Periodic table
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The periodic table, also known as the periodic table of the elements, is an ordered arrangement of the chemical elements into rows ("periods") and columns ("groups"). An icon of chemistry, the periodic table is also widely used in physics and other sciences. It is a depiction of the periodic law, which states that when the elements are arranged in order of their atomic numbers an approximate recurrence of their properties is evident. The table is divided into four roughly rectangular areas called blocks. Elements in the same group tend to show similar chemical characteristics.
Vertical, horizontal and diagonal trends characterize the periodic table. Metallic character increases going down a group and from right to left across a period. Nonmetallic character increases going from the bottom left of the periodic table to the top right.
The first periodic table to become generally accepted was that of the Russian chemist Dmitri Mendeleev in 1869; he formulated the periodic law as a dependence of chemical properties on atomic mass. As not all elements were then known, there were gaps in his periodic table, and Mendeleev successfully used the periodic law to predict some properties of some of the missing elements. The periodic law was recognized as a fundamental discovery in the late 19th century. It was explained early in the 20th century, with the discovery of atomic numbers and associated pioneering work in quantum mechanics, both ideas serving to illuminate the internal structure of the atom. A recognisably modern form of the table was reached in 1945 with Glenn T. Seaborg's discovery that the actinides were in fact f-block rather than d-block elements. The periodic table and law have become a central and indispensable part of modern chemistry.
The periodic table continues to evolve with the progress of science. In nature, only elements up to atomic number 94 exist;[a] elements beyond that are synthetic elements that can only be obtained artificially. By 2010, the first 118 elements were known, thereby completing the first seven rows of the table;[1] however, chemical characterization is still needed for the heaviest elements to confirm that their properties match their positions. New discoveries will extend the table beyond these seven rows, though it is not yet known how many more elements are possible; moreover, theoretical calculations suggest that this unknown region will not follow the patterns of the known part of the table. Some scientific discussion also continues regarding whether some elements are correctly positioned in the table. Many alternative representations of the periodic law exist, and there is some discussion as to whether there is an optimal form of the periodic table.
Structure
| Group | 1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |
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| Hydrogen & alkali metals |
Alkaline earth metals | Triels | Tetrels | Pnictogens | Chalcogens | Halogens | Noble gases | ||||||||||||
| Period |
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| 6 |  |
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| 7 |  |
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Primordial From decay Synthetic Border shows natural occurrence of the element
- Ca: 40.078 — Abridged value (uncertainty omitted here)[3]
- Po: [209] — mass number of the most stable isotope
Each chemical element has a unique atomic number (Z— for "Zahl", German for "number") representing the number of protons in its nucleus.[4] Each distinct atomic number therefore corresponds to a class of atom: these classes are called the chemical elements.[5] The chemical elements are what the periodic table classifies and organizes. Hydrogen is the element with atomic number 1; helium, atomic number 2; lithium, atomic number 3; and so on. Each of these names can be further abbreviated by a one- or two-letter chemical symbol; those for hydrogen, helium, and lithium are respectively H, He, and Li.[6] Neutrons do not affect the atom's chemical identity, but do affect its weight. Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element.[6] Naturally occurring elements usually occur as mixes of different isotopes; since each isotope usually occurs with a characteristic abundance, naturally occurring elements have well-defined atomic weights, defined as the average mass of a naturally occurring atom of that element.[7] All elements have multiple isotopes, variants with the same number of protons but different numbers of neutrons. For example, carbon has three naturally occurring isotopes: all of its atoms have six protons and most have six neutrons as well, but about one per cent have seven neutrons, and a very small fraction have eight neutrons. Isotopes are never separated in the periodic table; they are always grouped together under a single element. When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses.[8]
In the standard periodic table, the elements are listed in order of increasing atomic number. A new row (period) is started when a new electron shell has its first electron. Columns (groups) are determined by the electron configuration of the atom; elements with the same number of electrons in a particular subshell fall into the same columns (e.g. oxygen, sulfur, and selenium are in the same column because they all have four electrons in the outermost p-subshell). Elements with similar chemical properties generally fall into the same group in the periodic table, although in the f-block, and to some respect in the d-block, the elements in the same period tend to have similar properties, as well. Thus, it is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.[9]
Today, 118 elements are known, the first 94 of which are known to occur naturally on Earth.[10][a] The remaining 24, americium to oganesson (95–118), occur only when synthesized in laboratories. Of the 94 naturally occurring elements, 83 are primordial and 11 occur only in decay chains of primordial elements. A few of the latter are so rare that they were not discovered in nature, but were synthesized in the laboratory before it was determined that they exist in nature: technetium (element 43), promethium (element 61), astatine (element 85), neptunium (element 93), and plutonium (element 94).[12] No element heavier than einsteinium (element 99) has ever been observed in macroscopic quantities in its pure form, nor has astatine; francium (element 87) has been only photographed in the form of light emitted from microscopic quantities.[14] Of the 94 natural elements, eighty have a stable isotope and one more (bismuth) has an almost-stable isotope (with a half-life of 2.01×1019 years, over a billion times the age of the universe).[15][b] Two more, thorium and uranium, have isotopes undergoing radioactive decay with a half-life comparable to the age of the Earth. The stable elements plus bismuth, thorium, and uranium make up the 83 primordial elements that survived from the Earth's formation.[c] The remaining eleven natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of thorium and uranium.[d] All 24 known artificial elements are radioactive.[6]
Group names and numbers
Under an international naming convention, the groups are numbered numerically from 1 to 18 from the leftmost column (the alkali metals) to the rightmost column (the noble gases). The f-block groups are ignored in this numbering.[22] Groups can also be named by their first element, e.g. the "scandium group" for group 3.[22] Previously, groups were known by Roman numerals. In the United States, the Roman numerals were followed by either an "A" (if the group was in the s- or p-block) or a "B" (if the group was in the d-block). The Roman numerals used correspond to the last digit of today's naming convention (e.g., the group 4 elements were group IVB, and the group 14 elements were group IVA). In Europe, "A" was used for groups 1 through 7, and "B" was used for groups 11 through 17. In addition, groups 8, 9, and 10 used to be treated as one triple-sized group, known collectively in both notations as group VIII. In 1988, the new IUPAC (International Union of Pure and Applied Chemistry) naming system (1–18) was put into use, and the old group names (I–VIII) were deprecated.[23]
| IUPAC group | 1a | 2 | —b | 3c | 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 | d | |||
| CAS (US, A-B-A) | IA | IIA | IIIB | IVB | VB | VIB | VIIB | VIIIB | IB | IIB | IIIA | IVA | VA | VIA | VIIA | VIIIA | |||
| Old IUPAC (Europe, A–B) | IA | IIA | IIIA | IVA | VA | VIA | VIIA | VIII | IB | IIB | IIIB | IVB | VB | VIB | VIIB | 0 | |||
| Trivial namer | H and alkali metals | alkaline earth metals | lanthanoids actinoids |
triels | tetrels | pnictogens | chalcogens | halogens | noble gases | ||||||||||
| Name by elementr | lithium group | beryllium group | scandium group | titanium group | vanadium group | chromium group | manganese group | iron group | cobalt group | nickel group | copper group | zinc group | boron group | carbon group | nitrogen group | oxygen group | fluorine group | helium or neon group | |
| Period 1 | H | He | |||||||||||||||||
| Period 2 | Li | Be | B | C | N | O | F | Ne | |||||||||||
| Period 3 | Na | Mg | Al | Si | P | S | Cl | Ar | |||||||||||
| Period 4 | K | Ca | Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | Ga | Ge | As | Se | Br | Kr | |
| Period 5 | Rb | Sr | Y | Zr | Nb | Mo | Tc | Ru | Rh | Pd | Ag | Cd | In | Sn | Sb | Te | I | Xe | |
| Period 6 | Cs | Ba | La–Yb | Lu | Hf | Ta | W | Re | Os | Ir | Pt | Au | Hg | Tl | Pb | Bi | Po | At | Rn |
| Period 7 | Fr | Ra | Ac–No | Lr | Rf | Db | Sg | Bh | Hs | Mt | Ds | Rg | Cn | Nh | Fl | Mc | Lv | Ts | Og |
b The 14 f-block groups (columns) do not have a group number.
c The correct composition of group 3 is scandium (Sc), yttrium (Y), lutetium (Lu), and lawrencium (Lr), as shown here: this is endorsed by 1988[23] and 2021[24] IUPAC reports on the question. General inorganic chemistry texts often put scandium (Sc), yttrium (Y), lanthanum (La), and actinium (Ac) in group 3, so that Ce–Lu and Th–Lr become the f-block between groups 3 and 4; this was based on incorrectly measured electron configurations from history,[25] and Lev Landau and Evgeny Lifshitz already considered it incorrect in 1948.[26] Arguments can still occasionally be encountered in the contemporary literature purporting to defend it, but most authors consider them logically inconsistent.[27][28][29] Some sources follow a compromise that puts La–Lu and Ac–Lr as the f-block rows (despite that giving 15 f-block elements in each row, which contradicts quantum mechanics), leaving the heavier members of group 3 ambiguous.[24] See also Group 3 element#Composition.
d Group 18, the noble gases, had not been discovered at the time of Mendeleev's original table. Later (1902), Mendeleev accepted the evidence for their existence, and they could be placed in a new "group 0", consistently and without breaking the periodic table principle.
r Group name as recommended by IUPAC.