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