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Periodic table

This article is about the table used in chemistry. For other uses, see Periodic table (disambiguation).

Modern periodic table, in 18-column layout (colour legend below) The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number (number of protons), electron configurations, and recurring chemical properties. This ordering shows periodic trends, such as elements with similar behaviour in the same column. It also shows four rectangular blocks with some approximately similar chemical properties. In general, within one row (period) the elements are metals on the left, and non-metals on the right.

The rows of the table are called periods; the columns are called groups. Six groups (columns) have names as well as numbers: for example, group 17 elements are the halogens; and group 18, the noble gases. 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.

Dmitri Mendeleev published in 1869 the first widely recognized periodic table. He developed his table to illustrate periodic trends in the properties of the then-known elements. Mendeleev also predicted some properties of 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 the discovery or synthesis of further new elements and the development of new theoretical models to explain chemical behaviour.

All elements from atomic numbers 1 (hydrogen) to 118 (ununoctium) have been discovered or synthesized, with the most recent additions (elements 113, 115, 117, and 118) being confirmed by the IUPAC on December 30, 2015: they complete the first seven rows of the periodic table.[1] The first 94 elements exist naturally, although some are found only in trace amounts and were synthesized in laboratories before being found in nature.[n 1] Elements with atomic numbers from 95 to 118 have only been synthesized in laboratories or nuclear reactors.[2] Synthesis of elements having higher atomic numbers is being pursued. Numerous synthetic radionuclides of naturally occurring elements have also been produced in laboratories.

Contents Overview

For large cell versions, see Periodic table (large cells). v t e Periodic table Group 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Alkali metals Alkaline earth metals Pnicto­gens Chal­co­gens Halo­gens Noble gases Period 1

Hydro­gen 1 H He­lium 2 He 2 Lith­ium 3 Li Beryl­lium 4 Be Boron 5 B Carbon 6 C Nitro­gen 7 N Oxy­gen 8 O Fluor­ine 9 F Neon 10 Ne 3 So­dium 11 Na Magne­sium 12 Mg Alumin­ium 13 Al Sili­con 14 Si Phos­phorus 15 P Sulfur 16 S Chlor­ine 17 Cl Argon 18 Ar 4 Potas­sium 19 K Cal­cium 20 Ca Scan­dium 21 Sc Tita­nium 22 Ti Vana­dium 23 V Chrom­ium 24 Cr Manga­nese 25 Mn Iron 26 Fe Cobalt 27 Co Nickel 28 Ni Copper 29 Cu Zinc 30 Zn Gallium 31 Ga Germa­nium 32 Ge Arsenic 33 As Sele­nium 34 Se Bromine 35 Br Kryp­ton 36 Kr 5 Rubid­ium 37 Rb Stront­ium 38 Sr Yttrium 39 Y Zirco­nium 40 Zr Nio­bium 41 Nb Molyb­denum 42 Mo Tech­netium 43 Tc Ruthe­nium 44 Ru Rho­dium 45 Rh Pallad­ium 46 Pd Silver 47 Ag Cad­mium 48 Cd Indium 49 In Tin 50 Sn Anti­mony 51 Sb Tellur­ium 52 Te Iodine 53  I  Xenon 54 Xe 6 Cae­sium 55 Cs Ba­rium 56 Ba 1 asterisk Lute­tium 71 Lu Haf­nium 72 Hf Tanta­lum 73 Ta Tung­sten 74 W Rhe­nium 75 Re Os­mium 76 Os Iridium 77 Ir Plat­inum 78 Pt Gold 79 Au Mer­cury 80 Hg Thallium 81 Tl Lead 82 Pb Bis­muth 83 Bi Polo­nium 84 Po Asta­tine 85 At Radon 86 Rn 7 Fran­cium 87 Fr Ra­dium 88 Ra 1 asterisk Lawren­cium 103 Lr Ruther­fordium 104 Rf Dub­nium 105 Db Sea­borgium 106 Sg Bohr­ium 107 Bh Has­sium 108 Hs Meit­nerium 109 Mt Darm­stadtium 110 Ds Roent­genium 111 Rg Coper­nicium 112 Cn Unun­trium 113 Uut Flerov­ium 114 Fl Unun­pentium 115 Uup Liver­morium 116 Lv Unun­septium 117 Uus Unun­octium 118 Uuo 1 asterisk Lan­thanum 57 La Cerium 58 Ce Praseo­dymium 59 Pr Neo­dymium 60 Nd Prome­thium 61 Pm Sama­rium 62 Sm Europ­ium 63 Eu Gadolin­ium 64 Gd Ter­bium 65 Tb Dyspro­sium 66 Dy Hol­mium 67 Ho Erbium 68 Er Thulium 69 Tm Ytter­bium 70 Yb

1 asterisk Actin­ium 89 Ac Thor­ium 90 Th Protac­tinium 91 Pa Ura­nium 92 U Neptu­nium 93 Np Pluto­nium 94 Pu Ameri­cium 95 Am Curium 96 Cm Berkel­ium 97 Bk Califor­nium 98 Cf Einstei­nium 99 Es Fer­mium 100 Fm Mende­levium 101 Md Nobel­ium 102 No

black=solid green=liquid red=gas gray=unknown Color of the atomic number shows state of matter (at 0 °C and 1 atm) Primordial From decay Synthetic Border shows natural occurrence of the element Background color shows subcategory in the metal–metalloid–nonmetal trend: Metal Metalloid Nonmetal Unknown chemical properties Alkali metal Alkaline earth metal Lan­thanide Actinide Transition metal Post-​transition metal Polyatomic nonmetal Diatomic nonmetal Noble gas Each chemical element has a unique atomic number (Z) representing the number of protons in its nucleus.[n 2] Most elements have differing numbers of neutrons among different atoms, with these variants being referred to as isotopes. For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and a very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element. Elements with no stable isotopes have the atomic masses of their most stable isotopes, where such masses are shown, listed in parentheses.[3]

In the standard periodic table, the elements are listed in order of increasing atomic number (the number of protons in the nucleus of an atom). A new row (period) is started when a new electron shell has its first electron. Columns (groups) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen and selenium are in the same column because they both have four electrons in the outermost p-subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.[4]

As of 2016, the periodic table has 118 confirmed elements, from element 1 (hydrogen) to 118 (ununoctium). Elements 113, 115, 117 and 118 were officially confirmed by the International Union of Pure and Applied Chemistry (IUPAC) in December 2015. Their proposed names, nihonium (Nh), moscovium (Mc), tennessine (Ts) and oganesson (Og) respectively, were announced by the IUPAC in June 2016.[5][6] These names will not be formally approved until after the five-month public comment period ends in November 2016.[7] Until then, they are formally identified by their atomic number (e.g., "element 113"), or by their provisional systematic name ("ununtrium", symbol "Uut").[8]

The first 94 elements occur naturally; the remaining 24, americium to ununoctium (95–118) occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements.[2] 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).[9]

Grouping methods

Groups 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 valence shell.[10] Consequently, elements in the same group tend to have a shared chemistry and exhibit a clear trend in properties with increasing atomic number.[11] 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.[12][13][14]

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).[15] 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.[16]

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.[15]

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.[17] There are exceptions to these trends, however, an example of which occurs in group 11 where electronegativity increases farther down the group.[18]

Group numbera 1 2 3d 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Mendeleev (I–VIII) I II III IV V VI VII VIII I II III IV V VI VII b CAS (US, pattern A-B-A) IA IIA IIIB IVB VB VIB VIIB VIIIB IB IIB IIIA IVA VA VIA VIIA VIIIA old IUPAC (Europe, pattern A-B) IA IIA IIIA IVA VA VIA VIIA VIII IB IIB IIIB IVB VB VIB VIIB 0 Trivial name Alkali metals Alkaline earth metals Coin­age metalse Vola­tile metalse Icosa­gense Crys­tallo­gense Pnicto­gens Chal­co­gens Halo­gens Noble gases Name by element Lith­ium group Beryl­lium groupsp Scan­dium group Titan­ium group Vana­dium group Chro­mium group Man­ga­nese group Iron group Co­balt group Nickel group Cop­per group Zinc group Boron group Car­bon group Nitro­gen group Oxy­gen group Fluor­ine group Helium or Neon group Period 1 H c He Period 2 Li Be B C N O F Ne Period 3 Na Mg Al Si P S Cl Ar Period 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr Period 5 Rb Sr d Y Zr Nb Mo Tc Ru