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{{Short description|Chemical substance not composed of simpler ones}}
{{pp-vandalism|small=yes}}
{{Use British English|date=October 2021}}
{{Use dmy dates|date=August 2021}}
[[File:Periodic table (32-col, enwiki), black and white.png|thumb|upright=1.33|The chemical elements ordered in the [[periodic table]]]]
{{Sidebar periodic table}}
A '''chemical element''' is a [[chemical substance]] whose [[atom]]s all have the same number of [[proton]]s. The number of protons is called the [[atomic number]] of that element. For example, oxygen has an atomic number of 8, meaning each oxygen atom has 8 protons in its nucleus. Atoms of the same element can have different numbers of [[neutron]]s in their nuclei, known as [[isotope]]s of the element. Two or more atoms can combine to form [[molecule]]s. Some elements are formed from [[Homonuclear molecule|molecules of identical atoms]], e. g. atoms of hydrogen (H) form diatomic molecules (H<sub>2</sub>). [[Chemical compound]]s are substances made of atoms of different elements; they can have molecular or non-molecular structure. [[Mixture]]s are materials containing different chemical substances; that means (in case of molecular substances) that they contain different types of molecules. Atoms of one element can be transformed into atoms of a different element in [[nuclear reaction]]s, which change an atom's atomic number.

Historically, the term "chemical element" meant a substance that cannot be broken down into constituent substances by chemical reactions, and for most practical purposes this definition still has validity. There was some controversy in the 1920s over whether isotopes deserved to be recognized as separate elements if they could be separated by chemical means.<ref>Helge Kragh (2000). Conceptual Changes in Chemistry: The Notion of a Chemical Element, ca. 1900-1925</ref>

The term "(chemical) element" is used in two different but closely related meanings:<ref name="iupac">{{cite journal |last1=Chemistry (IUPAC) |first1=The International Union of Pure and Applied |title=IUPAC – chemical element (C01022) |url=https://goldbook.iupac.org/terms/view/C01022 |website=goldbook.iupac.org|doi=10.1351/goldbook.C01022 |doi-access=free }}</ref> it can mean a chemical substance consisting of a single kind of atoms, or it can mean that kind of atoms as a component of various chemical substances. For example, molecules of water (H<sub>2</sub>O) contain atoms of hydrogen (H) and oxygen (O), so water can be said as a compound consisting of the elements hydrogen (H) and oxygen (O) even though it does not contain the chemical substances (di)hydrogen (H<sub>2</sub>) and (di)oxygen (O<sub>2</sub>), as H<sub>2</sub>O molecules are different from H<sub>2</sub> and O<sub>2</sub> molecules. For the meaning "chemical substance consisting of a single kind of atoms", the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in English chemical literature, whereas in some other languages their equivalent is widely used. For example, the French chemical terminology distinguishes {{lang|fr|élément chimique}} (kind of atoms) and {{lang|fr|corps simple}} (chemical substance consisting of a single kind of atoms); the Russian chemical terminology distinguishes {{lang|ru|химический элемент}} and {{lang|ru|простое вещество}}.

Almost all [[baryonic matter]] in the universe is composed of elements (among rare exceptions are [[neutron star]]s). When different elements undergo chemical reactions, atoms are rearranged into new compounds held together by [[chemical bond]]s. Only a few elements, such as [[silver]] and [[gold]], are found uncombined as relatively pure [[native element mineral]]s. Nearly all other naturally occurring elements occur in the [[Earth]] as compounds or mixtures. [[Atmosphere of Earth|Air]] is mostly a mixture of molecular [[nitrogen]] and [[oxygen]], though it does contain compounds including [[carbon dioxide]] and [[water]], as well as atomic [[argon]], a [[noble gas]] which is [[chemically inert]] and therefore does not undergo chemical reactions.

The history of the discovery and use of elements began with early [[society|human societies]] that discovered native minerals like [[carbon]], [[sulfur]], [[copper]] and gold (though the modern concept of an element was not yet understood). Attempts to classify materials such as these resulted in the concepts of [[classical element]]s, [[alchemy]], and similar theories throughout history. Much of the modern understanding of elements developed from the work of [[Dmitri Mendeleev]], a Russian chemist who published the first recognizable [[periodic table]] in 1869. This table 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 property|physical]] and [[chemical property|chemical properties]]. The periodic table summarizes various properties of the elements, allowing chemists to derive relationships between them and to make predictions about elements not yet discovered, and potential new compounds.

By November 2016, the [[International Union of Pure and Applied Chemistry]] (IUPAC) had recognized a total of 118 elements. The first 94 occur naturally on [[Earth]], and the remaining 24 are [[synthetic element]]s produced in nuclear reactions. Save for unstable radioactive elements (radioelements) which [[radioactive decay|decay]] quickly, nearly all elements are available industrially in varying amounts. The [[timeline of chemical element discoveries|discovery and synthesis of further new elements]] is an ongoing area of scientific study.

== 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]] to the point 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 the formation of our [[Solar System]]. 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, 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) undergo radioactive decay 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. In this context, "known" 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}}</ref><ref name="Schewe">{{cite web|last1=Schewe|first1=P.|last2=Stein|first2=B.|date=17 October 2000|title=Elements 116 and 118 Are Discovered|url=http://www.aip.org/pnu/2006/797.html|work=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?hp|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|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 quantities: [[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 at the formation of Earth, 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 trace amounts 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|website=epa.gov|publisher=United States Environmental Protection Agency|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=live}} }}</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|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|List of the 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), or [[universal atomic mass unit]]s (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 ordinary copper consists of two stable isotopes, 69% {{sup|63}}Cu and 31% {{sup|65}}Cu, with different numbers of neutrons. 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 at 298.15K). 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 categorizations 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 characterize 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 colored 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 categorized 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 considered 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.

Although 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]].

== Nomenclature and symbols ==
{{anchor|Nomenclature}}

The various chemical elements are formally identified by their unique atomic numbers, their accepted names, and their [[chemical symbol]]s.

=== Atomic numbers ===
The known elements have atomic numbers from 1 to 118, conventionally presented as [[Arabic numerals]]. Since the elements can be uniquely sequenced by atomic number, conventionally from lowest to highest (as in a periodic table), sets of elements are sometimes specified by such notation as "through", "beyond", or "from ... through", as in "through iron", "beyond uranium", or "from lanthanum through lutetium". The terms "light" and "heavy" are sometimes also used informally to indicate relative atomic numbers (not densities), as in "lighter than carbon" or "heavier than lead", though the atomic masses of the elements (their atomic weights or atomic masses) do not always increase [[monotonic function|monotonically]] with their atomic numbers.

=== Element names ===
{{main|Naming of chemical elements}}

The naming of various substances now known as elements precedes the [[atomic theory]] of matter, as names were given locally by various cultures to various minerals, metals, compounds, alloys, mixtures, and other materials, though at the time it was not known which chemicals were elements and which compounds. As they were identified as elements, the existing names for anciently known elements (e.g., gold, mercury, iron) were kept in most countries. National differences emerged over the element names either for convenience, linguistic niceties, or nationalism. For example, German speakers use "Wasserstoff" (water substance) for "hydrogen", "Sauerstoff" (acid substance) for "oxygen" and "Stickstoff" (smothering substance) for "nitrogen"; English and some other languages use "sodium" for "natrium", and "potassium" for "kalium"; and the French, Italians, Greeks, Portuguese and Poles prefer "azote/azot/azoto" (from roots meaning "no life") for "nitrogen".

For purposes of international communication and trade, the [[Table of chemical elements|official names of the chemical elements]] both ancient and more recently recognized are decided by the [[International Union of Pure and Applied Chemistry]] (IUPAC), which has decided on a sort of international English language, drawing on traditional English names even when an element's chemical symbol is based on a Latin or other traditional word, for example adopting "gold" rather than "aurum" as the name for the 79th element (Au). IUPAC prefers the British spellings "[[aluminium]]" and "caesium" over the U.S. spellings "aluminum" and "cesium", and the U.S. "sulfur" over British "sulphur". However, elements that are practical to sell in bulk in many countries often still have locally used national names, and countries whose national language does not use the [[Latin alphabet]] are likely to use the IUPAC element names.

According to IUPAC, element names are not proper nouns; therefore, the full name of an element is not capitalized in English, even if derived from a [[proper noun]], as in [[californium]] and [[einsteinium]]. Isotope names are also uncapitalized if written out, ''e.g.,'' [[carbon-12]] or [[uranium-235]]. Chemical element ''symbols'' (such as Cf for californium and Es for einsteinium), are always capitalized (see below).

In the second half of the 20th century, physics laboratories became able to produce elements with half-lives too short for an appreciable amount of them to exist at any time. These are also named by IUPAC, which generally adopts the name chosen by the discoverer. This practice can lead to the controversial question of which research group actually discovered an element, a question that delayed the naming of elements with atomic number of 104 and higher for a considerable amount of time. (See [[element naming controversy]]).

Precursors of such controversies involved the nationalistic namings of elements in the late 19th century. For example, ''[[lutetium]]'' was named in reference to Paris, France. The Germans were reluctant to relinquish naming rights to the French, often calling it ''cassiopeium''. Similarly, the British discoverer of ''[[niobium]]'' originally named it ''columbium'', in reference to the [[New World]]. It was used extensively as such by American publications before the international standardization (in 1950).

=== Chemical symbols ===
{{For|lists of current chemical symbols, symbols not currently used, and other symbols that may look like chemical symbols|Chemical symbol}}

==== Specific elements ====
Before chemistry became a [[science]], [[alchemy|alchemists]] designed arcane symbols for both metals and common compounds. These were however used as abbreviations in diagrams or procedures; there was no concept of atoms combining to form [[molecule]]s. With his advances in the atomic theory of matter, [[John Dalton]] devised his own simpler symbols, based on circles, to depict molecules.

The current system of chemical notation was invented by [[Jöns Jacob Berzelius]] in 1814. In this system, chemical symbols are not mere abbreviations—though each consists of letters of the [[Latin alphabet]]. They are intended as universal symbols for people of all languages and alphabets.

Since Latin was the common language of science at Berzelius' time, his symbols were abbreviations based on the [[Latin]] names of elements (they may be Classical Latin names of elements known since antiquity or [[Neo-Latin]] coinages for later elements). The symbols are not followed by a period (full stop) as with abbreviations. In most cases, Latin names of elements as used by Berzelius have the same roots as the modern English name. For example, [[hydrogen]] has the symbol "H" from Neo-Latin {{lang|la|hydrogenium}}, which has the same Greek roots as English ''hydrogen''. However, in eleven cases Latin (as used by Berzelius) and English names of elements have different roots. Eight of them are the seven [[metals of antiquity]] and a metalloid also known since antiquity: "Fe" (Latin {{lang|la|ferrum}}) for [[iron]], "Hg" (Latin {{lang|la|hydrargyrum}}) for [[Mercury (element)|mercury]], "Sn" (Latin {{lang|la|stannum}}) for [[tin]], "Au" (Latin {{lang|la|aurum}}) for gold, "Ag" (Latin {{lang|la|argentum}}) for [[silver]], "Pb" (Latin {{lang|la|plumbum}}) for [[lead]], "Cu" (Latin {{lang|la|cuprum}}) for [[copper]], and "Sb" (Latin {{lang|la|stibium}}) for [[antimony]]. The three other mismatches between Neo-Latin (as used by Berzelius) and English names are "Na" (Neo-Latin {{lang|la|natrium}}) for [[sodium]], "K" (Neo-Latin {{lang|la|kalium}}) for [[potassium]], and "W" (Neo-Latin {{lang|la|wolframium}}) for [[tungsten]]. These mismatches came from different suggestings of naming the elements in the [[Modern era]]. Initially Berzelius had suggested "So" and "Po" for sodium and potassium, but he changed the symbols to "Na" and "K" later in the same year.

Elements discovered after 1814 were also assigned unique chemical symbols, based on the name of the element. The use of Latin as the universal language of science was fading, but chemical names of newly discovered elements came to be borrowed from language to language with little or no modifications. Symbols of elements discovered after 1814 match their names in English, French (ignoring the [[acute accent]] on {{angle bracket|é}}), and German (though German often allows alternate spellings with {{angle bracket|k}} or {{angle bracket|z}} instead of {{angle bracket|c}}: e.g., the name of [[calcium]] may be spelled {{lang|de|Calcium}} or {{lang|de|Kalzium}} in German, but its symbol is always "Ca"). Other languages sometimes modify element name spellings: Spanish {{lang|es|iterbio}} (ytterbium), Italian {{lang|it|afnio}} (hafnium), Swedish {{lang|sv|moskovium}} (moscovium); but those modifications do not affect chemical symbols: Yb, Hf, Mc.

Chemical symbols are understood internationally when element names might require translation. There have been some differences in the past. For example, Germans in the past have used "J" (for the name {{lang|de|Jod}}) for iodine, but now use "I" and {{lang|de|Iod}}.

The first letter of a chemical symbol is always capitalized, as in the preceding examples, and the subsequent letters, if any, are always lower case. Thus, the symbols for californium and einsteinium are Cf and Es.

==== General chemical symbols ====
There are also symbols in chemical equations for groups of elements, for example in comparative formulas. These are often a single capital letter, and the letters are reserved and not used for names of specific elements. For example, "'''X'''" indicates a variable group (usually a halogen) in a class of compounds, while "'''R'''" is a [[Radical (chemistry)|radical]], meaning a compound structure such as a hydrocarbon chain. The letter "'''Q'''" is reserved for "heat" in a chemical reaction. "'''Y'''" is also often used as a general chemical symbol, though it is also the symbol of [[yttrium]]. "'''Z'''" is also often used as a general variable group. "'''E'''" is used in organic chemistry to denote an [[Polar effect|electron-withdrawing group]] or an [[electrophile]]; similarly "'''Nu'''" denotes a [[nucleophile]]. "'''L'''" is used to represent a general [[ligand]] in [[inorganic chemistry|inorganic]] and [[organometallic chemistry]]. "'''M'''" is also often used in place of a general metal.

At least two other, two-letter generic chemical symbols are also in informal use, "'''Ln'''" for any [[lanthanide]] and "'''An'''" for any [[actinide]]. "'''Rg'''" was formerly used for any [[rare gas]] element, but the group of rare gases has now been renamed [[noble gas]]es and "'''Rg'''" now refers to [[roentgenium]].

==== Isotope symbols ====
Isotopes of an element are distinguished by mass number (total protons and neutrons), with this number combined with the element's symbol. IUPAC prefers that isotope symbols be written in superscript notation when practical, for example {{sup|12}}C and {{sup|235}}U. However, other notations, such as carbon-12 and uranium-235, or C-12 and U-235, are also used.

As a special case, the three naturally occurring isotopes of hydrogen are often specified as '''H''' for {{sup|1}}H ([[protium]]), '''D''' for {{sup|2}}H ([[deuterium]]), and '''T''' for {{sup|3}}H ([[tritium]]). This convention is easier to use in chemical equations, replacing the need to write out the mass number each time. Thus, the formula for [[heavy water]] may be written D{{sub|2}}O instead of {{sup|2}}H{{sub|2}}O.

== Origin of the elements ==
{{more citations needed|section|date=April 2021}}
{{main|Nucleosynthesis}}
[[File:Universe content pie chart.jpg|thumb|upright=1.7|Estimated distribution of dark matter and dark energy in the universe. Only the fraction of the mass and energy labeled "atoms" is composed of elements.]]
Only about 4% of the total mass of the universe is made of atoms or [[ion]]s, and thus represented by elements. This fraction is about 15% of the total matter, with the remainder of the matter (85%) being [[dark matter]]. The nature of dark matter is unknown, but it is not composed of atoms of elements because it contains no protons, neutrons, or electrons. (The remaining non-matter part of the mass of the universe is composed of the even less well understood [[dark energy]]).

The 94 naturally occurring elements were produced by at least four classes of astrophysical process. Most of the hydrogen, helium and a very small quantity of lithium were produced in the first few minutes of the [[Big Bang]]. This [[Big Bang nucleosynthesis]] happened only once; the other processes are ongoing. [[Nuclear fusion]] inside stars produces elements through stellar nucleosynthesis, including all elements from carbon to [[iron]] in atomic number. Elements higher in atomic number than iron, including heavy elements like uranium and plutonium, are produced by various forms of explosive nucleosynthesis in [[supernova]]e and [[neutron star merger]]s. The light elements [[lithium]], [[beryllium]] and [[boron]] are produced mostly through [[cosmic ray spallation]] (fragmentation induced by [[cosmic ray]]s) of carbon, nitrogen, and oxygen.

In the early phases of the Big Bang, nucleosynthesis of hydrogen resulted in the production of hydrogen-1 (protium, {{sup|1}}H) and helium-4 ({{sup|4}}He), as well as a smaller amount of deuterium ({{sup|2}}H) and tiny amounts (on the order of 10{{sup|−10}}) of lithium and beryllium. Even smaller amounts of boron may have been produced in the Big Bang, since it has been observed in some very old stars, while carbon has not.<ref>{{cite news|last=Wilford|first=J.N.|date=14 January 1992|title=Hubble Observations Bring Some Surprises|url=https://query.nytimes.com/gst/fullpage.html?res=9E0CE5D91F3AF937A25752C0A964958260|newspaper=[[The New York Times]]|access-date=15 February 2017|archive-date=5 March 2008|archive-url=https://web.archive.org/web/20080305211140/http://query.nytimes.com/gst/fullpage.html?res=9E0CE5D91F3AF937A25752C0A964958260|url-status=live}}</ref> No elements heavier than boron were produced in the Big Bang. As a result, the primordial abundance of atoms (or ions) consisted of ~75% {{sup|1}}H, 25% {{sup|4}}He, and 0.01% deuterium, with only tiny traces of lithium, beryllium, and perhaps boron.<ref>{{cite web|last=Wright|first=E. L.|date=12 September 2004|title=Big Bang Nucleosynthesis|url=http://www.astro.ucla.edu/~wright/BBNS.html|publisher=[[UCLA]], Division of Astronomy|access-date=22 February 2007|archive-date=13 January 2018|archive-url=https://web.archive.org/web/20180113051655/http://astro.ucla.edu/~wright/BBNS.html|url-status=live}}</ref> Subsequent enrichment of [[galactic spheroid|galactic halos]] occurred due to stellar nucleosynthesis and [[supernova nucleosynthesis]].<ref name="synthesis">{{cite journal|year=1999 |title=Synthesis of the elements in stars: forty years of progress |url=http://www.cococubed.com/papers/wallerstein97.pdf |journal=[[Reviews of Modern Physics]] |volume=69 |issue=4 |pages=995–1084 |doi=10.1103/RevModPhys.69.995 |bibcode=1997RvMP...69..995W |last1=Wallerstein |first1=George |last2=Iben |first2=Icko |last3=Parker |first3=Peter |last4=Boesgaard |first4=Ann |last5=Hale |first5=Gerald |last6=Champagne |first6=Arthur |last7=Barnes |first7=Charles |last8=Käppeler |first8=Franz |last9=Smith |first9=Verne |display-authors=8 |url-status=dead |archive-url=https://web.archive.org/web/20060928043229/http://www.cococubed.com/papers/wallerstein97.pdf |archive-date=28 September 2006 |hdl=2152/61093 |hdl-access=free }}</ref> However, the element abundance in [[intergalactic space]] can still closely resemble primordial conditions, unless it has been enriched by some means.

[[File:Nucleosynthesis periodic table.svg|thumb|upright=1.78|Periodic table showing the cosmogenic origin of each element in the Big Bang, or in large or small stars. Small stars can produce certain elements up to sulfur, by the [[alpha process]]. Supernovae are needed to produce "heavy" elements (those beyond iron and nickel) rapidly by neutron buildup, in the [[r-process]]. Certain large stars slowly produce other elements heavier than iron, in the [[s-process]]; these may then be blown into space in the off-gassing of [[planetary nebulae]]]]

On Earth (and elsewhere), trace amounts of various elements continue to be produced from other elements as products of [[nuclear transmutation]] processes. These include some produced by [[cosmic ray]]s or other nuclear reactions (see [[cosmogenic]] and [[nucleogenic]] nuclides), and others produced as [[decay product]]s of long-lived primordial nuclides.<ref name="Earnshaw1997">{{cite book|last1=Earnshaw|first1=A.|last2=Greenwood|first2=N.|year=1997|title=Chemistry of the Elements|edition=2nd|publisher=[[Butterworth-Heinemann]]}}</ref> For example, trace (but detectable) amounts of [[carbon-14]] ({{sup|14}}C) are continually produced in the air by cosmic rays impacting nitrogen atoms, and argon-40 ({{sup|40}}Ar) is continually produced by the decay of primordially occurring but unstable potassium-40 ({{sup|40}}K). Also, three primordially occurring but radioactive actinides, thorium, uranium, and plutonium, decay through a series of recurrently produced but unstable elements such as radium and [[radon]], which are transiently present in any sample of containing these metals. Three other radioactive elements, technetium, promethium, and neptunium, occur only incidentally in natural materials, produced as individual atoms by [[nuclear fission]] of the nuclei of various heavy elements or in other rare nuclear processes.

Besides the 94 naturally occurring elements, several [[artificial element]]s have been produced by [[nuclear physics]] technology. By 2016, these experiments had produced all elements up to atomic number 118.

== Abundance ==
{{Main|Abundance of the chemical elements}}

The following graph (note log scale) shows the abundance of elements in our Solar System. The table shows the 12 most common elements in our galaxy (estimated spectroscopically), as measured in [[parts per million]] by [[mass]].<ref name="croswell">{{cite book |last=Croswell |first=Ken |author-link=Ken Croswell |year=1996 |title=Alchemy of the Heavens |url=http://kencroswell.com/alchemy.html |publisher=Anchor |isbn=978-0-385-47214-2 |access-date=10 October 2007 |archive-date=13 May 2011 |archive-url=https://web.archive.org/web/20110513233910/http://www.kencroswell.com/alchemy.html |url-status=live }}</ref> Nearby galaxies that have evolved along similar lines have a corresponding enrichment of elements heavier than hydrogen and helium. The more distant galaxies are being viewed as they appeared in the past, so their abundances of elements appear closer to the primordial mixture. As physical laws and processes appear common throughout the [[visible universe]], however, scientists expect that these galaxies evolved elements in similar abundance.

The abundance of elements in the Solar System is in keeping with their origin from nucleosynthesis in the [[Big Bang]] and a number of progenitor supernova stars. Very abundant hydrogen and helium are products of the Big Bang, but the next three elements are rare since they had little time to form in the Big Bang and are not made in stars (they are, however, produced in small quantities by the breakup of heavier elements in interstellar dust, as a result of impact by cosmic rays). Beginning with carbon, elements are produced in stars by buildup from alpha particles (helium nuclei), resulting in an alternatingly larger abundance of elements with even atomic numbers (these are also more stable). In general, such elements up to iron are made in large stars in the process of becoming [[supernova]]s. Iron-56 is particularly common, since it is the most stable nuclide that can easily be made from alpha particles (being a product of decay of radioactive nickel-56, ultimately made from 14 helium nuclei). Elements heavier than iron are made in energy-absorbing processes in large stars, and their abundance in the universe (and on Earth) generally decreases with their atomic number.

The [[abundance of the chemical elements]] on Earth varies from air to crust to ocean, and in various types of life. The abundance of elements in Earth's crust differs from that in the Solar System (as seen in the Sun and massive planets like Jupiter) mainly in selective loss of the very lightest elements (hydrogen and helium) and also volatile neon, carbon (as hydrocarbons), nitrogen and sulfur, as a result of solar heating in the early formation of the Solar System. Oxygen, the most abundant Earth element by mass, is retained on Earth by combination with silicon. Aluminium at 8% by mass is more common in the Earth's crust than in the universe and solar system, but the composition of the far more bulky mantle, which has magnesium and iron in place of aluminium (which occurs there only at 2% of mass) more closely mirrors the elemental composition of the solar system, save for the noted loss of volatile elements to space, and loss of iron which has migrated to the Earth's core.

The [[composition of the human body]], by contrast, more closely follows the composition of [[seawater]]—save that the human body has additional stores of carbon and nitrogen necessary to form the [[protein]]s and [[nucleic acid]]s, together with [[phosphorus]] in the nucleic acids and energy transfer molecule [[adenosine triphosphate]] (ATP) that occurs in the cells of all living organisms. Certain kinds of [[organism]]s require particular additional elements, for example the [[magnesium]] in [[chlorophyll]] in green plants, the [[calcium]] in [[mollusc shell]]s, or the iron in the [[hemoglobin]] in [[vertebrate]]s' [[red blood cell]]s.

[[File:SolarSystemAbundances.svg|thumb|upright=2.7|Abundances of the chemical elements in the Solar System. Hydrogen and helium are most common, from the Big Bang. The next three elements (Li, Be, B) are rare because they are poorly synthesized in the Big Bang and also in stars. The two general trends in the remaining stellar-produced elements are: (1) an alternation of abundance in elements as they have even or odd atomic numbers (the [[Oddo–Harkins rule]]), and (2) a general decrease in abundance as elements become heavier. Iron is especially common because it represents the minimum energy nuclide that can be made by fusion of helium in supernovae.]]

{|class="wikitable"
|-
!Elements in our galaxy!!Parts per million<br />by mass
|-
|[[Hydrogen]]
|739,000
|-
|[[Helium]]
|240,000
|-
|[[Oxygen]]
|10,400
|-
|[[Carbon]]
|4,600
|-
|[[Neon]]
|1,340
|-
|[[Iron]]
|1,090
|-
|[[Nitrogen]]
|960
|-
|[[Silicon]]
|650
|-
|[[Magnesium]]
|580
|-
|[[Sulfur]]
|440
|-
|[[Potassium]]
|210
|-
|[[Nickel]]
|100
|}
{{clear}}

{{Periodic table (dietary elements)}}

== History<!--linked from 'Classical element'--> ==
=== Evolving definitions ===
The concept of an "element" as an indivisible substance has developed through three major historical phases: Classical definitions (such as those of the ancient Greeks), chemical definitions, and atomic definitions.

=== Classical definitions ===
{{Main|Classical element#Hellenistic philosophy}}
[[Ancient philosophy]] posited a set of [[classical element]]s to explain observed patterns in [[nature]]. These ''elements'' originally referred to ''[[Earth (classical element)|earth]]'', ''[[Water (classical element)|water]]'', ''[[Air (classical element)|air]]'' and ''[[Fire (classical element)|fire]]'' rather than the chemical elements of modern science.

The term 'elements' (''stoicheia'') was first used by Greek philosopher [[Plato]] around 360&nbsp;BCE in his dialogue [[Timaeus (dialogue)|Timaeus]], which includes a discussion of the composition of inorganic and organic bodies and is a speculative treatise on chemistry. Plato believed the elements introduced a century earlier by [[Empedocles]] were composed of small [[regular polyhedron|polyhedral]] [[Theory of Forms|forms]]: [[tetrahedron]] (fire), [[octahedron]] (air), [[icosahedron]] (water), and [[cube]] (earth).<ref>{{cite book|author=Plato|year=2008|orig-year=c. 360 BC|title=Timaeus|url=https://books.google.com/books?id=xSjvowNydN8C&q=Plato%20timaeus&pg=PA45|publisher=Forgotten Books|page=45|isbn=978-1-60620-018-6|access-date=9 November 2020|archive-date=14 April 2021|archive-url=https://web.archive.org/web/20210414140053/https://books.google.com/books?id=xSjvowNydN8C&q=Plato%20timaeus&pg=PA45|url-status=live}}</ref><ref>{{cite web|last=Hillar |first=M. |year=2004 |title=The Problem of the Soul in Aristotle's De anima |url=http://www.socinian.org/aristotles_de_anima.html |publisher=[[NASA]]/[[WMAP]] |access-date=10 August 2006 |url-status=dead |archive-url=https://web.archive.org/web/20060909004214/http://www.socinian.org/aristotles_de_anima.html |archive-date=9 September 2006 }}</ref>

[[Aristotle]], {{circa|350 BCE}}, also used the term ''stoicheia'' and added a fifth element, [[Aether (classical element)|aether]], which formed the heavens. Aristotle defined an element as:

{{blockquote|Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.<ref>{{cite book|last=Partington|first=J. R.|year=1937|title=A Short History of Chemistry|location=New York|publisher=Dover Publications|isbn=978-0-486-65977-0|url-access=registration|url=https://archive.org/details/shorthistoryofch0000part_q6h4}}</ref>}}

=== Chemical definitions ===

==== Robert Boyle ====
{{rewrite|section|date=March 2024}}
[[File:Portret van Robert Boyle, RP-P-OB-4578 (cropped).jpg|thumb|200x200px|Robert Boyle, {{circa|1740}}]]
[[File:Sceptical chymist 1661 Boyle Title page AQ18 (3).jpg|thumb|330x330px|Title page of ''The Sceptical Chymist,'' published in 1661]]

In 1661, in ''[[The Sceptical Chymist]]'', [[Robert Boyle]] proposed his theory of corpuscularism which favoured the analysis of matter as constituted of irreducible units of matter (atoms); and, choosing to side with neither Aristotle's view of the four elements nor [[Paracelsus]]' view of three fundamental elements, left open the question of the number of elements. Boyle argued against a pre-determined number of elements—directly against Paracelsus' three [[Principle (chemistry)|principles]] (sulfur, mercury, and salt), indirectly against the [[Classical element|"Aristotelian" elements]] (earth, water, air, and fire), for Boyle felt that the arguments against the former were at least as valid against the latter.
{{Blockquote|text=Much of what I am to deliver ... may be indifferently apply'd to the four Peripatetick Elements, and the three Chymical Principles ... the Chymical ''Hypothesis'' seeming to be much more countenanc'd by Experience then the other, it will be expedient to insist chiefly upon the disproving of that; especially since most of the Arguments that are imploy'd against it, may, by a little variation, be made ... at least as strongly against the less plausible, ''Aristotelian'' Doctrine.{{Sfn|Boyle|1661|p=36}}}}

Then Boyle stated his view in four propositions. In the first and second, he suggests that matter consists of particles, but that these particles may be difficult to separate. Boyle used the concept of "corpuscles"—or "atomes",{{Sfn|Boyle|1661|p=38}} as he also called them—to explain how a limited number of elements could combine into a vast number of compounds.

{{Blockquote|text='''Propos. I.''' ''... At the first Production of mixt Bodies, the Universal Matter whereof they ... consisted, was actually divided into little Particles.''{{Sfn|Boyle|1661|p=37}} ... The Generation ... and wasting of Bodies ... and ... the Chymical Resolutions of mixt Bodies, and ... Operations of ... Fires upon them ... manifest their consisting of parts very minute...  ''Epicurus'' ... as you well know, supposes ... all ... Bodies ... to be produc'd by ... Atomes, moving themselves to and fro ... in the ... Infinite ''Vacuum''.{{Sfn|Boyle|1661|pp=37–38}} ... '''Propos. II.''' ''... These minute Particles ... were ... associated into minute ... Clusters ... not easily dissipable into such Particles as compos'd them.''{{Sfn|Boyle|1661|pp=38–39}} ... If we assigne to the Corpuscles, whereof each Element consists, a peculiar size and shape ... such ... Corpuscles may be mingled in such various Proportions, and ... connected so many ... wayes, that an almost incredible number of ... Concretes may be compos'd of them.{{Sfn|Boyle|1661|p=42}}}}

Boyle explained that gold reacts with ''[[aqua regia]],'' and mercury with nitric acid, sulfuric acid, and sulfur to produce various "compounds", and that they could be recovered from those compounds, just as would be expected of elements. Yet, Boyle did not consider gold,{{Sfn|Boyle|1661|p=29}} mercury,{{Sfn|Boyle|1661|p=41}} or lead{{Sfn|Boyle|1661|p=29}} elements, but rather—together with wine{{Sfn|Boyle|1661|p=145}}—"perfectly mixt bodies". 
{{Blockquote|text=Quicksilver ... with ''Aqua fortis'' will be brought into a ... white Powder ... with Sulphur it will compose a blood-red ... Cinaber. And yet out of all these exotick Compounds, we may recover the very same running Mercury.{{Sfn|Boyle|1661|pp=40–41}} ... '''Propos. III.''' ''... From most of such mixt Bodies ... there may by the Help of the Fire, be actually obtain'd a determinate number (whether Three, Four or Five, or fewer or more) of Substances ...''
The Chymists are wont to call the Ingredients of mixt Bodies, ''Principles'', as the ''Aristotelians'' name them ''Elements''. ... ''Principles'' ... as not being compounded of any more primary Bodies: and ''Elements'', in regard that all mix'd Bodies are compounded of them.{{Sfn|Boyle|1661|p=46}}}}

Even though Boyle is primarily regarded as the first modern chemist, ''The Sceptical Chymist'' still contains old ideas about the elements, alien to a contemporary viewpoint. Sulfur, for example, is not only the familiar yellow non-metal but also an inflammable "spirit".{{Sfn|Boyle|1661|p=145}}

==== Isaac Watts ====
[[File:Portrait of Isaac Watts, D.D..jpg|thumb|216x216px|Portrait of Isaac Watts by John Shury, {{circa|1830}}]]
In 1724, in his book ''[[Logick]]'', the English minister and logician [[Isaac Watts]] enumerated the elements then recognized by chemists. Watts' list of elements included two of Paracelsus' ''principles'' (sulfur and salt) and two classical elements (earth and water) as well as "spirit". Watts did, however, note a lack of consensus among chemists.<ref>{{Cite book |last=Watts |first=Isaac |url=https://archive.org/details/logickorrightuse00wattuoft/page/13/mode/1up |title=Logick: Or, the right use of reason in the enquiry after truth, with a variety of rules to guard against error in the affairs of religion and human life, as well as in the sciences |publisher=Printed for John Clark and Richard Hett |year=1726 |pages=13–15 |orig-date=1724}}</ref>
{{Blockquote|text=Elements are such Substances as cannot be resolved, or reduced, into two or more Substances of different Kinds. ... Followers of Aristotle made Fire, Air, Earth and Water to be the four Elements, of which all earthly Things were compounded; and they suppos'd the Heavens to be a Quintessence, or fifth sort of Body, distinct from all these : But, since experimental Philosophy ... have been better understood, this Doctrine has been abundantly refuted. The Chymists make Spirit, Salt, Sulphur, Water and Earth to be their five Elements, because they can reduce all terrestrial Things to these five :.. tho' they are not all agreed.}}

==== Antoine Lavoisier, Jöns Jacob Berzelius, and Dmitri Mendeleev ====
[[File:Mendeleev's 1869 periodic table.svg|upright=1.35|thumb|[[Mendeleev]]'s 1869 periodic table: ''An experiment on a system of elements. Based on their atomic weights and chemical similarities.'']]The first modern list of elements was given in [[Antoine Lavoisier]]'s 1789 ''[[Traité Élémentaire de Chimie|Elements of Chemistry]]'', which contained 33 elements, including [[light]] and [[Caloric theory|caloric]].<ref>{{cite book|last=Lavoisier|first=A. L.|year=1790|title=Elements of chemistry translated by Robert Kerr|url=https://books.google.com/books?id=4BzAjCpEK4gC&pg=PA175|place=Edinburgh|pages=175–176|isbn=978-0-415-17914-0|access-date=24 August 2020|archive-date=14 April 2021|archive-url=https://web.archive.org/web/20210414100525/https://books.google.com/books?id=4BzAjCpEK4gC&pg=PA175|url-status=live}}</ref><ref>{{Cite book |last=Lavoisier |first=Antoine |author-link=Antoine Lavoisier |url=https://www.gutenberg.org/cache/epub/30775/pg30775-images.html#Page_175 |title=Elements of chemistry: In a new systematic order, containing all the modern discoveries: Illustrated with thirteen copperplates |publisher=William Creech |year=1790 |pages=175–176 |language=en |translator-last=Kerr |translator-first=Robert |orig-date=1789}}</ref> By 1818, [[Jöns Jacob Berzelius]] had determined atomic weights for 45 of the 49 then-accepted elements. [[Dmitri Mendeleev]] had 63 elements in his 1869 periodic table.
[[File:DIMendeleevCab.jpg|thumb|left|upright|Dmitri Mendeleev, 1897]]
From Boyle until the early 20th century, an element was defined as a pure substance that cannot be decomposed into any simpler substance and cannot be transformed into other elements by chemical processes. Elements at the time were generally distinguished by their atomic weights, a property measurable with fair accuracy by available analytical techniques.

=== Atomic definitions ===
[[File:Henry Moseley.jpg|thumb|upright=1.0|[[Henry Moseley]]]]

The 1913 discovery by English physicist [[Henry Moseley]] that the nuclear charge is the physical basis for the atomic number, further refined when the nature of protons and [[neutron]]s became appreciated, eventually led to the current definition of an element based on atomic number (number of protons). The use of atomic numbers, rather than atomic weights, to distinguish elements has greater predictive value (since these numbers are integers) and also resolves some ambiguities in the chemistry-based view due to varying properties of isotopes and [[allotrope]]s within the same element. Currently, IUPAC defines an element to exist if it has isotopes with a lifetime longer than the 10{{sup|−14}} seconds it takes the nucleus to form an electronic cloud.<ref>[http://www.kernchemie.de/Transactinides/Transactinide-2/transactinide-2.html Transactinide-2] {{Webarchive|url=https://web.archive.org/web/20160303222840/http://www.kernchemie.de/Transactinides/Transactinide-2/transactinide-2.html |date=3 March 2016 }}. www.kernchemie.de</ref>

By 1914, eighty-seven elements were known, all naturally occurring (see [[Discovery of chemical elements]]). The remaining naturally occurring elements were discovered or isolated in subsequent decades, and various additional elements have also been produced synthetically, with much of that work pioneered by [[Glenn T. Seaborg]]. In 1955, element 101 was discovered and named [[mendelevium]] in honor of D. I. Mendeleev, the first to arrange the elements periodically.

=== Discovery and recognition of various elements ===
{{For timeline|Discovery of chemical elements}}

Ten materials familiar to various prehistoric cultures are now known to be elements: Carbon, copper, [[gold]], iron, lead, mercury, silver, sulfur, [[tin]], and [[zinc]]. Three additional materials now accepted as elements, [[arsenic]], [[antimony]], and [[bismuth]], were recognized as distinct substances before 1500 AD. [[Phosphorus]], [[cobalt]], and [[platinum]] were isolated before 1750.

Most of the remaining naturally occurring elements were identified and characterized by 1900, including:
* Such now-familiar [[Industry (manufacturing)|industrial]] materials as [[aluminium]], [[silicon]], [[nickel]], [[chromium]], magnesium, and tungsten
* Reactive metals such as [[lithium]], [[sodium]], potassium, and [[calcium]]
* The [[halogen]]s [[fluorine]], [[chlorine]], [[bromine]], and [[iodine]]
* Gases such as hydrogen, oxygen, nitrogen, helium, [[argon]], and [[neon]]
* Most of the [[rare-earth elements]], including [[cerium]], [[lanthanum]], [[gadolinium]], and [[neodymium]]
* The more common [[radioactive]] elements, including uranium, thorium, and [[radium]]

Elements isolated or produced since 1900 include:
* The three remaining undiscovered stable elements: [[hafnium]], [[lutetium]], and [[rhenium]]
* [[Plutonium]], which was first produced synthetically in 1940 by [[Glenn T. Seaborg]], but is now also known from a few long-persisting natural occurrences
* The three incidentally occurring natural elements ([[neptunium]], promethium, and technetium), which were all first produced synthetically but later discovered in trace amounts in geological samples
* Four scarce decay products of uranium or thorium (astatine, francium, [[actinium]], and [[protactinium]]), and
* All synthetic [[transuranic]] elements, beginning with [[americium]] and [[curium]]

=== Recently discovered elements ===
The first [[transuranium element]] (element with an atomic number greater than 92) discovered was [[neptunium]] in 1940. Since 1999, the [[IUPAC/IUPAP Joint Working Party]] has considered claims for the discovery of new elements. As of January 2016, all 118 elements have been confirmed by IUPAC as being discovered. The discovery of element 112 was acknowledged in 2009, and the name ''copernicium'' and the chemical symbol ''Cn'' were suggested for it.<ref>{{cite web|date=20 July 2009|title=IUPAC Announces Start of the Name Approval Process for the Element of Atomic Number 112|url=http://media.iupac.org/news/112_Naming_Process_20090720.pdf|publisher=IUPAC|access-date=27 August 2009|archive-date=13 March 2012|archive-url=https://web.archive.org/web/20120313014559/http://media.iupac.org/news/112_Naming_Process_20090720.pdf|url-status=live}}</ref> The name and symbol were officially endorsed by IUPAC on 19 February 2010.<ref>{{cite web|date=20 February 2010 |title=IUPAC (International Union of Pure and Applied Chemistry): Element 112 is Named Copernicium |url=http://www.iupac.org/web/nt/2010-02-20_112_Copernicium |publisher=IUPAC |url-status=dead |archive-url=https://web.archive.org/web/20100224054826/http://www.iupac.org/web/nt/2010-02-20_112_Copernicium |archive-date=24 February 2010 }}</ref> The heaviest element that is believed to have been synthesized to date is element 118, [[oganesson]], on 9 October 2006, by the [[Joint Institute for Nuclear Research|Flerov Laboratory]] of Nuclear Reactions in [[Dubna]], Russia.{{r|Schewe}}<ref>{{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.|last9=Voinov|first9=A.|display-authors=8|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> [[Tennessine]], element 117 was the latest element claimed to be discovered, in 2009.<ref name="jinr">{{cite web|last=Greiner |first=W. |title=Recommendations |url=http://www.jinr.ru/img_sections/PAC/NP/31/PAK_NP_31_recom_eng.pdf |work=31st meeting, PAC for Nuclear Physics |publisher=[[Joint Institute for Nuclear Research]] |url-status=dead |archive-url=https://web.archive.org/web/20100414173735/http://www.jinr.ru/img_sections/PAC/NP/31/PAK_NP_31_recom_eng.pdf |archive-date=14 April 2010 }}</ref> On 28 November 2016, scientists at the IUPAC officially recognized the names for the four newest elements, with atomic numbers 113, 115, 117, and 118.<ref name="IUPAC-20161130">{{cite news |author=Staff |title=IUPAC Announces the Names of the Elements 113, 115, 117, and 118 |url=https://iupac.org/iupac-announces-the-names-of-the-elements-113-115-117-and-118/ |date=30 November 2016 |work=IUPAC |access-date=1 December 2016 |archive-date=29 July 2018 |archive-url=https://web.archive.org/web/20180729210456/https://iupac.org/iupac-announces-the-names-of-the-elements-113-115-117-and-118/ |url-status=live }}</ref><ref name="NYT-20161201">{{cite news |last=St. Fleur |first=Nicholas |title=Four New Names Officially Added to the Periodic Table of Elements |url=https://www.nytimes.com/2016/12/01/science/periodic-table-new-elements.html |archive-url=https://ghostarchive.org/archive/20220101/https://www.nytimes.com/2016/12/01/science/periodic-table-new-elements.html |archive-date=2022-01-01 |url-access=limited |date=1 December 2016 |work=[[The New York Times]] |access-date=1 December 2016 }}{{cbignore}}</ref>

== List of the 118 known chemical elements ==
{{Main|List of chemical elements}}
The following sortable table shows the 118 known elements.
* '''Atomic number''', '''Element''', and '''Symbol''' all serve independently as unique identifiers.
* '''Element''' names are those accepted by [[IUPAC]]. <!-- ; provisional names for recently produced elements not yet formally named are in parentheses. -->
* '''Block''' indicates the periodic table [[Block (periodic table)|block]] for each element: red = s-block, yellow = p-block, blue = d-block, green = f-block.
* '''Group''' and '''period''' refer to an element's position in the periodic table. Group numbers here show the currently accepted numbering; for older numberings, see [[Group (periodic table)]].
{{List of chemical elements}}

== See also ==
{{div col|colwidth=22em}}
* [[Biological roles of the elements]]
* [[Chemical database]]
* [[Discovery of chemical elements]]
* [[Element collecting]]
* [[Fictional element]]
* [[Goldschmidt classification]]
* [[Island of stability]]
* [[List of nuclides]]
* [[Densities of the elements (data page)|List of the elements' densities]]
* [[Mineral (nutrient)]]
* [[Periodic systems of small molecules]]
* [[Prices of chemical elements]]
* [[Systematic element name]]
* [[Table of nuclides]]
* [[Roles of chemical elements]]
{{div col end}}

== References ==
{{Reflist}}

== Bibliography ==
* {{Cite book |last=Boyle |first=R.|year=1661 |url=https://www.gutenberg.org/files/22914/22914-h/22914-h.htm |title=The Sceptical Chymist: or Chymico-Physical Doubts & Paradoxes, Touching the Spagyrist's Principles Commonly call'd Hypostatical; As they are wont to be Propos'd and Defended by the Generality of Alchymists. Whereunto is præmis'd Part of another Discourse relating to the same Subject |publisher=Printed by J. Cadwell for J. Crooke}}

== Further reading ==
{{Commons category|Chemical elements}}{{Wikiquote}}{{Wikijunior|The Elements}}{{Wikibooks|General Chemistry/Chemistries of Various Elements}}

* {{cite book|last=Ball|first=P.|year=2004|title=The Elements: A Very Short Introduction|publisher=Oxford University Press|isbn=978-0-19-284099-8}}

* {{cite book|last=Emsley|first=J.|year=2003|title=Nature's Building Blocks: An A–Z Guide to the Elements|publisher=Oxford University Press|isbn=978-0-19-850340-8|url-access=registration|url=https://archive.org/details/naturesbuildingb0000emsl}}
* {{cite book|last=Gray|first=T.|year=2009|title=The Elements: A Visual Exploration of Every Known Atom in the Universe|publisher=Black Dog & Leventhal Publishers Inc|isbn=978-1-57912-814-2|url-access=registration|url=https://archive.org/details/elementsvisualex0000gray}}
* {{cite book|last=Scerri|first=E.R.|year=2007|title=The Periodic Table, Its Story and Its Significance|url=https://archive.org/details/periodictableits0000scer|url-access=registration|publisher=Oxford University Press|isbn=978-0-19-530573-9}}
* {{cite book|last=Strathern|first=P.|year=2000|title=Mendeleyev's Dream: The Quest for the Elements|publisher=[[Hamish Hamilton Ltd]]|isbn=978-0-241-14065-9}}
* {{cite book|title=The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table of the Elements|last=Kean|first=Sam|year=2011|publisher=Back Bay Books}}
* {{cite book|title=Compendium of Chemical Terminology |edition=2nd |editor1=A.D. McNaught |editor2=A. Wilkinson |publisher=Blackwell Scientific Publications |location=Oxford |date=1997 |isbn=978-0-9678550-9-7 |doi=10.1351/goldbook}} XML on-line corrected version: created by M. Nic, J. Jirat, B. Kosata; updates compiled by A. Jenkins

== External links ==
* [http://periodicvideos.com/ Videos for each element] by the University of Nottingham
* [http://www.bbc.co.uk/programmes/p00546sz "Chemical Elements"], ''In Our Time'', BBC Radio 4 discussion with Paul Strathern, Mary Archer and [[John Murrell (chemist)|John Murrell]] (25 May 2000)

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