The periodic table, also known as the periodic table of elements, arranges the chemical elements such as hydrogen, silicon, iron, and uranium according to their recurring properties. The number of each element corresponds to the number of protons in its nucleus (which is the same as the number of electrons orbiting that nucleus). The modern periodic table provides a useful framework for analyzing chemical reactions, and is widely used in chemistry, physics and other sciences.
The seven rows of the table, called periods, generally have metals on the left and nonmetals on the right. The columns, called groups, contain elements with similar chemical behaviours. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens; and group 18 are the noble gases. Also displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals. The organization of the periodic table can be used to derive relationships between the properties of the various elements, and to predict chemical properties and behaviours of undiscovered or newly synthesized elements.
Russian chemist Dmitri Mendeleev published the first recognizable periodic table in 1869, developed mainly to illustrate periodic trends of the then-known elements. He also predicted some properties of unidentified elements that were expected to fill gaps within the table. Most of his forecasts soon proved to be correct, culminating with the discovery of gallium and germanium in 1875 and 1886 respectively, which corroborated his predictions. Mendeleev's idea has been slowly expanded and refined with the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour.
The table here shows a widely used layout. Other forms (discussed below) show different structures in detail. Some discussion remains ongoing regarding the placement and categorisation of specific elements, the future extension and limits of the table, and whether there is an optimal form of the table.
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The number of each element—its atomic number—corresponds to the number of protons in its nucleus and the number of electrons orbiting that nucleus.
Sets of elements
- This section outlines metals and nonmetals (and metalloids); categories of elements; groups and periods; and periodic table blocks.
While the recognition of metals as solid, fusible and generally malleable substances dates from antiquity, Antoine Lavoisier may have the first to formally distinguish between metals and nonmetals ('non-métalliques') in 1789 with the publication of his 'revolutionary' Elementary Treatise on Chemistry. In 1811, Berzelius referred to nonmetallic elements as metalloids, in reference to their ability to form oxyanions. In 1825, in a revised German edition of his Textbook of Chemistry, he subdivided the metalloids into three classes. These were: constantly gaseous 'gazolyta' (hydrogen, nitrogen, oxygen); real metalloids (sulfur, phosphorus, carbon, boron, silicon); and salt-forming 'halogenia' (fluorine, chlorine, bromine, iodine). Only recently, since the mid-20th century, has the term metalloid been widely used to refer to elements with intermediate or borderline properties between metals and nonmetals. Mendeleev published his periodic table in 1869, along with references to groups of families of elements, and rows or periods of his periodic table. At the same time, Hinrichs wrote that simple lines could be drawn on a periodic table in order to delimit properties of interest, such as elements having metallic lustre (in contrast to those not having such lustre). Charles Janet, in 1928, appears to have been the first to refer to the periodic table's blocks.
Metals, metalloids and nonmetals
According to their shared physical and chemical properties, the elements can be classified into the major categories of metals, metalloids and nonmetals. Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals (other than noble gases). A majority of nonmetals are colored or colorless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding. In between metals and nonmetals are metalloids, which have intermediate or mixed properties.
Metal and nonmetals can be further classified into subcategories that show a gradation from metallic to non-metallic properties, when going left to right in the rows. The metals may be subdivided into the highly reactive alkali metals, through the less reactive alkaline earth metals, lanthanides and actinides, via the archetypal transition metals, and ending in the physically and chemically weak post-transition metals. Nonmetals may be simply subdivided into the polyatomic nonmetals, being nearer to the metalloids and show some incipient metallic character; the essentially nonmetallic diatomic nonmetals, nonmetallic and the almost completely inert, monatomic noble gases. Specialized groupings such as refractory metals and noble metals, are examples of subsets of transition metals, also known and occasionally denoted.
Placing elements into categories and subcategories based just on shared properties is imperfect. There is a large disparity of properties within each category with notable overlaps at the boundaries, as is the case with most classification schemes. Beryllium, for example, is classified as an alkaline earth metal although its amphoteric chemistry and tendency to mostly form covalent compounds are both attributes of a chemically weak or post-transition metal. Radon is classified as a nonmetallic noble gas yet has some cationic chemistry that is characteristic of metals. Other classification schemes are possible such as the division of the elements into mineralogical occurrence categories, or crystalline structures. Categorizing the elements in this fashion dates back to at least 1869 when Hinrichs wrote that simple boundary lines could be placed on the periodic table to show elements having shared properties, such as metals, nonmetals, or gaseous elements.
Set of elements that share some common properties are commonly put together in chemical categories. Some of these categories are better known than others; among the best-known ones are transition metals, noble gases, and halogens. Pnictogens make one category recognized by the most authoritative body of chemical nomenclature, the International Union of Pure and Applied Chemistry (IUPAC), but this name is not very common in the literature; in contrast, it avoids using the term metalloid, which is rather well-known in the literature. It is common to use a subset of the range of categories to colour-code elements on the periodic table.
The underlying rationale for common behavior across a category can usually be explained by the position of these elements in the periodic table: for example, noble gases, well known for their chemical inertness, are all in the rightmost column, meaning the have complete electron shells and thus very unwilling to participate in chemical reactions, whereas halogens, which are known as very reactive elements and located just to the left of noble gases, lack one electron to attain such a configuration and thus are very likely to attract one. For this reason, many categories match groups in the periodic table, though there are exceptions. Categories may overlap, and their names need not necessarily reflect their shared properties; for example, the rare earths are not particularly rare.
Different authors may use different categories depending on the properties of interest. Additionally, different authors may disagree on which elements belong to which categories, particularly around the boundaries. The approximate correspondence between groups and similar chemical properties may break down for some of the heaviest elements due to strong relativistic effects, and although it is common to extend the categories to the whole group regardless, some questions have been raised about this practice.
A group or family is a vertical column in the periodic table. Groups usually have more significant periodic trends than periods and blocks, explained below. Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group generally have the same electron configurations in their valence shell. Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number. In some parts of the periodic table, such as the d-block and the f-block, horizontal similarities can be as important as, or more pronounced than, vertical similarities.
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). Previously, they were known by roman numerals. In America, the roman numerals were followed by either an "A" if the group was in the s- or p-block, or a "B" if the group was in the d-block. The roman numerals used correspond to the last digit of today's naming convention (e.g. the group 4 elements were group IVB, and the group 14 elements were group IVA). In Europe, the lettering was similar, except that "A" was used if the group was before group 10, and "B" was used for groups including and after group 10. In addition, groups 8, 9 and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII. In 1988, the new IUPAC naming system was put into use, and the old group names were deprecated.
Some of these groups have been given trivial (unsystematic) names, as seen in the table below, although some are rarely used. Groups 3–10 have no trivial names and are referred to simply by their group numbers or by the name of the first member of their group (such as "the scandium group" for group 3), since they display fewer similarities and/or vertical trends.
Elements in the same group tend to show patterns in atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group has a top-to-bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus. There are exceptions to these trends: for example, in group 11, electronegativity increases farther down the group.
A period is a horizontal row in the periodic table. Although groups generally have more significant periodic trends, there are regions where horizontal trends are more significant than vertical group trends, such as the f-block, where the lanthanides and actinides form two substantial horizontal series of elements.
Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron, which causes the electron to be drawn closer to the nucleus. 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. 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.
Specific regions of the periodic table can be referred to as blocks in recognition of the sequence in which the electron shells of the elements are filled. Elements are assigned to blocks by what orbitals their valence electrons or vacancies lie in. The s-block comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. The p-block comprises the last six groups, which are groups 13 to 18 in IUPAC group numbering (3A to 8A in American group numbering) and contains, among other elements, all of the metalloids. The d-block comprises groups 3 to 12 (or 3B to 2B in American group numbering) and contains all of the transition metals. The f-block, often offset below the rest of the periodic table, has no group numbers and comprises most of the lanthanides and actinides. A hypothetical g-block is expected to begin around element 121, a few elements away from what is currently known.
Periodic trends and patterns
The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of electron shells (numbered 1, 2, and so on). Each shell consists of one or more subshells (named s, p, d, f and g). As atomic number increases, electrons progressively fill these shells and subshells more or less according to the Madelung rule or energy ordering rule, as shown in the diagram. The electron configuration for neon, for example, is 1s2 2s2 2p6. With an atomic number of ten, neon has two electrons in the first shell, and eight electrons in the second shell; there are two electrons in the s subshell and six in the p subshell. In periodic table terms, the first time an electron occupies a new shell corresponds to the start of each new period, these positions being occupied by hydrogen and the alkali metals.
Since the properties of an element are mostly determined by its electron configuration, the properties of the elements likewise show recurring patterns or periodic behaviour, some examples of which are shown in the diagrams below for atomic radii, ionization energy and electron affinity. It is this periodicity of properties, manifestations of which were noticed well before the underlying theory was developed, that led to the establishment of the periodic law (the properties of the elements recur at varying intervals) and the formulation of the first periodic tables. The periodic law may then be successively clarified as: depending on atomic weight; depending on atomic number; and depending on the total number of s, p, d, and f electrons in each atom. The cycles last 2, 6, 10, and 14 elements respectively.
There is additionally an internal "double periodicity" that splits the shells in half; this arises because the first half of the electrons going into a particular type of subshell fill unoccupied orbitals, but the second half have to fill already occupied orbitals, following Hund's rule of maximum multiplicity. The second half thus suffer additional repulsion that causes the trend to split between first-half and second-half elements; this is for example evident when observing the ionisation energies of the 2p elements, in which the triads B-C-N and O-F-Ne show increases, but oxygen actually has a first ionisation slightly lower than that of nitrogen as it is easier to remove the extra, paired electron.
Atomic radii vary in a predictable and explainable manner across the periodic table. For instance, the radii generally decrease along each period of the table, from the alkali metals to the noble gases; and increase down each group. The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory.
The electrons in the 4f-subshell, which is progressively filled from lanthanum (element 57) to ytterbium (element 70),[n 2] are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii that are smaller than would be expected and that are almost identical to the atomic radii of the elements immediately above them. Hence lutetium has virtually the same atomic radius (and chemistry) as yttrium, hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. This is an effect of the lanthanide contraction: a similar actinide contraction also exists. The effect of the lanthanide contraction is noticeable up to platinum (element 78), after which it is masked by a relativistic effect known as the inert pair effect. The d-block contraction, which is a similar effect between the d-block and p-block, is less pronounced than the lanthanide contraction but arises from a similar cause.
Such contractions exist throughout the table, but are chemically most relevant for the lanthanides with their almost constant +3 oxidation state.
The first ionization energy is the energy it takes to remove one electron from an atom, the second ionization energy is the energy it takes to remove a second electron from the atom, and so on. For a given atom, successive ionization energies increase with the degree of ionization. For magnesium as an example, the first ionization energy is 738 kJ/mol and the second is 1450 kJ/mol. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table.
Large jumps in the successive molar ionization energies occur when removing an electron from a noble gas (complete electron shell) configuration. For magnesium again, the first two molar ionization energies of magnesium given above correspond to removing the two 3s electrons, and the third ionization energy is a much larger 7730 kJ/mol, for the removal of a 2p electron from the very stable neon-like configuration of Mg2+. Similar jumps occur in the ionization energies of other third-row atoms.
Electronegativity is the tendency of an atom to attract a shared pair of electrons. An atom's electronegativity is affected by both its atomic number and the distance between the valence electrons and the nucleus. The higher its electronegativity, the more an element attracts electrons. It was first proposed by Linus Pauling in 1932. In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is the most electronegative of the elements,[n 3] while caesium is the least, at least of those elements for which substantial data is available.
There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon respectively because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity. The anomalously high electronegativity of lead, particularly when compared to thallium and bismuth, is an artifact of electronegativity varying with oxidation state: its electronegativity conforms better to trends if it is quoted for the +2 state instead of the +4 state.
The electron affinity of an atom is the amount of energy released when an electron is added to a neutral atom to form a negative ion. Although electron affinity varies greatly, some patterns emerge. Generally, nonmetals have more positive electron affinity values than metals. Chlorine most strongly attracts an extra electron. The electron affinities of the noble gases have not been measured conclusively, so they may or may not have slightly negative values.
Electron affinity generally increases across a period. This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and is therefore more stable.
A trend of decreasing electron affinity going down groups would be expected. The additional electron will be entering an orbital farther away from the nucleus. As such this electron would be less attracted to the nucleus and would release less energy when added. In going down a group, around one-third of elements are anomalous, with heavier elements having higher electron affinities than their next lighter congenors. Largely, this is due to the poor shielding by d and f electrons. A uniform decrease in electron affinity only applies to group 1 atoms.
The lower the values of ionization energy, electronegativity and electron affinity, the more metallic character the element has. Conversely, nonmetallic character increases with higher values of these properties. Given the periodic trends of these three properties, metallic character tends to decrease going across a period (or row) and, with some irregularities (mostly) due to poor screening of the nucleus by d and f electrons, and relativistic effects, tends to increase going down a group (or column or family). Thus, the most metallic elements (such as caesium) are found at the bottom left of traditional periodic tables and the most nonmetallic elements (such as neon) at the top right. The combination of horizontal and vertical trends in metallic character explains the stair-shaped dividing line between metals and nonmetals found on some periodic tables, and the practice of sometimes categorizing several elements adjacent to that line, or elements adjacent to those elements, as metalloids.
With some minor exceptions, oxidation numbers among the elements show four main trends according to their periodic table geographic location: left; middle; right; and south. On the left (groups 1 to 4, not including the f-block elements, and also niobium, tantalum, and probably dubnium in group 5), the highest most stable oxidation number is the group number, with lower oxidation states being less stable. In the middle (groups 3 to 11), higher oxidation states become more stable going down each group. Group 12 is an exception to this trend; they behave as if they were located on the left side of the table. On the right, higher oxidation states tend to become less stable going down a group. The shift between these trends is continuous: for example, group 3 also has lower oxidation states most stable in its lightest member (scandium, with CsScCl3 for example known in the +2 state), and group 12 is predicted to have copernicium more readily showing oxidation states above +2.
The lanthanides positioned along the south of the table are distinguished by having the +3 oxidation state in common; this is their most stable state. The early actinides show a pattern of oxidation states somewhat similar to those of their period 6 and 7 transition metal congeners; the later actinides are more similar to the lanthanides, though the last ones (excluding lawrencium) have an increasingly important +2 oxidation state that becomes the most stable state for nobelium.
Linking or bridging groups
From left to right across the four blocks of the long- or 32-column form of the periodic table are a series of linking or bridging groups of elements, located approximately between each block. In general, groups at the peripheries of blocks display similarities to the groups of the neighbouring blocks as well as to the other groups in their own blocks, as expected as most periodic trends are continuous. These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, lanthanides, and heavy group 4 and 5 elements show some behaviour similar to the alkaline earth metals or, more generally, s block metals but have some of the physical properties of d block transition metals. In fact, the metals all the way up to group 6 are united by being class-A cations ("hard" acids) that form more stable complexes with ligands whose donor atoms are the most electronegative nonmetals nitrogen, oxygen, and fluorine; metals later in the table form a transition to class-B cations ("soft" acids) that form more stable complexes with ligands whose donor atoms are the less electronegative heavier elements of groups 15 through 17.
Meanwhile, lutetium behaves chemically as a lanthanide (with which it is often classified) but shows a mix of lanthanide and transition metal physical properties (as does yttrium). Lawrencium, as an analogue of lutetium, would presumably display like characteristics.[n 4] The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals. The volatile group 12 metals, zinc, cadmium and mercury are sometimes regarded as linking the d block to the p block. Notionally they are d block elements but they have few transition metal properties and are more like their p block neighbors in group 13. The relatively inert noble gases, in group 18, bridge the most reactive groups of elements in the periodic table—the halogens in group 17 and the alkali metals in group 1.
The 1s, 2p, 3d, 4f, and 5g shells are each the first to have their value of ℓ, the azimuthal quantum number that determines a subshell's orbital angular momentum. This gives them some special properties, that has been referred to as kainosymmetry (from Greek καινός "new"). Elements filling these orbitals are usually less metallic than their heavier homologues, prefer lower oxidation states, and have smaller atomic and ionic radii. As kainosymmetric orbitals appear in the even rows (except for 1s), this creates an even–odd difference between periods from period 2 onwards: elements in even periods are smaller and have more oxidising higher oxidation states (if they exist), whereas elements in odd periods differ in the opposite direction.
First systemization attempts
In 1789, Antoine Lavoisier published a list of 33 chemical elements, grouping them into gases, metals, nonmetals, and earths. Chemists spent the following century searching for a more precise classification scheme. In 1829, Johann Wolfgang Döbereiner observed that many of the elements could be grouped into triads based on their chemical properties. Lithium, sodium, and potassium, for example, were grouped together in a triad as soft, reactive metals. Döbereiner also observed that, when arranged by atomic weight, the second member of each triad was roughly the average of the first and the third. This became known as the Law of Triads. German chemist Leopold Gmelin worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. Jean-Baptiste Dumas published work in 1857 describing relationships between various groups of metals. Although various chemists were able to identify relationships between small groups of elements, they had yet to build one scheme that encompassed them all. In 1857, German chemist August Kekulé observed that carbon often has four other atoms bonded to it. Methane, for example, has one carbon atom and four hydrogen atoms. This concept eventually became known as valency, where different elements bond with different numbers of atoms.
In 1862, the French geologist Alexandre-Émile Béguyer de Chancourtois published an early form of the periodic table, which he called the telluric helix or screw. He was the first person to notice the periodicity of the elements. With the elements arranged in a spiral on a cylinder by order of increasing atomic weight, de Chancourtois showed that elements with similar properties seemed to occur at regular intervals. His chart included some ions and compounds in addition to elements. His paper also used geological rather than chemical terms and did not include a diagram. As a result, it received little attention until the work of Dmitri Mendeleev.
In 1864, Julius Lothar Meyer, a German chemist, published a table with 28 elements. Realizing that an arrangement according to atomic weight did not exactly fit the observed periodicity in chemical properties he gave valency priority over minor differences in atomic weight. A missing element between Si and Sn was predicted with atomic weight 73 and valency 4. Concurrently, English chemist William Odling published an arrangement of 57 elements, ordered on the basis of their atomic weights. With some irregularities and gaps, he noticed what appeared to be a periodicity of atomic weights among the elements and that this accorded with "their usually received groupings". Odling alluded to the idea of a periodic law but did not pursue it. He subsequently proposed (in 1870) a valence-based classification of the elements.
English chemist John Newlands produced a series of papers from 1863 to 1866 noting that when the elements were listed in order of increasing atomic weight, similar physical and chemical properties recurred at intervals of eight. He likened such periodicity to the octaves of music. This so termed Law of Octaves was ridiculed by Newlands' contemporaries, and the Chemical Society refused to publish his work. Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as germanium. The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.
In 1867, Gustavus Hinrichs, a Danish born academic chemist based in America, published a spiral periodic system based on atomic spectra and weights, and chemical similarities. His work was regarded as idiosyncratic, ostentatious and labyrinthine and this may have militated against its recognition and acceptance.
Russian chemistry professor Dmitri Mendeleev and German chemist Julius Lothar Meyer independently published their periodic tables in 1869 and 1870, respectively. Mendeleev's table, dated March 1 [O.S. February 17] 1869, was his first published version. That of Meyer was an expanded version of his (Meyer's) table of 1864. They both constructed their tables by listing the elements in rows or columns in order of atomic weight and starting a new row or column when the characteristics of the elements began to repeat.
The recognition and acceptance afforded to Mendeleev's table came from two decisions he made. The first was to leave gaps in the table when it seemed that the corresponding element had not yet been discovered. Mendeleev was not the first chemist to do so, but he was the first to be recognized as using the trends in his periodic table to predict the properties of those missing elements, such as gallium and germanium. The second decision was to occasionally ignore the order suggested by the atomic weights and switch adjacent elements, such as tellurium and iodine, to better classify them into chemical families.
Mendeleev published in 1869, using atomic weight to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough to allow Mendeleev to accurately predict the properties of missing elements.
Mendeleev took the unusual step of naming missing elements using the Sanskrit numerals eka (1), dvi (2), and tri (3) to indicate that the element in question was one, two, or three rows removed from a lighter congener. It has been suggested that Mendeleev, in doing so, was paying homage to ancient Sanskrit grammarians, in particular Pāṇini, who devised a periodic alphabet for the language.
Following the discovery of the atomic nucleus by Ernest Rutherford in 1911, it was proposed that the integer count of the nuclear charge is identical to the sequential place of each element in the periodic table. In 1913, English physicist Henry Moseley using X-ray spectroscopy confirmed this proposal experimentally. Moseley determined the value of the nuclear charge of each element and showed that Mendeleev's ordering actually places the elements in sequential order by nuclear charge. Nuclear charge is identical to proton count and determines the value of the atomic number (Z) of each element. Using atomic number gives a definitive, integer-based sequence for the elements. Moseley predicted, in 1913, that the only elements still missing between aluminium (Z = 13) and gold (Z = 79) were Z = 43, 61, 72, and 75, all of which were later discovered. The atomic number is the absolute definition of an element and gives a factual basis for the ordering of the periodic table.
Second version and further development
In 1871, Mendeleev published his periodic table in a new form, with groups of similar elements arranged in columns rather than in rows, and those columns numbered I to VIII corresponding with the element's oxidation state. He also gave detailed predictions for the properties of elements he had earlier noted were missing, but should exist. These gaps were subsequently filled as chemists discovered additional naturally occurring elements. It is often stated that the last naturally occurring element to be discovered was francium (referred to by Mendeleev as eka-caesium) in 1939, but it was technically only the last element to be discovered in nature as opposed to by synthesis. Plutonium, produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971.
The popular periodic table layout, also known as the common or standard form (as shown at various other points in this article), is attributable to Horace Groves Deming. In 1923, Deming, an American chemist, published short (Mendeleev style) and medium (18-column) form periodic tables.[n 5] Merck and Company prepared a handout form of Deming's 18-column medium table, in 1928, which was widely circulated in American schools. By the 1930s Deming's table was appearing in handbooks and encyclopedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company.
With the development of modern quantum mechanical theories of electron configurations within atoms, it became apparent that each period (row) in the table corresponded to the filling of a quantum shell of electrons. Larger atoms have more electron sub-shells, so later tables have required progressively longer periods.
In 1945, Glenn Seaborg, an American scientist, made the suggestion that the actinide elements, like the lanthanides, were filling an f sub-level. Before this time the actinides were thought to be forming a fourth d-block row. Seaborg's colleagues advised him not to publish such a radical suggestion as it would most likely ruin his career. As Seaborg considered he did not then have a career to bring into disrepute, he published anyway. Seaborg's suggestion was found to be correct and he subsequently went on to win the 1951 Nobel Prize in chemistry for his work in synthesizing actinide elements.[n 6]
Although minute quantities of some transuranic elements occur naturally, they were all first discovered in laboratories. Their production has expanded the periodic table significantly, the first of these being neptunium, synthesized in 1939. Because many of the transuranic elements are highly unstable and decay quickly, they are challenging to detect and characterize when produced. There have been controversies concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights. In 2010, a joint Russia–US collaboration at Dubna, Moscow Oblast, Russia, claimed to have synthesized six atoms of tennessine (element 117), making it the most recently claimed discovery. It, along with nihonium (element 113), moscovium (element 115), and oganesson (element 118), are the four most recently named elements, whose names all became official on 28 November 2016.
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".
Different periodic tables
The long- or 32-column table
The modern periodic table is sometimes expanded into its long or 32-column form by reinstating the footnoted f-block elements into their natural position between the s- and d-blocks, as proposed by Alfred Werner in 1905. Unlike the 18-column form, this arrangement results in "no interruptions in the sequence of increasing atomic numbers". The relationship of the f-block to the other blocks of the periodic table also becomes easier to see. William B. Jensen advocates a form of table with 32 columns on the grounds that the lanthanides and actinides are otherwise relegated in the minds of students as dull, unimportant elements that can be quarantined and ignored. Despite these advantages, the 32-column form is generally avoided by editors on account of its undue rectangular ratio compared to a book page ratio, and the familiarity of chemists with the modern form, as introduced by Seaborg.
Placement of hydrogen and helium
Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium (1s22s1) and beryllium (1s22s2). Such a placement is common for hydrogen, as its chemistry has some similarities to the other group 1 elements: like them, hydrogen is univalent. But there are also some significant differences: for example, hydrogen is a nonmetal, unlike the other group 1 elements that are all metals. For this reason hydrogen has sometimes been placed instead in group 17, given hydrogen's strictly univalent and largely non-metallic chemistry, and the strictly univalent and non-metallic chemistry of fluorine (the element otherwise at the top of group 17). Sometimes, to show hydrogen has properties corresponding to both those of the alkali metals and the halogens, it is shown at the top of the two columns simultaneously. Finally, hydrogen is sometimes placed separately from any group; this is based on its general properties being regarded as sufficiently different from those of the elements in any other group.
Helium's extraordinary inertness is extremely close to that of the other light noble gases neon and argon in group 18, and not at all close to the behaviour of the metallic and increasingly reactive group 2 elements, and therefore it is nearly universally placed in group 18. That said, helium is occasionally placed separately from any group as well, and there are even a few chemists who have argued for helium in group 2 on the grounds of various properties such as ionisation energies and reactivity where helium fits better into the group 2 trend than the group 18 trend.
Group 3 and its elements in periods 6 and 7
La and Ac below Y
Lu and Lr below Y
Markers below Y
Although scandium and yttrium are always the first two elements in group 3, the identity of the next two elements is not completely settled. They are commonly lanthanum and actinium, and less often lutetium and lawrencium. The two variants originate from historical difficulties in placing the lanthanides in the periodic table, and arguments as to where the f block elements start and end.[n 7] A third (compromise) variant shows the two positions below yttrium as being occupied by all lanthanides and all actinides.
The lanthanum-actinium option[n 8] is the most common one. It results in a group 3 that has all elements ionise to a noble-gas electron configuration and smooth vertical periodic trends. The lutetium-lawrencium option[n 9] results in a contiguous d-block, and the kink in the vertical periodic trends at lutetium matches those of other early d-block groups. The lanthanides-actinides option[n 10] emphasises chemical similarity between lanthanides (although actinides are not quite as similar).
Most working chemists are not aware there is any controversy, even though the matter has been debated periodically for decades without apparent resolution. IUPAC has not yet made a recommendation on the matter; in 2015, an IUPAC taskforce was established to provide one.
Further periodic table extensions
Currently, the periodic table has seven complete rows, with all spaces filled in with discovered elements. Future elements would have to begin an eighth row. As atomic nuclei get highly charged, special relativity becomes needed to gauge the effect of the nucleus on the electron cloud. This results in heavy elements increasingly having differing properties compared to their lighter homologues in the periodic table, which is already visible in the late sixth and early seventh period, and expected to become very strong in the late seventh and eighth periods. Therefore, there are some discussions if this future eighth period should follow the pattern set by the earlier periods or not. Heavier elements also become increasingly unstable as the strong force that binds the nucleus together becomes less able to counteract repulsion between the positively-charged protons in it, so it is also an open question how many of the eighth-period elements will be able to exist. 
Tables with different structures
Within 100 years of the appearance of Mendeleev's table in 1869, Edward G. Mazurs had collected an estimated 700 different published versions of the periodic table. As well as numerous rectangular variations, other periodic table formats have been shaped, for example,[n 11] like a circle, cube, cylinder, building, spiral, lemniscate, octagonal prism, pyramid, sphere, or triangle. Such alternatives are often developed to highlight or emphasize chemical or physical properties of the elements that are not as apparent in traditional periodic tables.
A popular alternative structure is that of Otto Theodor Benfey (1960). The elements are arranged in a continuous spiral, with hydrogen at the centre and the transition metals, lanthanides, and actinides occupying peninsulas.
Most periodic tables are two-dimensional; three-dimensional tables are known to as far back as at least 1862 (pre-dating Mendeleev's two-dimensional table of 1869). More recent examples include Courtines' Periodic Classification (1925), Wringley's Lamina System (1949), Giguère's Periodic helix (1965) and Dufour's Periodic Tree (1996). Going one further, Stowe's Physicist's Periodic Table (1989) has been described as being four-dimensional (having three spatial dimensions and one colour dimension).
The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum. Towards the chemistry end of the continuum can be found, as an example, Rayner-Canham's "unruly" Inorganic Chemist's Periodic Table (2002), which emphasizes trends and patterns, and unusual chemical relationships and properties. Near the physics end of the continuum is Janet's Left-Step Periodic Table (1928). This has a structure that shows a closer connection to the order of electron-shell filling and, by association, quantum mechanics. A somewhat similar approach has been taken by Alper, albeit criticized by Eric Scerri as disregarding the need to display chemical and physical periodicity. Somewhere in the middle of the continuum is the ubiquitous common or standard form of periodic table. This is regarded as better expressing empirical trends in physical state, electrical and thermal conductivity, and oxidation numbers, and other properties easily inferred from traditional techniques of the chemical laboratory. Its popularity is thought to be a result of this layout having a good balance of features in terms of ease of construction and size, and its depiction of atomic order and periodic trends.
The many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table, to which there is currently not a consensus answer.
- List of chemical elements
- List of periodic table-related articles
- Names for sets of chemical elements
- Standard model
- Abundance of the chemical elements
- Atomic electron configuration table
- Element collecting
- Table of nuclides
- The Mystery of Matter: Search for the Elements (PBS film)
- Timeline of chemical element discoveries
- No data was available for the noble gases, astatine, francium and elements heavier than americium.
- Although lanthanum does not have a 4f electron in the ground state, lanthanum metal shows 4f occupancy and it can show 4f character in its compounds.
- While fluorine is the most electronegative of the elements under the Pauling scale, neon is the most electronegative element under other scales, such as the Allen scale.
- While Lr is thought to have a p rather than d electron in its ground-state electron configuration, and would therefore be expected to be a volatile metal capable of forming a +1 cation in solution like thallium, no evidence of either of these properties has been able to be obtained despite experimental attempts to do so. It was originally expected to have a d electron in its electron configuration and this may still be the case for metallic lawrencium, whereas gas phase atomic lawrencium is very likely thought to have a p electron.
- An antecedent of Deming's 18-column table may be seen in Adams' 16-column Periodic Table of 1911. Adams omits the rare earths and the "radioactive elements" (i.e. the actinides) from the main body of his table and instead shows them as being "careted in only to save space" (rare earths between Ba and eka-Yt; radioactive elements between eka-Te and eka-I). See: Elliot Q. A. (1911). "A modification of the periodic table". Journal of the American Chemical Society. 33(5): 684–88 .
- A second extra-long periodic table row, to accommodate known and undiscovered elements with an atomic weight greater than bismuth (thorium, protactinium and uranium, for example), had been postulated as far back as 1892. Most investigators considered that these elements were analogues of the third series transition elements, hafnium, tantalum and tungsten. The existence of a second inner transition series, in the form of the actinides, was not accepted until similarities with the electron structures of the lanthanides had been established. See: van Spronsen, J. W. (1969). The periodic system of chemical elements. Amsterdam: Elsevier. pp. 315–16, ISBN 0-444-40776-6.
- The detachment of the lanthanides from the main body of the periodic table has been attributed to the Czech chemist Bohuslav Brauner who, in 1902, allocated all of them ("Ce etc.") to one position in group 4, below zirconium. This arrangement was referred to as the "asteroid hypothesis", in analogy to asteroids occupying a single orbit in the solar system. Before this time the lanthanides were generally (and unsuccessfully) placed throughout groups I to VIII of the older 8-column form of periodic table. Although predecessors of Brauner's 1902 arrangement are recorded from as early as 1895, he is known to have referred to the "chemistry of asteroids" in an 1881 letter to Mendeleev. Other authors assigned all of the lanthanides to either group 3, groups 3 and 4, or groups 2, 3 and 4. In 1922 Niels Bohr continued the detachment process by locating the lanthanides between the s- and d-blocks. In 1949 Glenn T. Seaborg (re)introduced the form of periodic table that is popular today, in which the lanthanides and actinides appear as footnotes. Seaborg first published his table in a classified report dated 1944. It was published again by him in 1945 in Chemical and Engineering News, and in the years up to 1949 several authors commented on, and generally agreed with, Seaborg's proposal. In that year he noted that the best method for presenting the actinides seemed to be by positioning them below, and as analogues of, the lanthanides. See: Thyssen P. and Binnemans K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 41. Amsterdam: Elsevier, pp. 1–94; Seaborg G. T. (1994). Origin of the Actinide Concept'. In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 18. Amsterdam: Elsevier, pp. 1–27.
- For examples of this table see Atkins et al. (2006). Shriver & Atkins Inorganic Chemistry (4th ed.). Oxford: Oxford University Press • Myers et al. (2004). Holt Chemistry. Orlando: Holt, Rinehart & Winston • Chang R. (2000). Essential Chemistry (2nd ed.). Boston: McGraw-Hill
- For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G. & Overton T. (2013). Descriptive Inorganic Chemistry (6th ed.). New York: W. H. Freeman and Company • Brown et al. (2009). Chemistry: The Central Science (11th ed.). Upper Saddle River, New Jersey: Pearson Education • Moore et al. (1978). Chemistry. Tokyo: McGraw-Hill Kogakusha
- For examples of the group 3 = Ln and An table see Housecroft C. E. & Sharpe A. G. (2008). Inorganic Chemistry (3rd ed.). Harlow: Pearson Education • Halliday et al. (2005). Fundamentals of Physics (7th ed.). Hoboken, New Jersey: John Wiley & Sons • Nebergall et al. (1980). General Chemistry (6th ed.). Lexington: D. C. Heath and Company
- See The Internet database of periodic tables for depictions of these kinds of variants.
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- Dynamic periodic table, with interactive layouts
- Eric Scerri, leading philosopher of science specializing in the history and philosophy of the periodic table
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- Periodic table of endangered elements
- Periodic table of samples
- Periodic table of videos