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{{About|the table used in chemistry|other uses|Periodic table (disambiguation)}}

Periodic table (colour legend below)

{{sidebar periodic table}}

The periodic table is a tabular arrangement of the 化學元素s, ordered by their 原子序数 (number of 質子s), 电子排布s, and recurring 化学性质. This ordering shows {{tsl|en|periodic trends||}}, such as elements with similar behaviour in the same column. It also shows four rectangular 元素分区s with some approximately similar chemical properties. In general, within one row (period) the elements are 金属s on the left, and 非金属元素 on the right.

The rows of the table are called 元素周期; the columns are called 族 (化学). Six groups have names as well as numbers: for example, group 17 elements are the 卤素s; and group 18, the 稀有气体es. The periodic table can be used to derive relationships between the properties of the elements, and predict the properties of new elements yet to be discovered or synthesized. The periodic table provides a useful framework for analyzing chemical behaviour, and is widely used in chemistry and other sciences.

The Russian chemist 德米特里·伊万诺维奇·门捷列夫 published the first widely recognized periodic table in 1869. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of {{tsl|en|Dmitri Mendeleev's predicted elements||then-unknown elements}} that would be expected to fill gaps in this table. Most of his predictions were proved correct when the elements in question were subsequently discovered. Mendeleev's periodic table has since been expanded and refined with 化學元素發現年表 and the development of new theoretical models to explain chemical behaviour.

All elements from atomic numbers 1 () to 118 (Og) have been discovered or synthesized, with the most recent additions (, , Ts, and Og) being confirmed by the 國際純化學和應用化學聯合會 (IUPAC) in 2015 and officially named in 2016: they complete the first seven rows of the periodic table.[1][2] The first 94 elements exist naturally, although some are found only in trace amounts and were synthesized in laboratories before being found in nature.{{#tag:ref|The elements discovered initially by synthesis and later in nature are technetium (Z=43), promethium (61), astatine (85), neptunium (93), and plutonium (94).|group=n}} Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors.[3] Synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic 放射性同位素s of naturally occurring elements have also been produced in laboratories.


{{for|large cell versions|Periodic table (large cells)}} {{Periodic table}}

Each chemical element has a unique atomic number (Z) representing the number of protons in its nucleus.{{#tag:ref|An 0號元素 (i.e. a substance composed purely of neutrons), is included in a few alternate presentations, for example, in the 化學星空.|group=n}} Most elements have differing numbers of 中子s among different atoms, with these variants being referred to as 同位素s. For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and a very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element. Elements with no stable isotopes have the atomic masses of their most stable isotopes, where such masses are shown, listed in parentheses.[4]

In the standard periodic table, the elements are listed in order of increasing atomic number (the number of 質子s in the 原子核 of an atom). A new row (元素周期) is started when a new electron shell has its first electron. Columns (族 (化学)) are determined by the 电子排布 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 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.[5]

As of 2016, the periodic table has 118 confirmed elements, from element 1 (hydrogen) to 118 (oganesson). Elements 113, 115, 117 and 118 were officially confirmed by the 國際純化學和應用化學聯合會 (IUPAC) in December 2015. Their proposed names, nihonium (Nh), moscovium (Mc), tennessine (Ts) and oganesson (Og) respectively, were announced by the IUPAC in June 2016 and made official in November 2016.[6][7][8][9]

The first 94 elements occur naturally; the remaining 24, americium to oganesson (95–118) occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are {{tsl|en|primordial element||primordial}} and 11 occur only in decay chains of primordial elements.[3] No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form, nor has astatine (element 85); francium (element 87) has been only photographed in the form of light emitted from microscopic quantities (300,000 atoms).[10]

Grouping methods[编辑]


{{Main article|Group (periodic table)}} 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 電子層.[11] Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number.[12] However, 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.[13][14][15]

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).[16] Previously, they were known by 罗马数字. In America, the roman numerals were followed by either an "A" if the group was in the 元素分区 or 元素分区, or a "B" if the group was in the 元素分区. The roman numerals used correspond to the last digit of today's naming convention (e.g. the 4族元素s were group IVB, and the 碳族元素s were group IVA). In Europe, the lettering was similar, except that "A" was used if the group was before 10族元素, 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.[17]

Some of these groups have been given 化学序, 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 3族元素), since they display fewer similarities and/or vertical trends.[16]

Elements in the same group tend to show patterns in 原子半径, 电离能, and 电负性. 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.[18] There are exceptions to these trends, however, an example of which occurs in 11族元素 where electronegativity increases farther down the group.[19]

{{Periodic table (group names)}}


{{Main article|Period (periodic table)}} 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 镧系元素s and 锕系元素s form two substantial horizontal series of elements.[20]

Elements in the same period show trends in atomic radius, ionization energy, 电子亲合能, 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.[21] 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.[18] 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.[22]


{{Main article|Block (periodic table)}}

Left to right: s-, f-, d-, p-block in the periodic table

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. Each block is named according to the subshell in which the "last" electron notionally resides.[23]{{#tag:ref|There is an inconsistency and some irregularities in this convention. Thus, helium is shown in the p-block but is actually an s-block element, and (for example) the d-subshell in the d-block is actually filled by the time group 11 is reached, rather than group 12.|group=n}} The 元素分区 comprises the first two groups (alkali metals and alkaline earth metals) as well as hydrogen and helium. The 元素分区 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 类金属s. The 元素分区 comprises groups 3 to 12 (or 3B to 2B in American group numbering) and contains all of the 过渡金属. The 元素分区, often offset below the rest of the periodic table, has no group numbers and comprises lanthanides and actinides.[24]

Metals, metalloids and nonmetals[编辑]

border=none}}}} in the periodic table. Sources disagree on the classification of some of these elements.

According to their shared physical and chemical properties, the elements can be classified into the major categories of 金属s, 类金属s and 非金属元素s. Metals are generally shiny, highly conducting solids that form alloys with one another and salt-like ionic compounds with nonmetals (other than the 稀有气体es). The majority of nonmetals are coloured or colourless insulating gases; nonmetals that form compounds with other nonmetals feature covalent bonding. In between metals and nonmetals are metalloids, which have intermediate or mixed properties.[25]

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 are 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. The nonmetals are simply subdivided into the 非金属元素, which, being nearest to the metalloids, show some incipient metallic character; the 非金属元素, which are essentially nonmetallic; and the monatomic noble gases, which are nonmetallic and almost completely inert. Specialized groupings such as the 難熔金屬 and the 抗腐蚀金属, which are subsets (in this example) of the transition metals, are also known[26] and occasionally denoted.[27]

Placing the elements into categories and subcategories based on shared properties is imperfect. There is a spectrum of properties within each category and it is not hard to find overlaps at the boundaries, as is the case with most classification schemes.[28] 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 nonmetal and a noble gas yet has some cationic chemistry that is more characteristic of a metal. Other classification schemes are possible such as the division of the elements into {{tsl|en|Goldschmidt classification||mineralogical occurrence categories}}, or {{tsl|en|Periodic table (crystal structure)||crystalline structures}}. Categorizing the elements in this fashion dates back to at least 1869 when Hinrichs[29] wrote that simple boundary lines could be drawn on the periodic table to show elements having like properties, such as the metals and the nonmetals, or the gaseous elements.

Periodic trends and patterns[编辑]

{{Main article|Periodic trends}}

Electron configuration[编辑]

{{Main article|Electronic configuration}}

Approximate order in which shells and subshells are arranged by increasing energy according to the 構造原理

The electron configuration or organisation of electrons orbiting neutral atoms shows a recurring pattern or periodicity. The electrons occupy a series of 電子層s (numbered shell 1, shell 2, and so on). Each shell consists of one or more 電子層 (named s, p, d, f and g). As 原子序数 increases, electrons progressively fill these shells and subshells more or less according to the 構造原理 or energy ordering rule, as shown in the diagram. The electron configuration for , 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—two 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 and the 碱金属.[30][31]

Periodic table trends (arrows show an increase)

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 約翰·沃爾夫岡·德貝萊納 the 尼尔斯·玻尔, 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.[30][31]

Atomic radii[编辑]

{{main article|Atomic radius}}


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 量子力学.[32]

The electrons in the 4f-subshell, which is progressively filled from (element 58) to (element 70), 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.[33] Hence has virtually the same atomic radius (and chemistry) as , and has an atomic radius similar to , and so forth. This is known as the 镧系收缩. The effect of the lanthanide contraction is noticeable up to (element 78), after which it is masked by a 相对论量子化学 known as the 惰性电子对效应.[34] The {{tsl|en|d-block contraction||}}, which is a similar effect between the 元素分区 and 元素分区, is less pronounced than the lanthanide contraction but arises from a similar cause.[33]

Ionization energy[编辑]

Ionization energy: each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases

{{main article|Ionization energy}}

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.[34]

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 -like configuration of Mg2+. Similar jumps occur in the ionization energies of other third-row atoms.[34]


{{main article|Electronegativity}}

Graph showing increasing electronegativity with growing number of selected groups

Electronegativity is the tendency of an 原子 to attract 电子s.[35] An atom's electronegativity is affected by both its 原子序数 and the distance between the 價電子 and the nucleus. The higher its electronegativity, the more an element attracts electrons. It was first proposed by 萊納斯·鮑林 in 1932.[36] In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, is the most electronegative of the elements,{{#tag:ref|While fluorine is the most electronegative of the elements under the 电负性, is the most electronegative element under other scales, such as the 电负性.|group=n}} while is the least, at least of those elements for which substantial data is available.[19]

There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than and 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.[19] The anomalously high electronegativity of lead, 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.[37]

Electron affinity[编辑]

{{main article|Electron affinity}}

Dependence of electron affinity on atomic number.[38] Values generally increase across each period, culminating with the halogens before decreasing precipitously with the noble gases. Examples of localized peaks seen in hydrogen, the alkali metals and the 11族元素 are caused by a tendency to complete the s-shell (with the 6s shell of gold being further stabilized by relativistic effects and the presence of a filled 4f sub shell). Examples of localized troughs seen in the alkaline earth metals, and nitrogen, phosphorus, manganese and rhenium are caused by filled s-shells, or half-filled p- or d-shells.[39]

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, 非金属元素 have more positive electron affinity values than 金属s. 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.[40]

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.[40]

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. However, 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.[41]

Metallic character[编辑]

The lower the values of ionization energy, electronegativity and electron affinity, the more 金属lic character the element has. Conversely, nonmetallic character increases with higher values of these properties.[42] 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 相对论量子化学,[43] tends to increase going down a group (or column or family). Thus, the most metallic elements (such as and ) are found at the bottom left of traditional periodic tables and the most nonmetallic elements (, , ) at the top right. The combination of horizontal and vertical trends in metallic character explains the stair-shaped {{tsl|en|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 类金属s.[44][45]

Linking or bridging groups[编辑]

{{Periodic table (micro)|mark=Sc, Y, La, Ac, Lu, Lr, Cu, Ag, Au, Zn, Cd, Hg, He, Ne, Ar, Kr, Xe, Rn|caption=32-column periodic table showing, from left to right, the location of group 3; lutetium and lawrencium; groups 11–12; and the noble gases}}

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.[46] These groups, like the metalloids, show properties in between, or that are a mixture of, groups to either side. Chemically, the group 3 elements, scandium, yttrium, lanthanum and actinium behave largely like the alkaline earth metals[47] or, more generally, s block metals[48][49] but have some of the physical properties of d block transition metals.[50] Lutetium and lawrencium, at the end of the end of the f block, may constitute another linking or bridging group. Lutetium behaves chemically as a lanthanide but shows a mix of lanthanide and transition metal physical properties.[51][52] Lawrencium, as an analogue of lutetium, would presumable display like characteristics.{{#tag:ref|While Lr is thought to have p rather than d electron in its ground-state electron configuration, and would therefore be expected to be volatile metal capable of forming a +1 cation in solution, no evidence of either of these properties has been able to be obtained despite experimental attempts to do so.[53] It was originally expected to have a d electron in its electron configuration[53] and this may still be the case for metallic lawrencium, whereas gas phase atomic lawrencium is very likely thought to have a p electron.[54]|group=n}} The coinage metals in group 11 (copper, silver, and gold) are chemically capable of acting as either transition metals or main group metals.[55] 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.[56][57] 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.[46]


{{Main article|History of the periodic table}}

First systemization attempts[编辑]

The 化學元素發現年表 mapped to significant periodic table development dates (pre-, per- and post-)

In 1789, 安托万-洛朗·德·拉瓦锡 published a list of 33 化學元素s, grouping them into 气体es, 金属s, 非金属元素s, and {{tsl|en|Earth (chemistry)||earths}}.[58] Chemists spent the following century searching for a more precise classification scheme. In 1829, 約翰·沃爾夫岡·德貝萊納 observed that many of the elements could be grouped into triads based on their chemical properties. , , and , for example, were grouped together in a triad as soft, {{tsl|en|Reactivity (chemistry)||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;[59] this became known as the {{tsl|en|Döbereiner's triads||Law of Triads}}.[60] German chemist {{tsl|en|Leopold Gmelin||}} worked with this system, and by 1843 he had identified ten triads, three groups of four, and one group of five. 讓-巴蒂斯特·杜馬 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.[59]

In 1857, German chemist 弗里德里希·奥古斯特·凯库勒·冯·斯特拉多尼茨 observed that often has four other atoms bonded to it. 甲烷, for example, has one carbon atom and four hydrogen atoms.[61] This concept eventually became known as 化合价; different elements bond with different numbers of atoms.[62]

In 1862, {{tsl|en|Alexandre-Emile Béguyer de Chancourtois||}}, a French geologist, published an early form of 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 德米特里·伊万诺维奇·门捷列夫.[63]

In 1864, 尤利乌斯·洛塔尔·迈耶尔, a German chemist, published a table with 44 elements arranged by valency. The table showed that elements with similar properties often shared the same valency.[64] Concurrently, {{tsl|en|William Odling||}} (an English chemist) 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".[65] Odling alluded to the idea of a periodic law but did not pursue it.[66] He subsequently proposed (in 1870) a valence-based classification of the elements.[67]

约翰·纽兰兹 periodic table, as presented to the Chemical Society in 1866, and based on the law of octaves

English chemist 约翰·纽兰兹 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 純八度s of music.[68][69] This so termed {{tsl|en|Law of Octaves||}}, however, was ridiculed by Newlands' contemporaries, and the {{tsl|en|Chemical Society||}} refused to publish his work.[70] Newlands was nonetheless able to draft a table of the elements and used it to predict the existence of missing elements, such as .[71] The Chemical Society only acknowledged the significance of his discoveries five years after they credited Mendeleev.[72]

In 1867, {{tsl|en|Gustavus Detlef Hinrichs||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.[73][74]

Mendeleev's table[编辑]

{{multiple image|direction=vertical|align=left|width=147|image1=Medeleeff by repin.jpg|caption1=Dmitri Mendeleev, watercolour by 伊利亚·叶菲莫维奇·列宾|image2=D. Mendeleev's Periodic table from his book.JPG|caption2=Mendeleev's periodic table from his book An Attempt Towards a Chemical Conception of the Ether}}

A version of Mendeleev's 1869 periodic table: An experiment on a system of elements based on their atomic weights and chemical similarities. This early arrangement presents the periods vertically, and the groups horizontally.

Russian chemistry professor 德米特里·伊万诺维奇·门捷列夫 and German chemist 尤利乌斯·洛塔尔·迈耶尔 independently published their periodic tables in 1869 and 1870, respectively.[75] Mendeleev's table was his first published version; that of Meyer was an expanded version of his (Meyer's) table of 1864.[76] 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.[77]

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.[78] 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 {{tsl|en|Mendeleev's predicted elements||missing elements}}, such as and .[79] The second decision was to occasionally ignore the order suggested by the {{tsl|en|atomic weights||}} and switch adjacent elements, such as and , to better classify them into 族 (化学). Later in 1913, 亨利·莫塞莱 determined experimental values of the nuclear charge or 原子序数 of each element, and showed that Mendeleev's ordering actually corresponds to the order of increasing atomic number.[80]

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 used atomic weight instead of atomic number to organize the elements, information determinable to fair precision in his time. Atomic weight worked well enough in most cases to (as noted) give a presentation that was able to predict the properties of missing elements more accurately than any other method then known. Substitution of atomic numbers, once understood, gave a definitive, integer-based sequence for the elements, and 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 sequence of atomic numbers is still used today even as new synthetic elements are being produced and studied.[81]

Second version and further development[编辑]

Mendeleev's 1871 periodic table with eight groups of elements. Dashes represented elements unknown in 1871.
Eight-column form of periodic table, updated with all elements discovered to 2016

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.[82] These gaps were subsequently filled as chemists discovered additional naturally occurring elements.[83] It is often stated that the last naturally occurring element to be discovered was (referred to by Mendeleev as eka-caesium) in 1939.[84] However, , produced synthetically in 1940, was identified in trace quantities as a naturally occurring element in 1971.[85]

The popular[86] 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.[87]{{#tag:ref|An antecedent of Deming's 18-column table may be seen in Adams' 16-column Periodic Table of 1911. Adams omits the rare earths and the "radioactive elements" (i.e. the actinides) from the main body of his table and instead shows them as being "careted in only to save space" (rare earths between Ba and eka-Yt; radioactive elements between eka-Te and eka-I). See: Elliot Q. A. (1911). "A modification of the periodic table". Journal of the American Chemical Society. 33(5): 684–688 (687).|group=n}} 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 encyclopaedias of chemistry. It was also distributed for many years by the Sargent-Welch Scientific Company.[88][89][90]

With the development of modern 量子力学 theories of 电子 configurations within atoms, it became apparent that each period (row) in the table corresponded to the filling of a 電子層 of electrons. Larger atoms have more electron sub-shells, so later tables have required progressively longer periods.[91]

格倫·西奧多·西博格 who, in 1945, suggested a new periodic table showing the actinides as belonging to a second f-block series

In 1945, 格倫·西奧多·西博格, an American scientist, made the suggestion that the 锕系元素, like the 镧系元素, 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 诺贝尔奖 in chemistry for his work in synthesizing actinide elements.[92][93]{{#tag:ref|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, however, 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. p. 315–316, {{ISBN|0-444-40776-6}}.|group=n}}

Although minute quantities of some 超铀元素 occur naturally,[3] they were all first discovered in laboratories. Their production has expanded the periodic table significantly, the first of these being , synthesized in 1939.[94] Because many of the transuranic elements are highly unstable and 放射性 quickly, they are challenging to detect and characterize when produced. There have been {{tsl|en|element naming controversy||controversies}} concerning the acceptance of competing discovery claims for some elements, requiring independent review to determine which party has priority, and hence naming rights.[95] In 2010, a joint Russia–US collaboration at 杜布纳, 莫斯科州, Russia, claimed to have synthesized six atoms of Ts (element 117), making it the most recently claimed discovery. It, along with (element 113), (element 115), and Og (element 118), are the four most recently named elements, whose names all became official on 28 November 2016.[96]

Different periodic tables[编辑]

The long- or 32-column table[编辑]

The periodic table in 32-column format

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. Unlike the 18-column form this arrangement results in "no interruptions in the sequence of increasing atomic numbers".[97] The relationship of the f-block to the other blocks of the periodic table also becomes easier to see.[98] 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.[99] 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),[100] and the familiarity of chemists with the modern form (as introduced by Seaborg).[101]

Tables with different structures[编辑]

{{main article|Alternative periodic tables}} Within 100 years of the appearance of Mendeleev's table in 1869 it has been estimated that around 700 different periodic table versions were published.[102] As well as numerous rectangular variations, other periodic table formats have been shaped, for example,{{#tag:ref|See The Internet database of periodic tables for depictions of these kinds of variants.|group=n}} like a circle, cube, cylinder, building, spiral, {{tsl|en|lemniscate||}},[103] 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.[102]

Theodor Benfey's spiral periodic table

A popular[104] alternative structure is that of 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.[105]

Most periodic tables are two-dimensional;[3] however, 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),[106] Wringley's Lamina System (1949),[107] {{tsl|en|Paul-Antoine Giguère||Giguère}}'s Periodic helix (1965)[108] and Dufour's Periodic Tree (1996).[109] Going one further, Stowe's Physicist's Periodic Table (1989)[110] has been described as being four-dimensional (having three spatial dimensions and one colour dimension).[111]

The various forms of periodic tables can be thought of as lying on a chemistry–physics continuum.[112] Towards the chemistry end of the continuum can be found, as an example, Rayner-Canham's "unruly"[113] Inorganic Chemist's Periodic Table (2002),[114] which emphasizes trends and patterns, and unusual chemical relationships and properties. Near the physics end of the continuum is {{tsl|en|Charles Janet||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, 量子力学.[115] A somewhat similar approach has been taken by Alper,[116] albeit criticized by {{tsl|en|Eric Scerri||}} as disregarding the need to display chemical and physical periodicity.[117] 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.[118] 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.[66][119] {{clear}} {{periodic table (left step)}}

Open questions and controversies[编辑]

Placement of hydrogen and helium[编辑]

Simply following electron configurations, hydrogen (electronic configuration 1s1) and helium (1s2) should be placed in groups 1 and 2, above lithium ([He]2s1) and beryllium ([He]2s2).[23] However, such placing is rarely used outside of the context of electron configurations: When the 稀有气体 (then called "inert gases") were first discovered around 1900, they were known as "group 0", reflecting no chemical reactivity of these elements known at that point, and helium was placed on the top that group, as it did share the extreme chemical inertness seen throughout the group. As the group changed its formal number, many authors continued to assign helium directly above neon, in group 18; one of the examples of such placing is the current IUPAC table.[120]

Hydrogen's chemical properties are not very close to those of the alkali metals, which occupy group 1, and on that basis hydrogen is sometimes placed elsewhere: one of the most common alternatives is in group 17;[117] one of the factors behind it is the strictly univalent predominantly non-metallic chemistry of hydrogen, and that of fluorine (the element placed on the top of group 17) is strictly univalent and non-metallic. Sometimes, to show how hydrogen has properties both corresponding to those of the alkali metals and the halogens, it may be shown in two columns simultaneously.[121] Another suggestion is above carbon in group 14: placed that way, it fits well into the trend of increasing trends of ionization potential values and electron affinity values, and is not too far from the electronegativity trend, even though hydrogen cannot show the tetravalence characteristic of the heavier group 14 elements.[122] Finally, hydrogen is sometimes placed separately from any group; this is based on how general properties of hydrogen differ from that of any group: unlike hydrogen, the other group 1 elements show extremely metallic behaviour; the group 17 elements commonly form salts (hence the term "halogen"); elements of any other group show some multivalent chemistry. The other period 1 element, helium, is sometimes placed separately from any group as well.[123] The property that distinguishes helium from the rest of the noble gases (even though the extraordinary inertness of helium is extremely close to that of neon and argon[124]) is that in its closed electron shell, helium has only two electrons in the outermost electron orbital, while the rest of the noble gases have eight.

{{anchor|1=Period 6 and 7 elements in group 3}} Group 3 and its elements in periods 6 and 7[编辑]

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.[125]{{#tag:ref|But for the existence of the lanthanides the composition of group 3 would not have been a source of any special interest, since scandium, yttrium, lanthanum and actinium exhibit the same gradual change in properties as do calcium, strontium, barium and radium in group 2.[126]|group=n}}{{#tag:ref|The detachment of the lanthanides from the main body of the periodic table has been attributed to the Czech chemist {{tsl|en|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 尼尔斯·玻尔 continued the detachment process by locating the lanthanides between the s- and d-blocks. In 1949 格倫·西奧多·西博格 (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 化学化工新闻, and in the years up to 1949 several authors commented on, and generally agreed with, Seaborg's proposal. In that year he noted that the best method for presenting the actinides seemed to be by positioning them below, and as analogues of, the lanthanides. See: Thyssen P. and Binnemans K. (2011). "Accommodation of the Rare Earths in the Periodic Table: A Historical Analysis". In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 41. Amsterdam: Elsevier, pp. 1–94; Seaborg G. T. (1994). Origin of the Actinide Concept'. In K. A. Gschneider Jr. (ed). Handbook on the Physics and Chemistry of the Rare Earths. 18. Amsterdam: Elsevier, pp. 1–27.|group=n}} It has been claimed that such arguments are proof that, "it is a mistake to break the [periodic] system into sharply delimited blocks".[127] A third variant shows the two positions below yttrium as being occupied by the lanthanides and the actinides.

Chemical and physical arguments have been made in support of lutetium and lawrencium[128][129] but the majority of authors seem unconvinced.[130] Most working chemists are not aware there is any controversy.[131] In December 2015 an IUPAC project was established to make a recommendation on the matter.[132]

Lanthanum and actinium[编辑]

La and Ac below Y

Lanthanum and actinium are commonly depicted as the remaining group 3 members.[133]{{#tag:ref|For examples of this table see 彼得·阿特金斯 et al. (2006). Shriver & Atkins Inorganic Chemistry (4th ed.). Oxford: Oxford University Press • Myers et al. (2004). Holt Chemistry. Orlando: Holt, Rinehart & Winston • {{tsl|en|Raymond Chang (chemist)||Chang R.}} (2000). Essential Chemistry (2nd ed.). Boston: McGraw-Hill|group=n}} It has been suggested that this layout originated in the 1940s, with the appearance of periodic tables relying on the electron configurations of the elements and the notion of the differentiating electron. The configurations of caesium, barium and lanthanum are [Xe]6s1, [Xe]6s2 and [Xe]5d16s2. Lanthanum thus has a 5d differentiating electron and this establishes "it in group 3 as the first member of the d-block for period 6".[134] A consistent set of electron configurations is then seen in group 3: scandium [Ar]3d14s2, yttrium [Kr]4d15s2 and lanthanum [Xe]5d16s2. Still in period 6, ytterbium was assigned an electron configuration of [Xe]4f135d16s2 and lutetium [Xe]4f145d16s2, "resulting in a 4f differentiating electron for lutetium and firmly establishing it as the last member of the f-block for period 6".[134] Later {{tsl|en|electron spectroscopy||spectroscopic}} work found that the electron configuration of ytterbium was in fact [Xe]4f146s2. This meant that ytterbium and lutetium—the latter with [Xe]4f145d16s2—both had 14 f-electrons, "resulting in a d- rather than an f- differentiating electron" for lutetium and making it an "equally valid candidate" with [Xe]5d16s2 lanthanum, for the group 3 periodic table position below yttrium.[134] Lanthanum, however, has the advantage of incumbency since the 5d1 electron appears for the first time in its structure whereas it appears for the third time in lutetium, having also made a brief second appearance in gadolinium.[135]

In terms of chemical behaviour,[136] and trends going down group 3 for properties such as melting point, electronegativity and ionic radius,[137][138] scandium, yttrium, lanthanum and actinium are similar to their group 1–2 counterparts. In this variant, the number of f electrons in the most common (trivalent) ions of the f-block elements consistently matches their position in the f-block.[139] For example, the f-electron counts for the trivalent ions of the first three f-block elements are Ce 1, Pr 2 and Nd 3.[140]

Lutetium and lawrencium[编辑]

Lu and Lr below Y

In other tables, lutetium and lawrencium are the remaining group 3 members.{{#tag:ref|For examples of the group 3 = Sc-Y-Lu-Lr table see Rayner-Canham G. & Overton T. (2013). Descriptive Inorganic Chemistry (6th ed.). New York: W. H. Freeman and Company • Brown et al. (2009). Chemistry: The Central Science (11th ed.). Upper Saddle River, New Jersey: Pearson Education • Moore et al. (1978). Chemistry. Tokyo: McGraw-Hill Kogakusha|group=n}} Early techniques for chemically separating scandium, yttrium and lutetium relied on the fact that these elements occurred together in the so-called "yttrium group" whereas La and Ac occurred together in the "cerium group".[134] Accordingly, lutetium rather than lanthanum was assigned to group 3 by some chemists in the 1920s and 30s.{{#tag:ref|The phenomenon of different separation groups is caused by increasing basicity with increasing radius, and does not constitute a fundamental reason to show Lu, rather than La, below Y. Thus, among the Group 2 碱土金属s, Mg (less basic) belongs in the "soluble group" and Ca, Sr and Ba (more basic) occur in the "ammonium carbonate group". Nevertheless, Mg, Ca, Sr and Ba are routinely collocated in Group 2 of the periodic table. See: Moeller et al. (1989). Chemistry with Inorganic Qualitative Analysis (3rd ed.). SanDiego: Harcourt Brace Jovanovich, pp. 955–956, 958.|group=n}} Several physicists in the 1950s and 60s favoured lutetium, in light of a comparison of several of its physical properties with those of lanthanum.[134] This arrangement, in which lanthanum is the first member of the f-block, is disputed by some authors since lanthanum lacks any f-electrons. However, it has been argued that this is not valid concern given other periodic table anomalies—thorium, for example, has no f-electrons yet is part of the f-block.[141] As for lawrencium, its gas phase atomic electron configuration was confirmed in 2015 as [Rn]5f147s27p1. Such a configuration represents another periodic table anomaly, regardless of whether lawrencium is located in the f-block or the d-block, as the only potentially applicable p-block position has been reserved for nihonium with its predicted configuration of [Rn]5f146d107s27p1.[142]{{#tag:ref|Even if metallic lawrencium has a p electron, simple modelling studies suggest it will behave like a lanthanide,[143] (as do the rest of the late actinides).[140]|group=n}}

Chemically, scandium, yttrium and lutetium (and presumably lawrencium) behave like trivalent versions of the group 1–2 metals.[144] On the other hand, trends going down the group for properties such as melting point, electronegativity and ionic radius, are similar to those found among their group 4–8 counterparts.[134] In this variant, the number of f electrons in the gaseous forms of the f-block atoms usually matches their position in the f-block. For example, the f-electron counts for the first five f-block elements are La 0, Ce 1, Pr 3, Nd 4 and Pm 5.[134]

Lanthanides and actinides[编辑]

Markers below Y

A few authors position all thirty lanthanides and actinides in the two positions below yttrium (usually via footnote markers). This variant emphasizes similarities in the chemistry of the 15 lanthanide elements (La–Lu), possibly at the expense of ambiguity as to which elements occupy the two group 3 positions below yttrium, and a 15-column wide f block (there can only be 14 elements in any row of the f block).{{#tag:ref|For examples of the group 3 = Ln and An table see Housecroft C. E. & Sharpe A. G. (2008). Inorganic Chemistry (3rd ed.). Harlow: Pearson Education • Halliday et al. (2005). Fundamentals of Physics (7th ed.). Hoboken, NewJersey: John Wiley & Sons • Nebergall et. al. (1980). General Chemistry (6th ed.). Lexington: D. C. Heath and Company|group=n}}

Groups included in the transition metals[编辑]

The definition of a 过渡金属, as given by IUPAC, is an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell.[145] By this definition all of the elements in groups 3–11 are transition metals. The IUPAC definition therefore excludes group 12, comprising zinc, cadmium and mercury, from the transition metals category.

Some chemists treat the categories "元素分区 elements" and "transition metals" interchangeably, thereby including groups 3–12 among the transition metals. In this instance the group 12 elements are treated as a special case of transition metal in which the d electrons are not ordinarily involved in chemical bonding. The 2007 report of 四氟化汞 (HgF4), a compound in which mercury would use its d electrons for bonding, has prompted some commentators to suggest that mercury can be regarded as a transition metal.[146] Other commentators, such as Jensen,[147] have argued that the formation of a compound like HgF4 can occur only under highly abnormal conditions; indeed, its existence is currently disputed. As such, mercury could not be regarded as a transition metal by any reasonable interpretation of the ordinary meaning of the term.[147]

Still other chemists further exclude the 3族元素s from the definition of a transition metal. They do so on the basis that the group 3 elements do not form any ions having a partially occupied d shell and do not therefore exhibit any properties characteristic of transition metal chemistry.[148] In this case, only groups 4–11 are regarded as transition metals. However, though the group 3 elements show few of the characteristic chemical properties of the transition metals, they do show some of their characteristic physical properties (on account of the presence in each atom of a single d electron).[50]

Elements with unknown chemical properties[编辑]

Although all elements up to oganesson have been discovered, of the elements above 𨭆 (element 108), only (element 112) and (element 114) have known chemical properties. The other elements may behave differently from what would be predicted by extrapolation, due to 相对论量子化学 effects; for example, flerovium has been predicted to possibly exhibit some noble-gas-like properties, even though it is currently placed in the 碳族元素.[149] More recent experiments have suggested, however, that flerovium behaves chemically like lead, as expected from its periodic table position.[150]

Further periodic table extensions[编辑]

{{main article|Extended periodic table}}

{{Periodic table (micro)|number=120|caption=Periodic table with eight rows, extended to element 172[151]}}

It is unclear whether new elements will continue the pattern of the current periodic table as 扩展元素周期表, or require further adaptations or adjustments. 格倫·西奧多·西博格 expected the eighth period to follow the previously established pattern exactly, so that it would include a two-element s-block for elements Uue and Ubn, a new 扩展元素周期表 for the next 18 elements, and 30 additional elements continuing the current f-, d-, and p-blocks, culminating in element 168, the next noble gas.[152] More recently, physicists such as {{tsl|en|Pekka Pyykkö||}} have theorized that these additional elements do not follow the 構造原理, which predicts how electron shells are filled and thus affects the appearance of the present periodic table. There are currently several competing theoretical models for the placement of the elements of atomic number less than or equal to 172.[153]

Element with the highest possible atomic number[编辑]

The number of possible elements is not known. A very early suggestion made by Elliot Adams in 1911, and based on the arrangement of elements in each horizontal periodic table row, was that elements of atomic weight greater than 256± (which would equate to between elements 99 and 100 in modern-day terms) did not exist.[154] A higher—more recent—estimate is that the periodic table may end soon after the 稳定岛,[155] which is expected to centre around 扩展元素周期表, as the extension of the periodic and nuclides tables is restricted by proton and neutron 原子核滴线.[156] Other predictions of an end to the periodic table include at element 128 by 约翰·埃姆斯利,[3] at element 137 by 理查德·費曼,[157] and at element 155 by Albert Khazan.[3]{{#tag:ref|Karol (2002, p. 63) contends that gravitational effects would become significant when atomic numbers become astronomically large, thereby overcoming other super-massive nuclei instability phenomena, and that 中子星s (with atomic numbers on the order of 1021) can arguably be regarded as representing the heaviest known elements in the universe. See: Karol P. J. (2002). "The Mendeleev–Seaborg periodic table: Through Z = 1138 and beyond". Journal of Chemical Education 79 (1): 60–63.|group=n}}

Bohr model

The 玻尔模型 exhibits difficulty for atoms with atomic number greater than 137, as any element with an atomic number greater than 137 would require 1s electrons to be travelling faster than c, the 光速.[158] Hence the non-relativistic Bohr model is inaccurate when applied to such an element.

Relativistic Dirac equation

The 相对论 Dirac equation has problems for elements with more than 137 protons. For such elements, the wave function of the Dirac ground state is oscillatory rather than bound, and there is no gap between the positive and negative energy spectra, as in the {{tsl|en|Klein paradox||}}.[159] More accurate calculations taking into account the effects of the finite size of the nucleus indicate that the binding energy first exceeds the limit for elements with more than 173 protons. For heavier elements, if the innermost orbital (1s) is not filled, the electric field of the nucleus will pull an electron out of the vacuum, resulting in the 正电子发射;[160] however, this does not happen if the innermost orbital is filled, so that element 173 is not necessarily the end of the periodic table.[157]

Optimal form[编辑]

The many different forms of periodic table have prompted the question of whether there is an optimal or definitive form of periodic table. The answer to this question is thought to depend on whether the chemical periodicity seen to occur among the elements has an underlying truth, effectively hard-wired into the universe, or if any such periodicity is instead the product of subjective human interpretation, contingent upon the circumstances, beliefs and predilections of human observers. An objective basis for chemical periodicity would settle the questions about the location of hydrogen and helium, and the composition of group 3. Such an underlying truth, if it exists, is thought to have not yet been discovered. In its absence, the many different forms of periodic table can be regarded as variations on the theme of chemical periodicity, each of which explores and emphasizes different aspects, properties, perspectives and relationships of and among the elements.{{#tag:ref|{{tsl|en|Eric Scerri||Scerri}}, one of the foremost authorities on the history of the periodic table,[161] favoured the concept of an optimal form of periodic table but has recently changed his mind and now supports the value of a plurality of periodic tables.[162]|group=n}}

See also[编辑]

{{Wikipedia books|Periodic table}} {{div col|2}}

{{div col end}}







  • {{cite book|last=Ball|first=P.|authorlink=Philip Ball|title=The Ingredients: A Guided Tour of the Elements |location=Oxford|publisher=Oxford University Press |year=2002 |isbn=0-19-284100-9}}
  • {{cite book|last=Chang|first=R.|authorlink=Raymond Chang (chemist)|title=Chemistry|year=2002|edition=7th|publisher=McGraw-Hill Higher Education|location=New York|isbn=978-0-19-284100-1}}
  • {{cite book|last=Gray|first=T.|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|last1=Greenwood|first1=N. N.|authorlink1=Norman Greenwood|last2=Earnshaw|first2=A.|year=1984|title=Chemistry of the Elements|place=Oxford|publisher=Pergamon Press|isbn=0-08-022057-6}}
  • {{cite book |last1=Huheey|first1=J. E.|last2=Keiter|first2=E. A.|last3=Keiter|first3=R. L.|title=Principles of structure and reactivity|publisher=Harper Collins College Publishers|location=New York|edition=4th|isbn=0-06-042995-X}}
  • {{cite book |last=Moore |first=J. T.|edition=1st|title=Chemistry For Dummies|series = {{tsl|en|For Dummies||}} |year=2003 |publisher=Wiley Publications |location=New York |isbn=978-0-7645-5430-8}}
  • {{cite book|last=Scerri |first=E.|authorlink=Eric Scerri|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=Scerri|first=E.|authorlink=Eric Scerri|title=The periodic table: A very short introduction|year=2011|publisher=Oxford University Press|location=Oxford|isbn=978-0-19-958249-5}}
  • {{cite book |last=Venable|first=F. P.|authorlink=Francis Preston Venable|title=The Development of the Periodic Law|year=1896|url= Publishing Company|location=Easton, Pennsylvania}}


External links[编辑]

{{Sister project links|Periodic table}}

{{Periodic table (32 columns, compact)}} {{PeriodicTablesFooter}} {{BranchesofChemistry}} {{Use dmy dates|date=June 2015}} {{Authority control}}

Category:元素周期表 Category:化学元素 {{tsl|en|Category:Classification systems||}} {{tsl|en|Category:Dmitri Mendeleev||}} Category:俄羅斯發明 Category:1869年作品

  1. ^ {{cite web|url= Four elements added to periodic table|work=BBC News|date=January 4, 2016}}
  2. ^ {{cite web |first= Nicholas |last=St. Fleur |url= New Names Officially Added to the Periodic Table of Elements |work=New York Times|date=December 1, 2016}}
  3. ^ 3.0 3.1 3.2 3.3 3.4 3.5 {{cite book|last=Emsley|first=J.|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}}
  4. ^ Greenwood & Earnshaw, pp. 24–27
  5. ^ Gray, p. 6
  6. ^ {{Cite web|url= elements on the periodic table are named|last=CNN|first=Ashley Strickland|website=CNN|access-date=2016-06-11}}
  7. ^ {{cite web|url= and assignment of elements with atomic numbers 113, 115, 117 and 118|publisher=國際純化學和應用化學聯合會|date=2015-12-30}}
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