Editing Chemical element (section)

== Description ==
The lightest elements are [[hydrogen]] and [[helium]], both created by [[Big Bang nucleosynthesis]] in the [[Chronology of the universe|first 20 minutes of the universe]]<ref>See the timeline on p.10 in {{cite journal|year=2006|title=Evidence for Dark Matter|url=http://gaitskell.brown.edu/physics/talks/0408_SLAC_SummerSchool/Gaitskell_DMEvidence_v16.pdf|journal=Physical Review C|volume=74|issue=4|pages=044602|doi=10.1103/PhysRevC.74.044602|bibcode=2006PhRvC..74d4602O|last1=Oganessian|first1=Yu. Ts.|last2=Utyonkov|first2=V.|last3=Lobanov|first3=Yu.|last4=Abdullin|first4=F.|last5=Polyakov|first5=A.|last6=Sagaidak|first6=R.|last7=Shirokovsky|first7=I.|last8=Tsyganov|first8=Yu.|display-authors=|doi-access=free|access-date=8 October 2007|archive-date=13 February 2021|archive-url=https://web.archive.org/web/20210213212406/http://gaitskell.brown.edu/physics/talks/0408_SLAC_SummerSchool/Gaitskell_DMEvidence_v16.pdf|url-status=live}}</ref> in a [[ratio]] of around 3:1 by mass (or 12:1 by number of atoms),<ref>{{cite web
|url=http://pdgusers.lbl.gov/~pslii/uabackup/big_bang/elementabundancies/2300400.html|title=The Universe Adventure Hydrogen and Helium|publisher=Lawrence Berkeley National Laboratory U.S. Department of Energy |year=2005|url-status=dead|archive-url=https://web.archive.org/web/20130921054844/http://pdgusers.lbl.gov/~pslii/uabackup/big_bang/elementabundancies/2300400.html|archive-date=21 September 2013}}</ref><ref>{{cite web|url=http://www.astro.soton.ac.uk/~pac/PH112/notes/notes/node181.html|title=Formation of the light elements|author=astro.soton.ac.uk|publisher=University of Southampton|date=3 January 2001|url-status=dead|archive-url=https://web.archive.org/web/20130921054428/http://www.astro.soton.ac.uk/~pac/PH112/notes/notes/node181.html|archive-date=21 September 2013}}</ref> along with tiny traces of the next two elements, [[lithium]] and [[beryllium]]. Almost all other elements found in nature were made by various natural methods of [[nucleosynthesis]].<ref>{{cite web|url=http://www.foothill.edu/attach/938/Nucleosynthesis.pdf|title=How Stars Make Energy and New Elements|publisher=Foothill College|date=18 October 2006|access-date=17 February 2013|archive-date=11 August 2020|archive-url=https://web.archive.org/web/20200811064522/https://foothill.edu/attach/938/Nucleosynthesis.pdf|url-status=live}}</ref> On Earth, small amounts of new atoms are naturally produced in [[nucleogenic]] reactions, or in [[cosmogenic]] processes, such as [[cosmic ray spallation]]. New atoms are also naturally produced on Earth as [[Radiogenic nuclide|radiogenic]] [[Decay product|daughter isotopes]] of ongoing [[radioactive decay]] processes such as [[alpha decay]], [[beta decay]], [[spontaneous fission]], [[cluster decay]], and other rarer modes of decay.

Of the 94 naturally occurring elements, those with atomic numbers 1 through 82 each have at least one [[stable isotope]] (except for [[technetium]], element 43 and [[promethium]], element 61, which have no stable isotopes). Isotopes considered stable are those for which no radioactive decay has yet been observed. Elements with atomic numbers 83 through 94 are [[Radionuclide|unstable]] enough that radioactive decay of all isotopes can be detected. Some of these elements, notably [[bismuth]] (atomic number 83), [[thorium]] (atomic number 90), and [[uranium]] (atomic number 92), have one or more isotopes with half-lives long enough to survive as remnants of the explosive [[stellar nucleosynthesis]] that produced the [[heavy metals]] before our [[Solar System]] formed. At 2{{e|19}} years, over 10{{sup|9}} times the estimated age of the universe, [[bismuth-209]] has the longest known alpha decay half-life of any nuclide, and is almost always considered on par with the 80 stable elements.<ref name="Dume2003">{{cite news|title=Bismuth breaks half-life record for alpha decay|last=Dumé|first=B.|date=23 April 2003|work=Physicsworld.com|publisher=Institute of Physics|location=Bristol, England|url=http://physicsworld.com/cws/article/news/2003/apr/23/bismuth-breaks-half-life-record-for-alpha-decay|access-date=14 July 2015|archive-date=13 December 2017|archive-url=https://web.archive.org/web/20171213214524/http://physicsworld.com/cws/article/news/2003/apr/23/bismuth-breaks-half-life-record-for-alpha-decay|url-status=live}}</ref><ref name="Marcillac2003">{{cite journal|last1=de Marcillac|first1=P.|last2=Coron|first2=N.|last3=Dambier|first3=G.|last4=Leblanc|first4=J.|last5=Moalic|first5=J-P|year=2003|title=Experimental detection of alpha-particles from the radioactive decay of natural bismuth|journal=Nature|volume=422|pages=876–878|doi=10.1038/nature01541|pmid=12712201|issue=6934|bibcode=2003Natur.422..876D|s2cid=4415582}}</ref> The heaviest elements (those beyond plutonium, element 94) are radioactive, with [[half-life|half-lives]] so short that they are not found in nature and must be [[synthetic element|synthesized]].

There are now 118 known elements. "Known" here means observed well enough, even from just a few decay products, to have been differentiated from other elements.<ref>{{cite journal|last=Sanderson|first=K.|date=17 October 2006|title=Heaviest element made – again|journal=News@nature|url=http://www.nature.com/news/2006/061016/full/061016-4.html|doi=10.1038/news061016-4|s2cid=121148847|access-date=8 March 2007|archive-date=16 May 2020|archive-url=https://web.archive.org/web/20200516072856/https://www.nature.com/news/2006/061016/full/061016-4.html|url-status=live|url-access=subscription}}</ref><ref name="Schewe">{{cite journal |last3=Castelvecchi |first3=D. |last1=Schewe|first1=P.|last2=Stein|first2=B.|date=October 16, 2006  |issue=797 |title=Elements 116 and 118 Are Discovered|url=http://www.aip.org/pnu/2006/797.html|journal =Physics News Update |publisher=American Institute of Physics|access-date=19 October 2006|url-status=dead|archive-url=https://web.archive.org/web/20120101144201/http://www.aip.org/pnu/2006/797.html|archive-date=1 January 2012}}</ref> Most recently, the synthesis of element 118 (since named [[oganesson]]) was reported in October 2006, and the synthesis of element 117 ([[tennessine]]) was reported in April 2010.<ref>{{cite news|last=Glanz|first=J.|date=6 April 2010|title=Scientists Discover Heavy New Element|url=https://www.nytimes.com/2010/04/07/science/07element.html |url-access=subscription |newspaper=The New York Times|access-date=15 February 2017|archive-date=19 June 2017|archive-url=https://web.archive.org/web/20170619122834/http://www.nytimes.com/2010/04/07/science/07element.html?hp|url-status=live}}</ref><ref>{{cite journal
|last1=Oganessian|first1=Yu. Ts.|date=April 2010|title=Synthesis of a New Element with Atomic Number Z=117 |bibcode-access=free |publisher=Physical Review Journals |journal=Physical Review Letters|volume=104|page=142502|doi=10.1103/PhysRevLett.104.142502 |last2=Abdullin|first2=F. Sh.|last3=Bailey|first3=P. D.|last4=Benker|first4=D. E.|last5=Bennett|first5=M. E.|last6=Dmitriev|first6=S. N.|last7=Ezold|first7=J. G.|last8=Hamilton|first8=J. H.|last9=Henderson
|first9=R. A.|last10=Itkis|first10=M. G.|last11=Lobanov|first11=Yu. V.|last12=Mezentsev|first12=A. N.|last13=Moody|first13=K. J.|last14=Nelson|first14=S. L.|last15=Polyakov|first15=A. N.|last16=Porter
|first16=C. E.|last17=Ramayya|first17=A. V.|last18=Riley|first18=F. D.|last19=Roberto|first19=J. B.|last20=Ryabinin|first20=M. A.|last21=Rykaczewski|first21=K. P.|last22=Sagaidak|first22=R. N. |last23=Shaughnessy|first23=D. A.|last24=Shirokovsky|first24=I. V.|last25=Stoyer|first25=M. A.|last26=Subbotin|first26=V. G.|last27=Sudowe|first27=R.|last28=Sukhov|first28=A. M.|last29=Tsyganov|first29=Yu. S.
|last30=Utyonkov|first30=V. K.|pmid=20481935|issue=14|bibcode=2010PhRvL.104n2502O|display-authors=29|doi-access=free}}</ref> Of these 118 elements, 94 occur naturally on Earth. Six of these occur in extreme trace amounts: [[technetium]], atomic number 43; [[promethium]], number 61; [[astatine]], number 85; [[francium]], number 87; [[neptunium]], number 93; and [[plutonium]], number 94. These 94 elements have been detected in the universe at large, in the spectra of stars and also supernovae, where short-lived radioactive elements are newly being made. The first 94 elements have been detected directly on Earth as [[primordial nuclide]]s present from the formation of the [[Solar System]], or as naturally occurring fission or transmutation products of uranium and thorium.

The remaining 24 heavier elements, not found today either on Earth or in astronomical spectra, have been produced artificially: all are radioactive, with short half-lives; if any of these elements were present when the Earth formed, they are certain to have completely decayed, and if present in novae, are in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element synthesized, in 1937, though traces of technetium have since been found in nature (and also the element may have been discovered naturally in 1925).<ref>{{citation-attribution|1={{cite web|url=http://www.epa.gov/radiation/radionuclides/technetium.html|title=Technetium-99|publisher=United States Environmental Protection Agency Radiation Protection |access-date=26 February 2013|archive-date=1 September 2015|archive-url=https://web.archive.org/web/20150901222619/http://www.epa.gov/radiation/radionuclides/technetium.html|url-status=dead }} }}</ref> This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring rare elements.<ref>{{cite web|url=https://www.cfa.harvard.edu/~ejchaisson/cosmic_evolution/docs/text/text_stel_6.html |work=Cosmic Evolution - From Big Bang to Humankind |first1= Eric J. |last1=Chaisson |title=Origins of Heavy Elements|publisher=Harvard–Smithsonian Center for Astrophysics|access-date=26 February 2013|archive-date=25 September 2020|archive-url=https://web.archive.org/web/20200925165732/https://www.cfa.harvard.edu/~ejchaisson/cosmic_evolution/docs/text/text_stel_6.html|url-status=live}}</ref>

[[List of chemical elements|Lists of elements]] are available by name, atomic number, density, melting point, boiling point and [[chemical symbol]], as well as [[Molar ionization energies of the elements|ionization energy]]. The nuclides of stable and radioactive elements are also available as a [[list of nuclides]], sorted by length of half-life for those that are unstable. One of the most convenient, and certainly the most traditional presentation of the elements, is in the form of the periodic table, which groups together elements with similar chemical properties (and usually also similar electronic structures).

=== Atomic number ===
{{main|Atomic number}}

The [[atomic number]] of an element is equal to the number of protons in each atom, and defines the element.<ref>{{cite web | url =http://www.ndt-ed.org/EducationResources/HighSchool/Radiography/atomicmassnumber.htm | title =Atomic Number and Mass Numbers | publisher =ndt-ed.org | access-date =17 February 2013 | archive-url =https://web.archive.org/web/20140212155836/http://www.ndt-ed.org/EducationResources/HighSchool/Radiography/atomicmassnumber.htm | archive-date =12 February 2014 | url-status =dead }}</ref> For example, all carbon atoms contain 6 protons in their [[atomic nucleus]]; so the atomic number of carbon is 6.<ref>{{cite web | url =http://periodic.lanl.gov/6.shtml | title =Periodic Table of Elements: LANL Carbon | author =periodic.lanl.gov | publisher =[[Los Alamos National Laboratory]] | access-date =17 February 2013 | archive-date =25 January 2021 | archive-url =https://web.archive.org/web/20210125032252/https://periodic.lanl.gov/6.shtml | url-status =live }}</ref> Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as [[isotope]]s of the element.<ref>{{cite web | url =http://faculty.piercecollege.edu/yamadak/classes/Atomic%20mass.pdf | title =Atomic mass, isotopes, and mass number. | author =Katsuya Yamada | publisher =[[Los Angeles Pierce College]] | url-status =dead | archive-url =https://web.archive.org/web/20140111131537/http://faculty.piercecollege.edu/yamadak/classes/Atomic%20mass.pdf | archive-date =11 January 2014 }}</ref>

The number of protons in the nucleus also determines its [[electric charge]], which in turn determines the number of [[electron]]s of the atom in its [[ionization|non-ionized]] state. The electrons are placed into [[atomic orbital]]s that determine the atom's [[chemical property|chemical properties]]. The number of neutrons in a nucleus usually has very little effect on an element's chemical properties; except for hydrogen (for which the [[kinetic isotope effect]] is significant). Thus, all carbon isotopes have nearly identical chemical properties because they all have six electrons, even though they may have 6 to 8 neutrons. That is why atomic number, rather than [[mass number]] or [[atomic weight]], is considered the identifying characteristic of an element.

The symbol for atomic number is ''Z''.

=== Isotopes ===
{{Main|Isotope|Stable isotope ratio|List of nuclides}}

[[Isotope]]s are atoms of the same element (that is, with the same number of [[proton]]s in their nucleus), but having ''different'' numbers of [[neutron]]s. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, said three isotopes are known as [[carbon-12]], [[carbon-13]], and [[carbon-14]] ({{sup|12}}C, {{sup|13}}C, and {{sup|14}}C). Natural carbon is a [[mixture]] of {{sup|12}}C (about 98.9%), {{sup|13}}C (about 1.1%) and about 1 atom per trillion of {{sup|14}}C.

Most (54 of 94) naturally occurring elements have more than one stable isotope. Except for the [[isotopes of hydrogen]] (which differ greatly from each other in relative mass—enough to cause chemical effects), the isotopes of a given element are chemically nearly indistinguishable.

All elements have radioactive isotopes (radioisotopes); most of these radioisotopes do not occur naturally. Radioisotopes typically decay into other elements via [[alpha decay]], [[beta decay]], or [[inverse beta decay]]; some isotopes of the heaviest elements also undergo [[spontaneous fission]]. Isotopes that are not radioactive, are termed "stable" isotopes. All known stable isotopes occur naturally (see [[primordial nuclide]]). The many radioisotopes that are not found in nature have been characterized after being artificially produced. Certain elements have no stable isotopes and are composed ''only'' of radioisotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic number greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one stable isotope. The mean number of stable isotopes for the 80 stable elements is 3.1 stable isotopes per element. The largest number of stable isotopes for a single element is 10 (for [[tin]], element 50).

=== Isotopic mass and atomic mass ===
{{main|atomic mass|relative atomic mass}}

The [[mass number]] of an element, ''A'', is the number of [[nucleon]]s (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass number, which is written as a superscript on the left hand side of the chemical symbol (e.g., {{sup|238}}U). The mass number is always an integer and has units of "nucleons". Thus, [[Isotopes of magnesium|magnesium-24]] (24 is the mass number) is an atom with 24 nucleons (12 protons and 12 neutrons).

Whereas the mass number simply counts the total number of neutrons and protons and is thus an integer, the [[atomic mass]] of a particular isotope (or "nuclide") of the element is the mass of a single atom of that isotope, and is typically expressed in [[Dalton (unit)|dalton]]s (symbol: Da), aka universal atomic mass units (symbol: u). Its [[relative atomic mass]] is a dimensionless number equal to the atomic mass divided by the [[atomic mass constant]], which equals 1&nbsp;Da. In general, the mass number of a given nuclide differs in value slightly from its relative atomic mass, since the mass of each proton and neutron is not exactly 1&nbsp;Da; since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number; and because of the [[nuclear binding energy]] and electron binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969&nbsp;Da and that of chlorine-37 is 36.966&nbsp;Da. However, the relative atomic mass of each isotope is quite close to its mass number (always within 1%). The only isotope whose atomic mass is exactly a [[natural number]] is {{sup|12}}C, which has a mass of 12&nbsp;Da; because the dalton is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.

The [[standard atomic weight]] (commonly called "atomic weight") of an element is the ''average'' of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit. This number may be a fraction that is ''not'' close to a whole number. For example, the relative atomic mass of chlorine is 35.453&nbsp;u, which differs greatly from a whole number as it is an average of about 76% chlorine-35 and 24% chlorine-37. Whenever a relative atomic mass value differs by more than ~1% from a whole number, it is due to this averaging effect, as significant amounts of more than one isotope are naturally present in a sample of that element.

=== Chemically pure and isotopically pure ===
Chemists and nuclear scientists have different definitions of a ''pure element''. In chemistry, a pure element means a substance whose atoms all (or in practice almost all) have the same atomic number, or number of [[proton]]s. Nuclear scientists, however, define a pure element as one that consists of only one isotope.<ref>{{cite news |url=http://www.euronuclear.org/info/encyclopedia/p/pure-element.htm |title=Pure element |publisher=[[European Nuclear Society]] |access-date=13 August 2013 |archive-url=https://web.archive.org/web/20170613073021/http://www.euronuclear.org/info/encyclopedia/p/pure-element.htm |archive-date=13 June 2017 |url-status=dead }}</ref>

For example, a [[copper]] wire is 99.99% chemically pure if 99.99% of its atoms are copper, with 29 protons each. However it is not isotopically pure since natural copper consists of two stable isotopes, 69% {{sup|63}}Cu and 31% {{sup|65}}Cu, with different numbers of neutrons. (See [[Isotopes of copper]].) However, pure gold would be both chemically and isotopically pure, since ordinary gold consists only of one isotope, {{sup|197}}Au.

=== Allotropes ===
{{Main|Allotropy}}

Atoms of chemically pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple [[chemical structure]]s ([[Molecular geometry|spatial arrangements of atoms]]), known as [[allotrope]]s, which differ in their properties. For example, carbon can be found as [[diamond]], which has a tetrahedral structure around each carbon atom; [[graphite]], which has layers of carbon atoms with a hexagonal structure stacked on top of each other; [[graphene]], which is a single layer of graphite that is very strong; [[fullerene]]s, which have nearly spherical shapes; and [[carbon nanotube]]s, which are tubes with a hexagonal structure (even these may differ from each other in electrical properties). The ability of an element to exist in one of many structural forms is known as 'allotropy'.

The reference state of an element is defined by convention, usually as the thermodynamically most stable allotrope and physical state at a pressure of 1 [[bar (unit)|bar]] and a given temperature (typically 298.15[[Kelvin|K]]). However, for phosphorus, the reference state is white phosphorus even though it is not the most stable allotrope, and the reference state for carbon is graphite, because the structure of graphite is more stable than that of the other allotropes. In [[thermochemistry]], an element is defined to have an [[Standard enthalpy of formation|enthalpy of formation]] of zero in its reference state. 

=== Properties ===
Several kinds of descriptive categorisations can be applied broadly to the elements, including consideration of their general physical and chemical properties, their states of matter under familiar conditions, their melting and boiling points, their densities, their crystal structures as solids, and their origins.

==== General properties ====
Several terms are commonly used to characterise the general physical and chemical properties of the chemical elements. A first distinction is between [[metal]]s, which readily conduct [[electricity]], [[nonmetal]]s, which do not, and a small group, (the ''[[metalloid]]s''), having intermediate properties and often behaving as [[semiconductor]]s.

A more refined classification is often shown in coloured presentations of the periodic table. This system restricts the terms "metal" and "nonmetal" to only certain of the more broadly defined metals and nonmetals, adding additional terms for certain sets of the more broadly viewed metals and nonmetals. The version of this classification used in the periodic tables presented here includes: [[actinide]]s, [[alkali metal]]s, [[alkaline earth metal]]s, [[halogen]]s, [[lanthanide]]s, [[transition metal]]s, [[post-transition metal]]s, [[metalloid]]s, [[reactive nonmetal]]s, and [[noble gas]]es. In this system, the alkali metals, alkaline earth metals, and transition metals, as well as the lanthanides and the actinides, are special groups of the metals viewed in a broader sense. Similarly, the reactive nonmetals and the noble gases are nonmetals viewed in the broader sense. In some presentations, the halogens are not distinguished, with [[astatine]] identified as a metalloid and the others identified as nonmetals.

==== States of matter ====
Another commonly used basic distinction among the elements is their [[state of matter]] (phase), whether [[solid]], [[liquid]], or [[gas]], at [[standard temperature and pressure]] (STP). Most elements are solids at STP, while several are gases. Only [[bromine]] and [[mercury (element)|mercury]] are liquid at 0 degrees Celsius (32 degrees Fahrenheit) and 1 atmosphere pressure; [[caesium]] and [[gallium]] are solid at that temperature, but melt at 28.4°C (83.2°F) and 29.8°C (85.6°F), respectively.

==== Melting and boiling points ====
[[Melting point|Melting]] and [[boiling point]]s, typically expressed in degrees [[Celsius]] at a pressure of one atmosphere, are commonly used in characterizing the various elements. While known for most elements, either or both of these measurements is still undetermined for some of the radioactive elements available in only tiny quantities. Since helium remains a liquid even at [[absolute zero]] at atmospheric pressure, it has only a boiling point, and not a melting point, in conventional presentations.

==== Densities ====
{{Main|Densities of the elements (data page)}}

The [[density]] at selected [[standard temperature and pressure]] (STP) is often used in characterizing the elements. Density is often expressed in grams per cubic centimetre (g/cm{{sup|3}}). Since several elements are gases at commonly encountered temperatures, their densities are usually stated for their gaseous forms; when liquefied or solidified, the gaseous elements have densities similar to those of the other elements.

When an element has allotropes with different densities, one representative allotrope is typically selected in summary presentations, while densities for each allotrope can be stated where more detail is provided. For example, the three familiar [[allotropes of carbon]] ([[amorphous carbon]], [[graphite]], and [[diamond]]) have densities of 1.8–2.1, 2.267, and 3.515&nbsp;g/cm{{sup|3}}, respectively.

==== Crystal structures ====
{{Main|Crystal structure}}

The elements studied to date as solid samples have eight kinds of [[crystal structure]]s: [[cubic crystal system|cubic]], [[cubic crystal system|body-centered cubic]], face-centered cubic, [[Hexagonal crystal system|hexagonal]], [[Monoclinic crystal system|monoclinic]], [[orthorhombic crystal system|orthorhombic]], [[Trigonal crystal system|rhombohedral]], and [[Tetragonal crystal system|tetragonal]]. For some of the synthetically produced transuranic elements, available samples have been too small to determine crystal structures.

==== Occurrence and origin on Earth ====
{{Main|Abundance of elements in Earth's crust}}
Chemical elements may also be categorised by their origin on Earth, with the first 94 considered naturally occurring, while those with atomic numbers beyond 94 have only been produced artificially via human-made nuclear reactions.

Of the 94 naturally occurring elements, 83 are considered primordial and either [[stable isotope|stable]] or weakly radioactive. The longest-lived isotopes of the remaining 11 elements have [[Half-life|half lives]] too short for them to have been present at the beginning of the Solar System, and are therefore "transient elements". Of these 11 transient elements, five ([[polonium]], [[radon]], [[radium]], [[actinium]], and [[protactinium]]) are relatively common [[decay product]]s of [[thorium]] and [[uranium]]. The remaining six transient elements (technetium, promethium, astatine, [[francium]], [[neptunium]], and [[plutonium]]) occur only rarely, as products of rare decay modes or nuclear reaction processes involving uranium or other heavy elements.

Elements with atomic numbers 1 through 82, except 43 (technetium) and 61 (promethium), each have at least one isotope for which no radioactive decay has been observed. Observationally stable isotopes of some elements (such as [[tungsten]] and [[lead]]), however, are predicted to be slightly radioactive with very long half-lives:{{NUBASE2016|ref}} for example, the half-lives predicted for the observationally stable lead isotopes range from 10{{sup|35}} to 10{{sup|189}} years. Elements with atomic numbers 43, 61, and 83 through 94 are unstable enough that their radioactive decay can be detected. Three of these elements, bismuth (element 83), thorium (90), and uranium (92) have one or more isotopes with half-lives long enough to survive as remnants of the explosive [[stellar nucleosynthesis]] that produced the heavy elements before the formation of the Solar System. For example, at over 1.9{{e|19}} years, over a billion times longer than the estimated age of the universe, [[bismuth-209]] has the longest known [[alpha decay]] half-life of any isotope.{{r|Dume2003}}{{r|Marcillac2003}} The last 24 elements (those beyond plutonium, element 94) undergo radioactive decay with short half-lives and cannot be produced as daughters of longer-lived elements, and thus are not known to occur in nature at all.

=== Periodic table ===
{{Main|Periodic table}}
{{Periodic table}}

The properties of the elements are often summarized using the periodic table, which powerfully and elegantly organizes the elements by increasing atomic number into rows ([[period (periodic table)|"periods"]]) in which the columns ([[group (periodic table)|"groups"]]) share recurring ("periodic") physical and chemical properties. The table contains 118 confirmed elements as of 2021.

Though earlier precursors to this presentation exist, its invention is generally credited to Russian chemist [[Dmitri Mendeleev]] in 1869, who intended the table to illustrate recurring trends in the properties of the elements. The layout of the table has been refined and extended over time as new elements have been discovered and new theoretical models have been developed to explain chemical behavior.

Use of the periodic table is now ubiquitous in chemistry, providing an extremely useful framework to classify, systematize and compare all the many different forms of chemical behavior. The table has also found wide application in [[physics]], [[geology]], [[biology]], [[materials science]], [[engineering]], [[agriculture]], [[medicine]], [[nutrition]], [[environmental health]], and [[astronomy]]. Its principles are especially important in [[chemical engineering]].