on. Similar periodic variations appear for the compounds of the elements, which can be observed by comparing hydrides, oxides, sulfides, halides, and so on. Chemical properties are more difficult to describe quantitatively, but likewise exhibit their own periodicities. Examples include the variation in the acidic and basic properties of the elements and their compounds, the stabilities of compounds, and methods of isolating the elements. Periodicity is and has been used very widely to predict the properties of unknown new elements and new compounds, and is central to modern chemistry.

Classification of elements

A periodic table colour-coded to show some commonly used sets of similar elements. The categories and their boundaries differ somewhat between sources. Lutetium and lawrencium in group 3 are also transition metals.
   Alkali metals
   Alkaline earth metals
   Lanthanides
   Actinides
   Transition metals 		   Other metals
   Metalloids
   Other nonmetals
   Halogens
   Noble gases

Many terms have been used in the literature to describe sets of elements that behave similarly. The group names alkali metal, alkaline earth metal, triel, tetrel, pnictogen, chalcogen, halogen, and noble gas are acknowledged by IUPAC; the other groups can be referred to by their number, or by their first element (e.g., group 6 is the chromium group). Some divide the p-block elements from groups 13 to 16 by metallicity, although there is neither an IUPAC definition nor a precise consensus on exactly which elements should be considered metals, nonmetals, or semi-metals (sometimes called metalloids). Neither is there a consensus on what the metals succeeding the transition metals ought to be called, with post-transition metal and poor metal being among the possibilities having been used. Some advanced monographs exclude the elements of group 12 from the transition metals on the grounds of their sometimes quite different chemical properties, but this is not a universal practice and IUPAC does not presently mention it as allowable in its Principles of Chemical Nomenclature.

The lanthanides are considered to be the elements La–Lu, which are all very similar to each other: historically they included only Ce–Lu, but lanthanum became included by common usage. The rare earth elements (or rare earth metals) add scandium and yttrium to the lanthanides. Analogously, the actinides are considered to be the elements Ac–Lr (historically Th–Lr), although variation of properties in this set is much greater than within the lanthanides. IUPAC recommends the names lanthanoids and actinoids to avoid ambiguity, as the -ide suffix typically denotes a negative ion; however lanthanides and actinides remain common. With the increasing recognition of lutetium and lawrencium as d-block elements, some authors began to define the lanthanides as La–Yb and the actinides as Ac–No, matching the f-block. The transactinides or superheavy elements are the short-lived elements beyond the actinides, starting at lawrencium or rutherfordium (depending on where the actinides are taken to end).

Many more categorizations exist and are used according to certain disciplines. In astrophysics, a metal is defined as any element with atomic number greater than 2, i.e. anything except hydrogen and helium. The term "semimetal" has a different definition in physics than it does in chemistry: bismuth is a semimetal by physical definitions, but chemists generally consider it a metal. A few terms are widely used, but without any very formal definition, such as "heavy metal", which has been given such a wide range of definitions that it has been criticized as "effectively meaningless".

The scope of terms varies significantly between authors. For example, according to IUPAC, the noble gases extend to include the whole group, including the very radioactive superheavy element oganesson. However, among those who specialize in the superheavy elements, this is not often done: in this case "noble gas" is typically taken to imply the unreactive behaviour of the lighter elements of the group. Since calculations generally predict that oganesson should not be particularly inert due to relativistic effects, and may not even be a gas at room temperature if it could be produced in bulk, its status as a noble gas is often questioned in this context. Furthermore, national variations are sometimes encountered: in Japan, alkaline earth metals often do not include beryllium and magnesium as their behaviour is different from the heavier group 2 metals.

History

Main article: History of the periodic table See also: Timeline of chemical element discoveries

Early history

In 1817, German physicist Johann Wolfgang Döbereiner began to formulate one of the earliest attempts to classify the elements. In 1829, he found that he could form some of the elements into groups of three, with the members of each group having related properties. He termed these groups triads. Chlorine, bromine, and iodine formed a triad; as did calcium, strontium, and barium; lithium, sodium, and potassium; and sulfur, selenium, and tellurium. Today, all these triads form part of modern-day groups: the halogens, alkaline earth metals, alkali metals, and chalcogens. Various chemists continued his work and were able to identify more and more relationships between small groups of elements. However, they could not build one scheme that encompassed them all.

Newlands's table of the elements.
Newlands's table of the elements in 1866.

John Newlands published a letter in the Chemical News in February 1863 on the periodicity among the chemical elements. In 1864 Newlands published an article in the Chemical News showing that if the elements are arranged in the order of their atomic weights, those having consecutive numbers frequently either belong to the same group or occupy similar positions in different groups, and he pointed out that each eighth element starting from a given one is in this arrangement a kind of repetition of the first, like the eighth note of an octave in music (The Law of Octaves). However, Newlands's formulation only worked well for the main-group elements, and encountered serious problems with the others.

German chemist Lothar Meyer noted the sequences of similar chemical and physical properties repeated at periodic intervals. According to him, if the atomic weights were plotted as ordinates (i.e. vertically) and the atomic volumes as abscissas (i.e. horizontally)—the curve obtained a series of maximums and minimums—the most electropositive elements would appear at the peaks of the curve in the order of their atomic weights. In 1864, a book of his was published; it contained an early version of the periodic table containing 28 elements, and classified elements into six families by their valence—for the first time, elements had been grouped according to their valence. Works on organizing the elements by atomic weight had until then been stymied by inaccurate measurements of the atomic weights. In 1868, he revised his table, but this revision was published as a draft only after his death.

Mendeleev

Mendeleev's 1869 periodic tableMendeleev's 1871 periodic table

The definitive breakthrough came from the Russian chemist Dmitri Mendeleev. Although other chemists (including Meyer) had found some other versions of the periodic system at about the same time, Mendeleev was the most dedicated to developing and defending his system, and it was his system that most affected the scientific community. On 17 February 1869 (1 March 1869 in the Gregorian calendar), Mendeleev began arranging the elements and comparing them by their atomic weights. He began with a few elements, and over the course of the day his system grew until it encompassed most of the known elements. After he found a consistent arrangement, his printed table appeared in May 1869 in the journal of the Russian Chemical Society. When elements did not appear to fit in the system, he boldly predicted that either valencies or atomic weights had been measured incorrectly, or that there was a missing element yet to be discovered. In 1871, Mendeleev published a long article, including an updated form of his table, that made his predictions for unknown elements explicit. Mendeleev predicted the properties of three of these unknown elements in detail: as they would be missing heavier homologues of boron, aluminium, and silicon, he named them eka-boron, eka-aluminium, and eka-silicon ("eka" being Sanskrit for "one"). In 1875, the French chemist Paul-Émile Lecoq de Boisbaudran, working without knowledge of Mendeleev's prediction, discovered a new element in a sample of the mineral sphalerite, and named it gallium. He isolated the element and began determining its properties. Mendeleev, reading de Boisbaudran's publication, sent a letter claiming that gallium was his predicted eka-aluminium. Although Lecoq de Boisbaudran was initially sceptical, and suspected that Mendeleev was trying to take credit for his discovery, he later admitted that Mendeleev was correct. In 1879, the Swedish chemist Lars Fredrik Nilson discovered a new element, which he named scandium: it turned out to be eka-boron. Eka-silicon was found in 1886 by German chemist Clemens Winkler, who named it germanium. The properties of gallium, scandium, and germanium matched what Mendeleev had predicted. In 1889, Mendeleev noted at the Faraday Lecture to the Royal Institution in London that he had not expected to live long enough "to mention their discovery to the Chemical Society of Great Britain as a confirmation of the exactitude and generality of the periodic law". Even the discovery of the noble gases at the close of the 19th century, which Mendeleev had not predicted, fitted neatly into his scheme as an eighth main group.

Mendeleev nevertheless had some trouble fitting the known lanthanides into his scheme, as they did not exhibit the periodic change in valencies that the other elements did. After much investigation, the Czech chemist Bohuslav Brauner suggested in 1902 that the lanthanides could all be placed together in one group on the periodic table. He named this the "asteroid hypothesis" as an astronomical analogy: just as there is an asteroid belt instead of a single planet between Mars and Jupiter, so the place below yttrium was thought to be occupied by all the lanthanides instead of just one element.

Atomic number

Periodic table of Antonius van den Broek

After the internal structure of the atom was probed, amateur Dutch physicist Antonius van den Broek proposed in 1913 that the nuclear charge determined the placement of elements in the periodic table. The New Zealand physicist Ernest Rutherford coined the word "atomic number" for this nuclear charge. In van den Broek's published article he illustrated the first electronic periodic table showing the elements arranged according to the number of their electrons. Rutherford confirmed in his 1914 paper that Bohr had accepted the view of van den Broek.

The same year, English physicist Henry Moseley using X-ray spectroscopy confirmed van den Broek's proposal experimentally. Moseley determined the value of the nuclear charge of each element from aluminium to gold and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge. 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's research immediately resolved discrepancies between atomic weight and chemical properties; these were cases such as tellurium and iodine, where atomic number increases but atomic weight decreases. Although Moseley was soon killed in World War I, the Swedish physicist Manne Siegbahn continued his work up to uranium, and established that it was the element with the highest atomic number then known (92). Based on Moseley and Siegbahn's research, it was also known which atomic numbers corresponded to missing elements yet to be found: 43, 61, 72, 75, 85, and 87. (Element 75 had in fact already been found by Japanese chemist Masataka Ogawa in 1908 and named nipponium, but he mistakenly assigned it as element 43 instead of 75 and so his discovery was not generally recognized until later. The contemporarily accepted discovery of element 75 came in 1925, when Walter Noddack, Ida Tacke, and Otto Berg independently rediscovered it and gave it its present name, rhenium.)

The dawn of atomic physics also clarified the situation of isotopes. In the decay chains of the primordial radioactive elements thorium and uranium, it soon became evident that there were many apparent new elements that had different atomic weights but exactly the same chemical properties. In 1913, Frederick Soddy coined the term "isotope" to describe this situation, and considered isotopes to merely be different forms of the same chemical element. This furthermore clarified discrepancies such as tellurium and iodine: tellurium's natural isotopic composition is weighted towards heavier isotopes than iodine's, but tellurium has a lower atomic number.

Electron shells

The Danish physicist Niels Bohr applied Max Planck's idea of quantization to the atom. He concluded that the energy levels of electrons were quantised: only a discrete set of stable energy states were allowed. Bohr then attempted to understand periodicity through electron configurations, surmising in 1913 that the inner electrons should be responsible for the chemical properties of the element. In 1913, he produced the first electronic periodic table based on a quantum atom.

Bohr called his electron shells "rings" in 1913: atomic orbitals within shells did not exist at the time of his planetary model. Bohr explains in Part 3 of his famous 1913 paper that the maximum electrons in a shell is eight, writing, "We see, further, that a ring of n electrons cannot rotate in a single ring round a nucleus of charge ne unless n < 8." For smaller atoms, the electron shells would be filled as follows: "rings of electrons will only join if they contain equal numbers of electrons; and that accordingly the numbers of electrons on inner rings will only be 2, 4, 8." However, in larger atoms the innermost shell would contain eight electrons: "on the other hand, the periodic system of the elements strongly suggests that already in neon N = 10 an inner ring of eight electrons will occur." His proposed electron configurations for the atoms (shown to the right) mostly do not accord with those now known. They were improved further after the work of Arnold Sommerfeld and Edmund Stoner discovered more quantum numbers.

Bohr's electron configurations for light elements
Element Electrons per shell
4 2,2
6 2,4
7 4,3
8 4,2,2
9 4,4,1
10 8,2
11 8,2,1
16 8,4,2,2
18 8,8,2

The first one to systematically expand and correct the chemical potentials of Bohr's atomic theory was Walther Kossel in 1914 and in 1916. Kossel explained that in the periodic table new elements would be created as electrons were added to the outer shell. In Kossel's paper, he writes:

This leads to the conclusion that the electrons, which are added further, should be put into concentric rings or shells, on each of which ... only a certain number of electrons—namely, eight in our case—should be arranged. As soon as one ring or shell is completed, a new one has to be started for the next element; the number of electrons, which are most easily accessible, and lie at the outermost periphery, increases again from element to element and, therefore, in the formation of each new shell the chemical periodicity is repeated.

In a 1919 paper, Irving Langmuir postulated the existence of "cells" which we now call orbitals, which could each only contain two electrons each, and these were arranged in "equidistant layers" which we now call shells. He made an exception for the first shell to only contain two electrons. The chemist Charles Rugeley Bury suggested in 1921 that eight and eighteen electrons in a shell form stable configurations. Bury proposed that the electron configurations in transitional elements depended upon the valence electrons in their outer shell. He introduced the word transition to describe the elements now known as transition metals or transition elements. Bohr's theory was vindicated by the discovery of element 72: Georges Urbain claimed to have discovered it as the rare earth element celtium, but Bury and Bohr had predicted that element 72 could not be a rare earth element and had to be a homologue of zirconium. Dirk Coster and Georg von Hevesy searched for the element in zirconium ores and found element 72, which they named hafnium after Bohr's hometown of Copenhagen (Hafnia in Latin). Urbain's celtium proved to be simply purified lutetium (element 71). Hafnium and rhenium thus became the last stable elements to be discovered.

Prompted by Bohr, Wolfgang Pauli took up the problem of electron configurations in 1923. Pauli extended Bohr's scheme to use four quantum numbers, and formulated his exclusion principle which stated that no two electrons could have the same four quantum numbers. This explained the lengths of the periods in the periodic table (2, 8, 18, and 32), which corresponded to the number of electrons that each shell could occupy. In 1925, Friedrich Hund arrived at configurations close to the modern ones. As a result of these advances, periodicity became based on the number of chemically active or valence electrons rather than by the valences of the elements. The Aufbau principle that describes the electron configurations of the elements was first empirically observed by Erwin Madelung in 1926, though the first to publish it was Vladimir Karapetoff in 1930. In 1961, Vsevolod Klechkovsky derived the first part of the Madelung rule (that orbitals fill in order of increasing n + ℓ) from the Thomas–Fermi model; the complete rule was derived from a similar potential in 1971 by Yury N. Demkov and Valentin N. Ostrovsky.

Periodic table of Alfred Werner (1905), the first appearance of the long form

The quantum theory clarified the transition metals and lanthanides as forming their own separate groups, transitional between the main groups, although some chemists had already proposed tables showing them this way before then: the English chemist Henry Bassett did so in 1892, the Danish chemist Julius Thomsen in 1895, and the Swiss chemist Alfred Werner in 1905. Bohr used Thomsen's form in his 1922 Nobel Lecture; Werner's form is very similar to the modern 32-column form. In particular, this supplanted Brauner's asteroidal hypothesis.

The exact position of the lanthanides, and thus the composition of group 3, remained under dispute for decades longer because their electron configurations were initially measured incorrectly. On chemical grounds Bassett, Werner, and Bury grouped scandium and yttrium with lutetium rather than lanthanum (the former two left an empty space below yttrium as lutetium had not yet been discovered). Hund assumed in 1927 that all the lanthanide atoms had configuration 4f5d6s, on account of their prevailing trivalency. It is now known that the relationship between chemistry and electron configuration is more complicated than that. Early spectroscopic evidence seemed to confirm these configurations, and thus the periodic table was structured to have group 3 as scandium, yttrium, lanthanum, and actinium, with fourteen f-elements breaking up the d-block between lanthanum and hafnium. But it was later discovered that this is only true for four of the fifteen lanthanides (lanthanum, cerium, gadolinium, and lutetium), and that the other lanthanide atoms do not have a d-electron. In particular, ytterbium completes the 4f shell and thus Soviet physicists Lev Landau and Evgeny Lifshitz noted in 1948 that lutetium is correctly regarded as a d-block rather than an f-block element; that bulk lanthanum is an f-metal was first suggested by Jun Kondō in 1963, on the grounds of its low-temperature superconductivity. This clarified the importance of looking at low-lying excited states of atoms that can play a role in chemical environments when classifying elements by block and positioning them on the table. Many authors subsequently rediscovered this correction based on physical, chemical, and electronic concerns and applied it to all the relevant elements, thus making group 3 contain scandium, yttrium, lutetium, and lawrencium and having lanthanum through ytterbium and actinium through nobelium as the f-block rows: this corrected version achieves consistency with the Madelung rule and vindicates Bassett, Werner, and Bury's initial chemical placement.

In 1988, IUPAC released a report supporting this composition of group 3, a decision that was reaffirmed in 2021. Variation can still be found in textbooks on the composition of group 3, and some argumentation against this format is still published today, but chemists and physicists who have considered the matter largely agree on group 3 containing scandium, yttrium, lutetium, and lawrencium and challenge the counterarguments as being inconsistent.

Synthetic elements

Glenn T. Seaborg

By 1936, the pool of missing elements from hydrogen to uranium had shrunk to four: elements 43, 61, 85, and 87 remained missing. Element 43 eventually became the first element to be synthesized artificially via nuclear reactions rather than discovered in nature. It was discovered in 1937 by Italian chemists Emilio Segrè and Carlo Perrier, who named their discovery technetium, after the Greek word for "artificial". Elements 61 (promethium) and 85 (astatine) were likewise produced artificially in 1945 and 1940 respectively; element 87 (francium) became the last element to be discovered in nature, by French chemist Marguerite Perey in 1939. The elements beyond uranium were likewise discovered artificially, starting with Edwin McMillan and Philip Abelson's 1940 discovery of neptunium (via bombardment of uranium with neutrons). Glenn T. Seaborg and his team at the Lawrence Berkeley National Laboratory (LBNL) continued discovering transuranium elements, starting with plutonium in 1941, and discovered that contrary to previous thinking, the elements from actinium onwards were congeners of the lanthanides rather than transition metals. Bassett (1892), Werner (1905), and the French engineer Charles Janet (1928) had previously suggested this, but their ideas did not then receive general acceptance. Seaborg thus called them the actinides. Elements up to 101 (named mendelevium in honour of Mendeleev) were synthesized up to 1955, either through neutron or alpha-particle irradiation, or in nuclear explosions in the cases of 99 (einsteinium) and 100 (fermium).

A significant controversy arose with elements 102 through 106 in the 1960s and 1970s, as competition arose between the LBNL team (now led by Albert Ghiorso) and a team of Soviet scientists at the Joint Institute for Nuclear Research (JINR) led by Georgy Flyorov. Each team claimed discovery, and in some cases each proposed their own name for the element, creating an element naming controversy that lasted decades. These elements were made by bombardment of actinides with light ions. IUPAC at first adopted a hands-off approach, preferring to wait and see if a consensus would be forthcoming. But as it was also the height of the Cold War, it became clear that this would not happen. As such, IUPAC and the International Union of Pure and Applied Physics (IUPAP) created a Transfermium Working Group (TWG, fermium being element 100) in 1985 to set out criteria for discovery, which were published in 1991. After some further controversy, these elements received their final names in 1997, including seaborgium (106) in honour of Seaborg.

Yuri Oganessian

The TWG's criteria were used to arbitrate later element discovery claims from LBNL and JINR, as well as from research institutes in Germany (GSI) and Japan (Riken). Currently, consideration of discovery claims is performed by a IUPAC/IUPAP Joint Working Party. After priority was assigned, the elements were officially added to the periodic table, and the discoverers were invited to propose their names. By 2016, this had occurred for all elements up to 118, therefore completing the periodic table's first seven rows. The discoveries of elements beyond 106 were made possible by techniques devised by Yuri Oganessian at the JINR: cold fusion (bombardment of lead and bismuth by heavy ions) made possible the 1981–2004 discoveries of elements 107 through 112 at GSI and 113 at Riken, and he led the JINR team (in collaboration with American scientists) to discover elements 114 through 118 using hot fusion (bombardment of actinides by calcium ions) in 1998–2010. The heaviest known element, oganesson (118), is named in Oganessian's honour. Element 114 is named flerovium in honour of his predecessor and mentor Flyorov.

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". The discovery criteria set down by the TWG were updated in 2020 in response to experimental and theoretical progress that had not been foreseen in 1991. Today, the periodic table is among the most recognisable icons of chemistry. IUPAC is involved today with many processes relating to the periodic table: the recognition and naming of new elements, recommending group numbers and collective names, and the updating of atomic weights.

Future extension beyond the seventh period

Main article: Extended periodic table See also: Island of stability
Energy eigenvalues (in eV) for the outermost electrons of elements with Z = 100 through 172, predicted using Dirac–Fock calculations. The − and + signs refer to orbitals with decreased or increased azimuthal quantum number from spin–orbit splitting respectively: p− is p1/2, p+ is p3/2, d− is d3/2, d+ is d5/2, f− is f5/2, f+ is f7/2, g− is g7/2, and g+ is g9/2. The spacing of energy levels up to Z = 120 is normal, and becomes normal again at Z = 157; between them, a very different situation is observed.

The most recently named elements – nihonium (113), moscovium (115), tennessine (117), and oganesson (118) – completed the seventh row of the periodic table. Future elements would have to begin an eighth row. These elements may be referred to either by their atomic numbers (e.g. "element 164"), or by the IUPAC systematic element names adopted in 1978, which directly relate to the atomic numbers (e.g. "unhexquadium" for element 164, derived from Latin unus "one", Greek hexa "six", Latin quadra "four", and the traditional -ium suffix for metallic elements). All attempts to synthesize such elements have failed so far. An attempt to make element 119 has been ongoing since 2018 at the Riken research institute in Japan. The LBNL in the United States, the JINR in Russia, and the Heavy Ion Research Facility in Lanzhou (HIRFL) in China also plan to make their own attempts at synthesizing the first few period 8 elements.

If the eighth period followed the pattern set by the earlier periods, then it would contain fifty elements, filling the 8s, 5g, 6f, 7d, and finally 8p subshells in that order. But by this point, relativistic effects should result in significant deviations from the Madelung rule. Various different models have been suggested for the configurations of eighth-period elements, as well as how to show the results in a periodic table. All agree that the eighth period should begin like the previous ones with two 8s elements, 119 and 120. However, after that the massive energetic overlaps between the 5g, 6f, 7d, and 8p subshells means that they all begin to fill together, and it is not clear how to separate out specific 5g and 6f series. Elements 121 through 156 thus do not fit well as chemical analogues of any previous group in the earlier parts of the table, although they have sometimes been placed as 5g, 6f, and other series to formally reflect their electron configurations. Eric Scerri has raised the question of whether an extended periodic table should take into account the failure of the Madelung rule in this region, or if such exceptions should be ignored. The shell structure may also be fairly formal at this point: already the electron distribution in an oganesson atom is expected to be rather uniform, with no discernible shell structure.

The situation from elements 157 to 172 should return to normalcy and be more reminiscent of the earlier rows. The heavy p-shells are split by the spin–orbit interaction: one p-orbital (p1/2) is more stabilized, and the other two (p3/2) are destabilized. (Such shifts in the quantum numbers happen for all types of shells, but it makes the biggest difference to the order for the p-shells.) It is likely that by element 157, the filled 8s and 8p1/2 shells with four electrons in total have sunk into the core. Beyond the core, the next orbitals are 7d and 9s at similar energies, followed by 9p1/2 and 8p3/2 at similar energies, and then a large gap. Thus, the 9s and 9p1/2 orbitals in essence replace the 8s and 8p1/2 ones, making elements 157–172 probably chemically analogous to groups 3–18: for example, element 164 would appear two places below lead in group 14 under the usual pattern, but is calculated to be very analogous to palladium in group 10 instead. Thus, it takes fifty-four elements rather than fifty to reach the next noble element after 118. However, while these conclusions about elements 157 through 172's chemistry are generally agreed by models, there is disagreement on whether the periodic table should be drawn to reflect chemical analogies, or if it should reflect likely formal electron configurations, which should be quite different from earlier periods and are not agreed between sources. Discussion about the format of the eighth row thus continues.

Beyond element 172, calculation is complicated by the 1s electron energy level becoming imaginary. Such a situation does have a physical interpretation and does not in itself pose an electronic limit to the periodic table, but the correct way to incorporate such states into multi-electron calculations is still an open question needing to be solved to calculate the periodic table's structure beyond this point.

Nuclear stability will likely prove a decisive factor constraining the number of possible elements. It depends on the balance between the electric repulsion between protons and the strong force binding protons and neutrons together. Protons and neutrons are arranged in shells, just like electrons, and so a closed shell can significantly increase stability: the known superheavy nuclei exist because of such a shell closure, probably at around 114–126 protons and 184 neutrons. They are probably close to a predicted island of stability, where superheavy nuclides should be more long-lived than expected: predictions for the longest-lived nuclides on the island range from microseconds to millions of years. It should nonetheless be noted that these are essentially extrapolations into an unknown part of the chart of nuclides, and systematic model uncertainties need to be taken into account.

As the closed shells are passed, the stabilizing effect should vanish. Thus, superheavy nuclides with more than 184 neutrons are expected to have much shorter lifetimes, spontaneously fissioning within 10 seconds. If this is so, then it would not make sense to consider them chemical elements: IUPAC defines an element to exist only if the nucleus lives longer than 10 seconds, the time needed for it to gather an electron cloud. Nonetheless, theoretical estimates of half-lives are very model-dependent, ranging over many orders of magnitude. The extreme repulsion between protons is predicted to result in exotic nuclear topologies, with bubbles, rings, and tori expected: this further complicates extrapolation. It is not clear if any further-out shell closures exist, due to an expected smearing out of distinct nuclear shells (as is already expected for the electron shells at oganesson). Furthermore, even if later shell closures exist, it is not clear if they would allow such heavy elements to exist. As such, it may be that the periodic table practically ends around element 120, as elements become too short-lived to observe, and then too short-lived to have chemistry; the era of discovering new elements would thus be close to its end. If another proton shell closure beyond 126 does exist, then it probably occurs around 164; thus the region where periodicity fails more or less matches the region of instability between the shell closures.

Alternatively, quark matter may become stable at high mass numbers, in which the nucleus is composed of freely flowing up and down quarks instead of binding them into protons and neutrons; this would create a continent of stability instead of an island. Other effects may come into play: for example, in very heavy elements the 1s electrons are likely to spend a significant amount of time so close to the nucleus that they are actually inside it, which would make them vulnerable to electron capture.

Even if eighth-row elements can exist, producing them is likely to be difficult, and it should become even more difficult as atomic number rises. Although the 8s elements 119 and 120 are expected to be reachable with present means, the elements beyond that are expected to require new technology, if they can be produced at all. Experimentally characterizing these elements chemically would also pose a great challenge.

Alternative periodic tables

Main article: Types of periodic tables
Otto Theodor Benfey's spiral periodic table (1964)

The periodic law may be represented in multiple ways, of which the standard periodic table is only one. 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. Many forms retain the rectangular structure, including Charles Janet's left-step periodic table (pictured below), and the modernised form of Mendeleev's original 8-column layout that is still common in Russia. Other periodic table formats have been shaped much more exotically, such as spirals (Otto Theodor Benfey's pictured to the right), circles and triangles.

Alternative periodic tables are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables, with different ones skewed more towards emphasizing chemistry or physics at either end. The standard form, which remains by far the most common, is somewhere in the middle.

The many different forms of the periodic table have prompted the questions of whether there is an optimal or definitive form of the periodic table, and if so, what it might be. There are no current consensus answers to either question. Janet's left-step table is being increasingly discussed as a candidate for being the optimal or most fundamental form; Scerri has written in support of it, as it clarifies helium's nature as an s-block element, increases regularity by having all period lengths repeated, faithfully follows Madelung's rule by making each period correspond to one value of n + ℓ, and regularises atomic number triads and the first-row anomaly trend. While he notes that its placement of helium atop the alkaline earth metals can be seen a disadvantage from a chemical perspective, he counters this by appealing to the first-row anomaly, pointing out that the periodic table "fundamentally reduces to quantum mechanics", and that it is concerned with "abstract elements" and hence atomic properties rather than macroscopic properties.

Left-step periodic table (by Charles Janet)
f f f f f f f f f f f f f f d d d d d d d d d d p p p p p p s s
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
This form of periodic table is congruent with the order in which electron shells are ideally filled according to the Madelung rule, as shown in the accompanying sequence in the left margin (read from top to bottom, left to right). The experimentally determined ground-state electron configurations of the elements differ from the configurations predicted by the Madelung rule in twenty instances, but the Madelung-predicted configurations are always at least close to the ground state. The last two elements shown, elements 119 and 120, have not yet been synthesized.

See also

Notes

  1. ^ The question of how many natural elements there are is quite complicated and is not fully resolved. The heaviest element that occurs in large quantities on Earth is element 92, uranium. However, uranium can undergo spontaneous fission in nature, and the resulting neutrons can strike other uranium atoms. If neutron capture then occurs, elements 93 and 94, neptunium and plutonium, are formed via beta decay; these are in fact more common than some of the rarest elements in the first 92, such as promethium, astatine, and francium (see Abundance of elements in Earth's crust). Theoretically, neutron capture on the resulting plutonium might produce even higher-numbered elements, but the quantities would be too small to be observed. In the early Solar System, shorter-lived elements had not yet decayed away, and consequently there were more than 94 naturally occurring elements. Curium (element 96) is the longest-lived element beyond the first 94, and is probably still being brought to Earth via cosmic rays, but it has not been found. Elements up to 99 (einsteinium) have been observed in Przybylski's Star. Elements up to 100 (fermium) probably occurred in the natural nuclear fission reactor at Oklo Mine, Gabon, but they have long since decayed away. Even heavier elements may be produced in the r-process via supernovae or neutron star mergers, but this has not been confirmed. It is not clear how far they would extend past 100 and how long they would last: calculations suggest that nuclides of mass number around 280 to 290 are formed in the r-process, but quickly beta decay to nuclides that suffer spontaneous fission, so that 99.9% of the produced superheavy nuclides would decay within a month. If instead they were sufficiently long-lived, they might similarly be brought to Earth via cosmic rays, but again none have been found.
  2. Some isotopes currently considered stable are theoretically expected to be radioactive with extremely long half-lives: for instance, all the stable isotopes of elements 62 (samarium), 63 (europium), and all elements from 67 (holmium) onward are expected to undergo alpha decay or double beta decay. However, the predicted half-lives are extremely long (e.g. the alpha decay of Pb to the ground state of Hg is expected to have a half-life greater than 10 years), and the decays have never been observed.
  3. The half-life of plutonium's most stable isotope is just long enough that it should also be a primordial element. A 1971 study claimed to have detected primordial plutonium, but a more recent study from 2012 could not detect it. Based on its likely initial abundance in the Solar System, present experiments as of 2022 are likely about an order of magnitude away from detecting live primordial Pu.
  4. Tiny traces of plutonium are also continually brought to Earth via cosmic rays.
  5. See for example the periodic table poster sold by Sigma-Aldrich.
  6. Strictly speaking, one cannot draw an orbital such that the electron is guaranteed to be inside it, but it can be drawn to guarantee a 90% probability of this for example.
  7. ^ Authors differ on whether the n + ℓ rule has yet been derived from quantum mechanics. Scerri claims that it has not, despite several attempts to do so. On the other hand, Ostrovsky, who has claimed such justification from 1971, wrote 'Some authors insist that "still nobody has deduced the n+l rule from the principles of quantum mechanics", while others present quantum justification of the rule that was not ever disputed.' Other authors argue that such a derivation is not necessary, because it admits exceptions.
  8. Once two to four electrons are removed, the d and f orbitals usually become lower in energy than the s ones:
    1s ≪ 2s < 2p ≪ 3s < 3p ≪ 3d < 4s < 4p ≪ 4d < 5s < 5p ≪ 4f < 5d < 6s < 6p ≪ 5f < 6d < 7s < 7p ≪ ...
    and in the limit for extremely highly charged ions, orbitals simply fill in the order of increasing n instead. There is a gradual transition between the limiting situations of highly charged ions (increasing n) and neutral atoms (Madelung's rule). Thus for example, the energy order for the 55th electron outside the xenon core proceeds as follows in the isoelectronic series of caesium (55 electrons):
    Cs: 6s < 6p < 5d < 7s < 4f
    Ba: 6s < 5d < 6p < 7s < 4f
    La: 5d < 4f < 6s < 6p < 7s
    Ce: 4f < 5d < 6s < 6p < 7s
    and in the isoelectronic series of holmium (67 electrons), a Ho atom is 4f6s, but Er is 4f6s, Tm through W are 4f, and from Re onward the configuration is 4f5p following the hydrogenic order.
    Also, the ordering of the orbitals between each ≪ changes somewhat throughout each period. For example, the ordering in argon and potassium is 3p ≪ 4s < 4p ≪ 3d; by calcium it has become 3p ≪ 4s < 3d < 4p; from scandium to copper it is 3p ≪ 3d < 4s < 4p; and from zinc to krypton it is 3p < 3d ≪ 4s < 4p as the d-orbitals fall into the core at gallium. Deeply buried core shells in heavy atoms thus come closer to the hydrogenic order: around osmium (Z = 76) 4f falls below 5p, and around bismuth (Z = 83) 4f falls below 5s as well.
  9. In fact, electron configurations represent a first-order approximation: an atom really exists in a superposition of multiple configurations, and electrons in an atom are indistinguishable. The elements in the d- and f-blocks have multiple configurations separated by small energies and can change configuration depending on the chemical environment. In some of the undiscovered g-block elements, mixing of configurations may become so important that the result can no longer be well-described by a single configuration.
  10. Compounds that would use the 6d orbitals of nihonium as valence orbitals have been theoretically investigated, but they are all expected to be too unstable to observe.
  11. Properties of the p-block elements nevertheless do affect the succeeding s-block elements. The 3s shell in sodium is above a kainosymmetric 2p core, but the 4s shell in potassium is above the much larger 3p core. Hence while one would have already expected potassium atoms to be larger than sodium atoms, the size difference is greater than usual.
  12. There are many lower oxides as well: for example, phosphorus in group 15 forms two oxides, P2O3 and P2O5.
  13. The normally "forbidden" intermediate oxidation states may be stabilized by forming dimers, as in (gallium in the +2 oxidation state) or S2F10 (sulfur in the +5 oxidation state). Some compounds that appear to be in such intermediate oxidation states are actually mixed-valence compounds, such as Sb2O4, which contains both Sb(III) and Sb(V).
  14. The boundary between dispersion forces and metallic bonding is gradual, like that between ionic and covalent bonding. Characteristic metallic properties do not appear in small mercury clusters, but do appear in large ones.
  15. All this describes the situation at standard pressure. Under sufficiently high pressure, the band gaps of any solid drop to zero and metallization occurs. Thus for example at about 170 kbar iodine becomes a metal, and metallic hydrogen should form at pressures of about four million atmospheres. See metallization pressure for values for all nonmetals.
  16. Descriptions of the structures formed by the elements can be found throughout Greenwood and Earnshaw. There are two borderline cases. Arsenic's most stable form conducts electricity like a metal, but the bonding is significantly more localized to the nearest neighbours than it is for the similar structures of antimony and bismuth, and unlike normal metals it does not have a long liquid range, but rather sublimes instead. Hence its structure is better treated as network covalent. Carbon as graphite shows metallic conduction parallel to its planes, but is a semiconductor perpendicular to them. Some computations predict copernicium and flerovium to be nonmetallic, but the most recent experiments on them suggest that they are metallic. Astatine is calculated to metallise at standard conditions, so presumably tennessine should as well.
  17. See melting points of the elements (data page). The same is probably true of francium, but due to its extreme instability, this has never been experimentally confirmed. Copernicium and flerovium are expected to be liquids, similar to mercury, and experimental evidence suggests that they are metals.
  18. ^ See lists of metalloids. For example, a periodic table used by the American Chemical Society includes polonium as a metalloid, but one used by the Royal Society of Chemistry does not, and that included in the Encyclopædia Britannica does not refer to metalloids or semi-metals at all. Classification can change even within a single work. For example, Sherwin and Weston's Chemistry of the Non-Metallic Elements (1966) has a periodic table on p. 7 classifying antimony as a nonmetal, but on p. 115 it is called a metal.
  19. Demkov and Ostrovsky consider the potential U 1 / 2 ( r ) = − 2 v r R ( r + R ) 2 {\displaystyle U_{1/2}(r)=-{\frac {2v}{rR(r+R)^{2}}}} where R {\displaystyle R} and v {\displaystyle v} are constant parameters; this approaches a Coulomb potential for small r {\displaystyle r} . When v {\displaystyle v} satisfies the condition v = v N = 1 4 R 2 N ( N + 1 ) {\displaystyle v=v_{N}={\frac {1}{4}}R^{2}N(N+1)} , where N = n + l {\displaystyle N=n+l} , the zero-energy solutions to the Schrödinger equation for this potential can be described analytically with Gegenbauer polynomials. As v {\displaystyle v} passes through each of these values, a manifold containing all states with that value of N {\displaystyle N} arises at zero energy and then becomes bound, recovering the Madelung order. Perturbation-theory considerations show that states with smaller n {\displaystyle n} have lower energy, and that the s-orbitals (with l = 0 {\displaystyle l=0} ) have their energies approaching the next n + l {\displaystyle n+l} group.
  20. For example, the early actinides continue to behave more like the d-block transition metals in their propensity towards high oxidation states all the way from actinium to uranium, even though it is actually only actinium and thorium that have d-block-like configurations in the gas phase; f-electrons appear already at protactinium. Uranium's actual configuration of 5f6d7s is in fact analogous to that Hund assumed for the lanthanides, but uranium does not favour the trivalent state, preferring to be tetravalent or hexavalent. On the other hand, lanthanide-like configurations for the actinides begin at plutonium, but the shift towards lanthanide-like behaviour is only clear at curium: the elements between uranium and curium form a transition from transition-metal-like behaviour to lanthanide-like behaviour. Thus chemical behaviour and electron configuration do not exactly match each other.
  21. Technetium, promethium, astatine, neptunium, and plutonium were eventually discovered to occur in nature as well, albeit in tiny traces. See timeline of chemical element discoveries.

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Bibliography

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). Graphic 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). Archived (PDF) from the original on 11 August 2020. 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.

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1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18
1 H He
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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
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