Achemical element is achemical substance whoseatoms all have the same number ofprotons. The number of protons is called theatomic 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 ofneutrons in their nuclei, known asisotopes of the element. Two or more atoms can combine to formmolecules. Some elements are formed frommolecules of identical atoms, e. g. atoms of hydrogen (H) form diatomic molecules (H2).Chemical compounds are substances made of atoms of different elements; they can have molecular or non-molecular structure.Mixtures 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 innuclear reactions, 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.[1]
The term "(chemical) element" is used in two different but closely related meanings:[2] it can mean a chemical substance consisting of a single kind of atoms (afree element), or it can mean that kind of atoms as a component of various chemical substances. For example, molecules of water (H2O) 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 (H2) and (di)oxygen (O2), as H2O molecules are different from H2 and O2 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 distinguishesélément chimique (kind of atoms) andcorps simple (chemical substance consisting of a single kind of atoms); the Russian chemical terminology distinguishesхимический элемент andпростое вещество.
Almost allbaryonic matter in the universe is composed of elements (among rare exceptions areneutron stars). When different elements undergo chemical reactions, atoms are rearranged into new compounds held together bychemical bonds. Only a few elements, such assilver andgold, are found uncombined as relatively purenative element minerals. Nearly all other naturally occurring elements occur in theEarth as compounds or mixtures.Air is mostly a mixture of molecularnitrogen andoxygen, though it does contain compounds includingcarbon dioxide andwater, as well as atomicargon, anoble gas which ischemically inert and therefore does not undergo chemical reactions.
The history of the discovery and use of elements began with earlyhuman societies that discovered native minerals likecarbon,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 ofclassical elements,alchemy, and similar theories throughout history. Much of the modern understanding of elements developed from the work ofDmitri Mendeleev, a Russian chemist who published the first recognizableperiodic table in 1869. This table organizes the elements by increasing atomic number into rows ("periods") in which the columns ("groups") share recurring ("periodic")physical andchemical 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.
Of the 94 naturally occurring elements, those with atomic numbers 1 through 82 each have at least onestable isotope (except fortechnetium, element 43 andpromethium, 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 areunstable to the point that radioactive decay of all isotopes can be detected. Some of these elements, notablybismuth (atomic number 83),thorium (atomic number 90), anduranium (atomic number 92), have one or more isotopes with half-lives long enough to survive as remnants of the explosivestellar nucleosynthesis that produced theheavy metals before the formation of ourSolar System. At over 1.9×1019 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.[7][8] The heaviest elements (those beyond plutonium, element 94) undergo radioactive decay withhalf-lives so short that they are not found in nature and must besynthesized.
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.[9][10] Most recently, the synthesis of element 118 (since namedoganesson) was reported in October 2006, and the synthesis of element 117 (tennessine) was reported in April 2010.[11][12] 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; andplutonium, 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 asprimordial nuclides present from the formation of theSolar 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).[13] This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally occurring rare elements.[14]
Theatomic number of an element is equal to the number of protons in each atom, and defines the element.[15] For example, all carbon atoms contain 6 protons in theiratomic nucleus; so the atomic number of carbon is 6.[16] Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known asisotopes of the element.[17]
The number of protons in the nucleus also determines itselectric charge, which in turn determines the number ofelectrons of the atom in itsnon-ionized state. The electrons are placed intoatomic orbitals that determine the atom'schemical properties. The number of neutrons in a nucleus usually has very little effect on an element's chemical properties; except for hydrogen (for which thekinetic 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 thanmass number oratomic weight, is considered the identifying characteristic of an element.
Isotopes are atoms of the same element (that is, with the same number ofprotons in their nucleus), but havingdifferent numbers ofneutrons. 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 ascarbon-12,carbon-13, andcarbon-14 (12C,13C, and14C). Natural carbon is amixture of12C (about 98.9%),13C (about 1.1%) and about 1 atom per trillion of14C.
Most (54 of 94) naturally occurring elements have more than one stable isotope. Except for theisotopes 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 viaalpha decay,beta decay, orinverse beta decay; some isotopes of the heaviest elements also undergospontaneous fission. Isotopes that are not radioactive, are termed "stable" isotopes. All known stable isotopes occur naturally (seeprimordial 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 composedonly 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 (fortin, element 50).
Themass number of an element,A, is the number ofnucleons (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.,238U). The mass number is always an integer and has units of "nucleons". Thus,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, theatomic mass of a particular isotope (or "nuclide") of the element is the mass of a single atom of that isotope, and is typically expressed indaltons (symbol: Da), oruniversal atomic mass units (symbol: u). Itsrelative atomic mass is a dimensionless number equal to the atomic mass divided by theatomic mass constant, which equals 1 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 Da; since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number; and because of thenuclear binding energy and electron binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969 Da and that of chlorine-37 is 36.966 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 anatural number is12C, which has a mass of 12 Da; because the dalton is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.
Thestandard atomic weight (commonly called "atomic weight") of an element is theaverage 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 isnot close to a whole number. For example, the relative atomic mass of chlorine is 35.453 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 apure element. In chemistry, a pure element means a substance whose atoms all (or in practice almost all) have the same atomic number, or number ofprotons. Nuclear scientists, however, define a pure element as one that consists of only one isotope.[18]
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%63Cu and 31%65Cu, with different numbers of neutrons. However, pure gold would be both chemically and isotopically pure, since ordinary gold consists only of one isotope,197Au.
Atoms of chemically pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiplechemical structures (spatial arrangements of atoms), known asallotropes, which differ in their properties. For example, carbon can be found asdiamond, 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;fullerenes, which have nearly spherical shapes; andcarbon nanotubes, 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 1bar 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. Inthermochemistry, an element is defined to have anenthalpy 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 betweenmetals, which readily conductelectricity,nonmetals, which do not, and a small group, (themetalloids), having intermediate properties and often behaving assemiconductors.
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:actinides,alkali metals,alkaline earth metals,halogens,lanthanides,transition metals,post-transition metals,metalloids,reactive nonmetals, andnoble gases. 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, withastatine identified as a metalloid and the others identified as nonmetals.
States of matter
Another commonly used basic distinction among the elements is theirstate of matter (phase), whethersolid,liquid, orgas, atstandard temperature and pressure (STP). Most elements are solids at STP, while several are gases. Onlybromine andmercury are liquid at 0 degrees Celsius (32 degrees Fahrenheit) and 1 atmosphere pressure;caesium andgallium 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 andboiling points, typically expressed in degreesCelsius 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 atabsolute zero at atmospheric pressure, it has only a boiling point, and not a melting point, in conventional presentations.
Thedensity at selectedstandard temperature and pressure (STP) is often used in characterizing the elements. Density is often expressed in grams per cubic centimetre (g/cm3). 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 familiarallotropes of carbon (amorphous carbon,graphite, anddiamond) have densities of 1.8–2.1, 2.267, and 3.515 g/cm3, respectively.
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 eitherstable or weakly radioactive. The longest-lived isotopes of the remaining 11 elements havehalf 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, andprotactinium) are relatively commondecay products ofthorium anduranium. The remaining six transient elements (technetium, promethium, astatine,francium,neptunium, andplutonium) 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 astungsten andlead), however, are predicted to be slightly radioactive with very long half-lives:[19] for example, the half-lives predicted for the observationally stable lead isotopes range from 1035 to 10189 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 explosivestellar nucleosynthesis that produced the heavy elements before the formation of the Solar System. For example, at over 1.9×1019 years, over a billion times longer than the estimated age of the universe,bismuth-209 has the longest knownalpha decay half-life of any isotope.[7][8] 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.
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 ("periods") in which the columns ("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 chemistDmitri 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.
The various chemical elements are formally identified by their unique atomic numbers, their accepted names, and theirchemical symbols.
Atomic numbers
The known elements have atomic numbers from 1 to 118, conventionally presented asArabic 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 increasemonotonically with their atomic numbers.
The naming of various substances now known as elements precedes theatomic 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, theofficial names of the chemical elements both ancient and more recently recognized are decided by theInternational 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 theLatin 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 aproper noun, as incalifornium andeinsteinium. Isotope names are also uncapitalized if written out,e.g.,carbon-12 oruranium-235. Chemical elementsymbols (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. (Seeelement 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 itcassiopeium. Similarly, the British discoverer ofniobium originally named itcolumbium, in reference to theNew 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, seeChemical symbol.
Specific elements
Before chemistry became ascience,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 formmolecules. 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 byJöns Jacob Berzelius in 1814. In this system, chemical symbols are not mere abbreviations—though each consists of letters of theLatin 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 theLatin names of elements (they may be Classical Latin names of elements known since antiquity orNeo-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-Latinhydrogenium, which has the same Greek roots as Englishhydrogen. However, in eleven cases Latin (as used by Berzelius) and English names of elements have different roots. Eight of them are the sevenmetals of antiquity and a metalloid also known since antiquity: "Fe" (Latinferrum) foriron, "Hg" (Latinhydrargyrum) formercury, "Sn" (Latinstannum) fortin, "Au" (Latinaurum) for gold, "Ag" (Latinargentum) forsilver, "Pb" (Latinplumbum) forlead, "Cu" (Latincuprum) forcopper, and "Sb" (Latinstibium) forantimony. The three other mismatches between Neo-Latin (as used by Berzelius) and English names are "Na" (Neo-Latinnatrium) forsodium, "K" (Neo-Latinkalium) forpotassium, and "W" (Neo-Latinwolframium) fortungsten. These mismatches came from different suggestings of naming the elements in theModern 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 theacute accent on⟨é⟩), and German (though German often allows alternate spellings with⟨k⟩ or⟨z⟩ instead of⟨c⟩: e.g., the name ofcalcium may be spelledCalcium orKalzium in German, but its symbol is always "Ca"). Other languages sometimes modify element name spellings: Spanishiterbio (ytterbium), Italianafnio (hafnium), Swedishmoskovium (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 nameJod) for iodine, but now use "I" andIod.
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 aradical, 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 ofyttrium. "Z" is also often used as a general variable group. "E" is used in organic chemistry to denote anelectron-withdrawing group or anelectrophile; similarly "Nu" denotes anucleophile. "L" is used to represent a generalligand ininorganic andorganometallic 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 anylanthanide and "An" for anyactinide. "Rg" was formerly used for anyrare gas element, but the group of rare gases has now been renamednoble gases and "Rg" now refers toroentgenium.
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 example12C and235U. 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 asH for1H (protium),D for2H (deuterium), andT for3H (tritium). This convention is easier to use in chemical equations, replacing the need to write out the mass number each time. Thus, the formula forheavy water may be written D2O instead of2H2O.
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 orions, and thus represented by elements. This fraction is about 15% of the total matter, with the remainder of the matter (85%) beingdark 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 understooddark 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 theBig Bang. ThisBig Bang nucleosynthesis happened only once; the other processes are ongoing.Nuclear fusion inside stars produces elements through stellar nucleosynthesis, including all elements from carbon toiron 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 insupernovae andneutron star mergers. The light elementslithium,beryllium andboron are produced mostly throughcosmic ray spallation (fragmentation induced bycosmic rays) of carbon, nitrogen, and oxygen.
In the early phases of the Big Bang, nucleosynthesis of hydrogen resulted in the production of hydrogen-1 (protium,1H) and helium-4 (4He), as well as a smaller amount of deuterium (2H) and tiny amounts (on the order of 10−10) 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.[22] No elements heavier than boron were produced in the Big Bang. As a result, the primordial abundance of atoms (or ions) consisted of ~75%1H, 25%4He, and 0.01% deuterium, with only tiny traces of lithium, beryllium, and perhaps boron.[23] Subsequent enrichment ofgalactic halos occurred due to stellar nucleosynthesis andsupernova nucleosynthesis.[24] However, the element abundance inintergalactic space can still closely resemble primordial conditions, unless it has been enriched by some means.
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 thealpha process. Supernovae are needed to produce "heavy" elements (those beyond iron and nickel) rapidly by neutron buildup, in ther-process. Certain large stars slowly produce other elements heavier than iron, in thes-process; these may then be blown into space in the off-gassing ofplanetary nebulae
On Earth (and elsewhere), trace amounts of various elements continue to be produced from other elements as products ofnuclear transmutation processes. These include some produced bycosmic rays or other nuclear reactions (seecosmogenic andnucleogenic nuclides), and others produced asdecay products of long-lived primordial nuclides.[25] For example, trace (but detectable) amounts ofcarbon-14 (14C) are continually produced in the air by cosmic rays impacting nitrogen atoms, and argon-40 (40Ar) is continually produced by the decay of primordially occurring but unstable potassium-40 (40K). Also, three primordially occurring but radioactive actinides, thorium, uranium, and plutonium, decay through a series of recurrently produced but unstable elements such as radium andradon, 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 bynuclear fission of the nuclei of various heavy elements or in other rare nuclear processes.
Besides the 94 naturally occurring elements, severalartificial elements have been produced bynuclear physics technology. By 2016, these experiments had produced all elements up to atomic number 118.
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 inparts per million bymass.[26] 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 thevisible 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 theBig 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 becomingsupernovas. 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.
Theabundance 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.
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 (theOddo–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.
No evidence for biological action in mammals, but essential or beneficial in some organisms
In the case of thelanthanides, the definition of an essential nutrient as being indispensable and irreplaceable is not completely applicable due to their extreme similarity. The stable early lanthanides La–Nd are known to stimulate the growth of various lanthanide-using organisms, and Sm–Gd show lesser effects for some such organisms. The later elements in the lanthanide series do not appear to have such effects.[33]
History
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.
The term 'elements' (stoicheia) was first used by Greek philosopherPlato around 360 BCE in his dialogueTimaeus, 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 byEmpedocles were composed of smallpolyhedralforms:tetrahedron (fire),octahedron (air),icosahedron (water), andcube (earth).[34][35]
Element – one of those bodies into which other bodies can decompose, and that itself is not capable of being divided into other.[36]
Chemical definitions
Robert Boyle
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Robert Boyle,c. 1740Title page ofThe Sceptical Chymist, published in 1661
In 1661, inThe 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 norParacelsus' 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' threeprinciples (sulfur, mercury, and salt), indirectly against the"Aristotelian" elements (earth, water, air, and fire), for Boyle felt that the arguments against the former were at least as valid against the latter.
Much of what I am to deliver ... may be indifferently apply'd to the four Peripatetick Elements, and the three Chymical Principles ... the ChymicalHypothesis 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.[37]
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",[38] as he also called them—to explain how a limited number of elements could combine into a vast number of compounds.
Propos. I.... At the first Production of mixt Bodies, the Universal Matter whereof they ... consisted, was actually divided into little Particles.[39] ... 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 ... InfiniteVacuum.[40] ...Propos. II.... These minute Particles ... were ... associated into minute ... Clusters ... not easily dissipable into such Particles as compos'd them.[41] ... 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.[42]
Boyle explained that gold reacts withaqua 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,[43] mercury,[44] or lead[43] elements, but rather—together with wine[45]—"perfectly mixt bodies".
Quicksilver ... withAqua 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.[46] ...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 theAristotelians name themElements. ...Principles ... as not being compounded of any more primary Bodies: andElements, in regard that all mix'd Bodies are compounded of them.[47]
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".[45]
In 1724, in his bookLogick, the English minister and logicianIsaac 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.[48]
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
Mendeleev's 1869 periodic table:An experiment on a system of elements. Based on their atomic weights and chemical similarities.
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.
The 1913 discovery by English physicistHenry Moseley that the nuclear charge is the physical basis for the atomic number, further refined when the nature of protons andneutrons 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 andallotropes within the same element. Currently, IUPAC defines an element to exist if it has isotopes with a lifetime longer than the 10−14 seconds it takes the nucleus to form an electronic cloud.[51]
By 1914, eighty-seven elements were known, all naturally occurring (seeDiscovery 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 byGlenn T. Seaborg. In 1955, element 101 was discovered and namedmendelevium in honor of D. I. Mendeleev, the first to arrange the elements periodically.
Ten materials familiar to various prehistoric cultures are now known to be elements: Carbon, copper,gold, iron, lead, mercury, silver, sulfur,tin, andzinc. Three additional materials now accepted as elements,arsenic,antimony, andbismuth, were recognized as distinct substances before 1500 AD.Phosphorus,cobalt, andplatinum were isolated before 1750.
Most of the remaining naturally occurring elements were identified and characterized by 1900, including:
Plutonium, which was first produced synthetically in 1940 byGlenn 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, andprotactinium), and
The firsttransuranium element (element with an atomic number greater than 92) discovered wasneptunium in 1940. Since 1999, theIUPAC/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 namecopernicium and the chemical symbolCn were suggested for it.[52] The name and symbol were officially endorsed by IUPAC on 19 February 2010.[53] The heaviest element that is believed to have been synthesized to date is element 118,oganesson, on 9 October 2006, by theFlerov Laboratory of Nuclear Reactions inDubna, Russia.[10][54]Tennessine, element 117 was the latest element claimed to be discovered, in 2009.[55] 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.[56][57]
Block indicates the periodic tableblock for each element: red = s-block, yellow = p-block, blue = d-block, green = f-block.
Group andperiod refer to an element's position in the periodic table. Group numbers here show the currently accepted numbering; for older numberings, seeGroup (periodic table).
^abSchewe, P.; Stein, B.; Castelvecchi, D. (16 October 2006)."Elements 116 and 118 Are Discovered".Physics News Update (797). American Institute of Physics. Archived fromthe original on 1 January 2012. Retrieved19 October 2006.
^Oganessian, Yu. Ts.; Abdullin, F. Sh.; Bailey, P. D.; Benker, D. E.; Bennett, M. E.; Dmitriev, S. N.; Ezold, J. G.; Hamilton, J. H.; Henderson, R. A.; Itkis, M. G.; Lobanov, Yu. V.; Mezentsev, A. N.; Moody, K. J.; Nelson, S. L.; Polyakov, A. N.; Porter, C. E.; Ramayya, A. V.; Riley, F. D.; Roberto, J. B.; Ryabinin, M. A.; Rykaczewski, K. P.; Sagaidak, R. N.; Shaughnessy, D. A.; Shirokovsky, I. V.; Stoyer, M. A.; Subbotin, V. G.; Sudowe, R.; Sukhov, A. M.; Tsyganov, Yu. S.; et al. (April 2010)."Synthesis of a New Element with Atomic Number Z=117".Physical Review Letters.104 (14). Physical Review Journals: 142502.Bibcode:2010PhRvL.104n2502O.doi:10.1103/PhysRevLett.104.142502.PMID20481935.
^ This article incorporates text from this source, which is in thepublic domain:"Technetium-99". United States Environmental Protection Agency Radiation Protection. Archived fromthe original on 1 September 2015. Retrieved26 February 2013.
^Chaisson, Eric J."Origins of Heavy Elements".Cosmic Evolution - From Big Bang to Humankind. Harvard–Smithsonian Center for Astrophysics.Archived from the original on 25 September 2020. Retrieved26 February 2013.
^Wright, E. L. (12 September 2004)."Big Bang Nucleosynthesis".UCLA, Division of Astronomy.Archived from the original on 13 January 2018. Retrieved22 February 2007.
^Nielsen, Forrest H. (1999). "Ultratrace minerals". In Maurice E. Shils; James A. Olsen; Moshe Shine; A. Catharine Ross (eds.).Modern nutrition in health and disease. Baltimore: Lippincott Williams & Wilkins. pp. 283–303.hdl:10113/46493.ISBN978-0683307696.
^Originally assessed as 0.7 by Pauling but never revised after other elements' electronegativities were updated for precision. Predicted to be higher than that of caesium.
Kean, Sam (2011).The Disappearing Spoon: And Other True Tales of Madness, Love, and the History of the World from the Periodic Table of the Elements. Back Bay Books.
A.D. McNaught; A. Wilkinson, eds. (1997).Compendium of Chemical Terminology (2nd ed.). Oxford: Blackwell Scientific Publications.doi:10.1351/goldbook.ISBN978-0-9678550-9-7. XML on-line corrected version: created by M. Nic, J. Jirat, B. Kosata; updates compiled by A. Jenkins
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