Which chemical element is represented by the symbol ‘Rn’?
TagMe TagMe
Chemical_element.txt Radon.txt
Chemical element Radon
A chemical element or element is a species of atoms having the same number of protons in their atomic nuclei (i.e. the same atomic number, Z).  There are 118 elements that have been identified, of which the first 94 occur naturally on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have exclusively radioactive isotopes, which decay over time into other elements. Iron is the most abundant element (by mass) making up Earth, while oxygen is the most common element in the crust of Earth. 

Chemical elements constitute all of the ordinary matter of the universe. However astronomical observations suggest that ordinary observable matter is only approximately 15% of the matter in the universe: the remainder is dark matter, the composition of which is unknown, but it is not composed of chemical elements. 
The two lightest elements, hydrogen and helium were mostly formed in the Big Bang and are the most common elements in the universe. The next three elements (lithium, beryllium and boron) were formed mostly by cosmic ray spallation, and are thus more rare than those that follow. Formation of elements with from six to twenty six protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis. The high abundance of oxygen, silicon, and iron on Earth reflects their common production in such stars. Elements with greater than twenty-six protons are formed by supernova nucleosynthesis in supernovae, which, when they explode, blast these elements far into space as supernova remnants, where they may become incorporated into planets when they are formed. 

The term "element" is used for a kind of atoms with a given number of protons (regardless of whether they are or they are not ionized or chemically bonded, e.g. hydrogen in water) as well as for a pure chemical substance consisting of a single element (e.g. hydrogen gas). For the second meaning, the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in the English-language chemical literature, whereas in some other languages their equivalent is widely used (e.g. French corps simple, Russian простое вещество). One element can form multiple substances different by their structure; they are called allotropes of the element.

When different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds.  Only a minority of elements are found uncombined as relatively pure minerals. Among the more common of such "native elements" are copper, silver, gold, carbon (as coal, graphite, or diamonds), and sulfur. All but a few of the most inert elements, such as noble gases and noble metals, are usually found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined forms, most of these occur as mixtures. For example, atmospheric air is primarily a mixture of nitrogen, oxygen, and argon, and native solid elements occur in alloys, such as that of iron and nickel.

The history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur, copper and gold. Later civilizations extracted elemental copper, tin, lead and iron from their ores by smelting, using charcoal. Alchemists and chemists subsequently identified many more, with almost all of the naturally-occurring elements becoming known by 1900.

The properties of the chemical elements are summarized on the periodic table, which organizes the elements by increasing atomic number into rows ("periods") in which the columns ("groups") share recurring ("periodic") physical and chemical properties. Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, most of them in high degrees of purity.

Description

The lightest chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe  in a ratio of around 3:1 by mass (or 12:1 by number of atoms),   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.  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 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 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), uranium (atomic number 92) and plutonium (atomic number 94), 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 our solar system. For example, at over 1.9 years, over a billion times longer than the current estimated age of the universe, bismuth-209 (atomic number 83) has the longest known alpha decay half-life of any naturally occurring element. The very heaviest elements (those beyond plutonium, element 94) undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized.

As of 2010, there are 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).  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 nuclides 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: these are all radioactive, with very short half-lives; if any atoms of these elements were present at the formation of Earth, they are extremely likely, to the point of certainty, to have already decayed, and if present in novae, have been in quantities too small to have been noted. Technetium was the first purportedly non-naturally occurring element synthesized, in 1937, although trace amounts of technetium have since been found in nature (and also the element may have been discovered naturally in 1925).  This pattern of artificial production and later natural discovery has been repeated with several other radioactive naturally-occurring rare elements. 

Lists of the elements are available by name, by symbol, by atomic number, by density, by melting point, and by boiling point as well as ionization energies of the elements. 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

The atomic number of an element is equal to the number of protons in each atom, and defines the element.  For example, all carbon atoms contain 6 protons in their atomic nucleus; so the atomic number of carbon is 6.  Carbon atoms may have different numbers of neutrons; atoms of the same element having different numbers of neutrons are known as isotopes of the element. 

The number of protons in the atomic nucleus also determines its electric charge, which in turn determines the number of electrons of the atom in its non-ionized state. The electrons are placed into atomic orbitals that determine the atom's various chemical properties. The number of neutrons in a nucleus usually has very little effect on an element's chemical properties (except in the case of hydrogen and deuterium). Thus, all carbon isotopes have nearly identical chemical properties because they all have six protons and six electrons, even though carbon atoms may, for example, have 6 or 8 neutrons. That is why the atomic number, rather than mass number or atomic weight, is considered the identifying characteristic of a chemical element.

The symbol for atomic number is Z.

Isotopes

Isotopes are atoms of the same element (that is, with the same number of protons in their atomic nucleus), but having different numbers of neutrons. Most (66 of 94) naturally occurring elements have more than one stable isotope. Thus, for example, there are three main isotopes of carbon. All carbon atoms have 6 protons in the nucleus, but they can have either 6, 7, or 8 neutrons. Since the mass numbers of these are 12, 13 and 14 respectively, the three isotopes of carbon are known as carbon-12, carbon-13, and carbon-14, often abbreviated to 12C, 13C, and 14C. Carbon in everyday life and in chemistry is a mixture of 12C (about 98.9%), 13C (about 1.1%) and about 1 atom per trillion of 14C.

Except in the case of 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 of the elements have some isotopes that are radioactive (radioisotopes), although not all of these radioisotopes occur naturally. The radioisotopes typically decay into other elements upon radiating an alpha or beta particle. If an element has isotopes that are not radioactive, these are termed "stable" isotopes. All of the known stable isotopes occur naturally (see primordial isotope). The many radioisotopes that are not found in nature have been characterized after being artificially made. Certain elements have no stable isotopes and are composed only of radioactive isotopes: specifically the elements without any stable isotopes are technetium (atomic number 43), promethium (atomic number 61), and all observed elements with atomic numbers greater than 82.

Of the 80 elements with at least one stable isotope, 26 have only one single 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 that occur for a single element is 10 (for tin, element 50).

Isotopic mass and atomic mass

The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a superscript on the left hand side of the atomic symbol (e.g., 238U). The mass number is always a simple whole number and has units of "nucleons." An example of a referral to a mass number is "magnesium-24," which 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 a natural (or whole) number, the atomic mass of a single atom is a real number for the mass of a particular isotope of the element, the unit being u. In general, when expressed in u it differs in value slightly from the mass number for a given nuclide (or isotope) since the mass of the protons and neutrons is not exactly 1 u, since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number, and (finally) because of the nuclear binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969 u and that of chlorine-37 is 36.966 u. However, the atomic mass in u of each isotope is quite close to its simple mass number (always within 1%). The only isotope whose atomic mass is exactly a natural number is 12C, which by definition has a mass of exactly 12, because u is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.

The relative atomic mass (historically and commonly also 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 (u). This number may be a fraction that is not close to a whole number, due to the averaging process. For example, the relative atomic mass of chlorine is 35.453 u, which differs greatly from a whole number due to being made of an average of 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 resulting from significant amounts of more than one isotope being naturally present in the sample of the element in question.

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 protons. Nuclear scientists, however, define a pure element as one that consists of only one stable isotope. 

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, a pure gold ingot would be both chemically and isotopically pure, since ordinary gold consists only of one isotope, 197Au.

Allotropes

Atoms of chemically pure elements may bond to each other chemically in more than one way, allowing the pure element to exist in multiple structures (spatial arrangements of atoms), known as allotropes, 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; fullerenes, which have nearly spherical shapes; and carbon 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 standard state, also known as reference state, of an element is defined as its thermodynamically most stable state at 1 bar at a given temperature (typically at 298.15 K). In thermochemistry, an element is defined to have an enthalpy of formation of zero in its standard state. For example, the reference state for carbon is graphite, because the structure of graphite is more stable than that of the other allotropes.

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 metals, which readily conduct electricity, nonmetals, which do not, and a small group, (the metalloids), having intermediate properties and often behaving as semiconductors.

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, polyatomic nonmetals, diatomic nonmetals, and noble 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 polyatomic nonmetals, diatomic 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 a selected standard temperature and pressure (STP). Most of the elements are solids at conventional temperatures and atmospheric pressure, while several are gases. Only bromine and mercury are liquids at 0 degrees Celsius (32 degrees Fahrenheit) and normal atmospheric pressure; caesium and gallium are solids 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 and boiling points, 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

The density at a selected standard temperature and pressure (STP) is frequently used in characterizing the elements. Density is often expressed in grams per cubic centimeter (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 familiar allotropes of carbon (amorphous carbon, graphite, and diamond) have densities of 1.8–2.1, 2.267, and 3.515 g/cm3, respectively.

Crystal structures

The elements studied to date as solid samples have eight kinds of crystal structures: cubic, body-centered cubic, face-centered cubic, hexagonal, monoclinic, orthorhombic, rhombohedral, and tetragonal. For some of the synthetically produced transuranic elements, available samples have been too small to determine crystal structures.

Occurrence and origin on Earth

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 as the synthetic products of man-made nuclear reactions.

Of the 94 naturally occurring elements, 84 are considered primordial and either stable or weakly radioactive. The remaining 10 naturally occurring elements possess half lives too short for them to have been present at the beginning of the Solar System, and are therefore considered transient elements. (Plutonium is sometimes also considered a transient element because primordial plutonium has by now decayed to almost undetectable traces.) Of these 10 transient elements, 5 (polonium, radon, radium, actinium, and protactinium) are relatively common decay products of thorium, uranium, and plutonium. The remaining 5 transient elements (technetium, promethium, astatine, francium, and neptunium) 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 40 are all stable, while those with atomic numbers 41 through 82 (except technetium and promethium) are metastable. The half-lives of these metastable "theoretical radionuclides" are so long (at least 100 million times longer than the estimated age of the universe) that their radioactive decay has yet to be detected by experiment. Elements with atomic numbers 83 through 94 are unstable to the point that their radioactive decay can be detected. Four of these elements, bismuth (element 83), thorium (element 90), uranium (element 92), and plutonium (element 94), 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 our solar system. For example, at over 1.9 years, over a billion times longer than the current estimated age of the universe, bismuth-209 has the longest known alpha decay half-life of any naturally occurring element. The very heaviest elements (those beyond plutonium, element 94) undergo radioactive decay with short half-lives and do not occur in nature.

The periodic table

The properties of the chemical 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 current standard table contains 118 confirmed elements as of 10 April 2010.

Although earlier precursors to this presentation exist, its invention is generally credited to the 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 within the academic discipline of 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

The various chemical elements are formally identified by their unique atomic numbers, by their accepted names, and by their symbols.

Atomic numbers

The known elements have atomic numbers from 1 through 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", although technically the weight or mass of atoms of an element (their atomic weights or atomic masses) do not always increase monotonically with their atomic numbers.

Element names

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, although 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 names of elements either for convenience, linguistic niceties, or nationalism. For a few illustrative examples: German speakers use "Wasserstoff" (water substance) for "hydrogen", "Sauerstoff" (acid substance) for "oxygen" and "Stickstoff" (smothering substance) for "nitrogen", while English and some romance 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 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 the 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, chemical elements are not proper nouns in English; consequently, the full name of an element is not routinely capitalized in English, even if derived from a proper noun, as in californium and einsteinium. Isotope names of chemical elements 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 twentieth century, physics laboratories became able to produce nuclei of chemical 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 prior to the international standardization.

Chemical symbols

Specific chemical elements

Before chemistry became a science, alchemists had 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 molecules. 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 Berzelius. In this typographical 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.

The first of these symbols were intended to be fully universal. Since Latin was the common language of science at that time, they were abbreviations based on the Latin names of metals. Cu comes from Cuprum, Fe comes from Ferrum, Ag from Argentum. The symbols were not followed by a period (full stop) as with abbreviations. Later chemical elements were also assigned unique chemical symbols, based on the name of the element, but not necessarily in English. For example, sodium has the chemical symbol 'Na' after the Latin natrium. The same applies to "W" (wolfram) for tungsten, "Fe" (ferrum) for iron, "Hg" (hydrargyrum) for mercury, "Sn" (stannum) for tin, "K" (kalium) for potassium, "Au" (aurum) for gold, "Ag" (argentum) for silver, "Pb" (plumbum) for lead, "Cu" (cuprum) for copper, and "Sb" (stibium) for antimony.

Chemical symbols are understood internationally when element names might require translation. There have sometimes been differences in the past. For example, Germans in the past have used "J" (for the alternate name Jod) for iodine, but now use "I" and "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 (small letters). Thus, the symbols for californium or einsteinium are Cf and Es.

General chemical symbols

There are also symbols in chemical equations for groups of chemical 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, an "X" indicates a variable group (usually a halogen) in a class of compounds, while "R" is a 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, although it is also the symbol of yttrium. "Z" is also frequently used as a general variable group. "E" is used in organic chemistry to denote an electron-withdrawing group.  "L" is used to represent a general ligand in inorganic and organometallic chemistry. "M" is also often used in place of a general metal.

At least two additional, two-letter generic chemical symbols are also in informal usage, "Ln" for any lanthanide element and "An" for any actinide element. "Rg" was formerly used for any rare gas element, but the group of rare gases has now been renamed noble gases and the symbol "Rg" has now been assigned to the element roentgenium.

Isotope symbols

Isotopes are distinguished by the atomic mass number (total protons and neutrons) for a particular isotope of an element, with this number combined with the pertinent element's symbol. IUPAC prefers that isotope symbols be written in superscript notation when practical, for example 12C and 235U. 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 the element hydrogen are often specified as H for 1H (protium), D for 2H (deuterium), and T for 3H (tritium). This convention is easier to use in chemical equations, replacing the need to write out the mass number for each atom. For example, the formula for heavy water may be written D2O instead of 2H2O.

Origin of the elements

Only about 4% of the total mass of the universe is made of atoms or ions, and thus represented by chemical 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 chemical 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 more mysterious dark energy).

The universe's 94 naturally occurring chemical elements are thought to have been produced by at least four cosmic processes. Most of the hydrogen and helium in the universe was produced primordially in the first few minutes of the Big Bang. Three recurrently occurring later processes are thought to have produced the remaining elements. Stellar nucleosynthesis, an ongoing process, produces all elements from carbon through iron in atomic number, but little lithium, beryllium, or boron. Elements heavier in atomic number than iron, as heavy as uranium and plutonium, are produced by explosive nucleosynthesis in supernovas and other cataclysmic cosmic events. Cosmic ray spallation (fragmentation) of carbon, nitrogen, and oxygen is important to the production of lithium, beryllium and boron.

During the early phases of the Big Bang, nucleosynthesis of hydrogen nuclei resulted in the production of hydrogen-1 (protium, 1H) and helium-4 (4He), as well as a smaller amount of deuterium (2H) and very minuscule 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.  It is generally agreed that no heavier elements than boron were produced in the Big Bang. As a result, the primordial abundance of atoms (or ions) consisted of roughly 75% 1H, 25% 4He, and 0.01% deuterium, with only tiny traces of lithium, beryllium, and perhaps boron.  Subsequent enrichment of galactic halos occurred due to stellar nucleosynthesis and supernova nucleosynthesis. However, the element abundance in intergalactic space can still closely resemble primordial conditions, unless it has been enriched by some means.

On Earth (and elsewhere), trace amounts of various elements continue to be produced from other elements as products of natural transmutation processes. These include some produced by cosmic rays or other nuclear reactions (see cosmogenic and nucleogenic nuclides), and others produced as decay products of long-lived primordial nuclides. For example, trace (but detectable) amounts of carbon-14 (14C) are continually produced in the atmosphere 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 radioactive elements such as radium and radon, which are transiently present in any sample of these metals or their ores or compounds. Three other radioactive elements, technetium, promethium, and neptunium, occur only incidentally in natural materials, produced as individual atoms by natural fission of the nuclei of various heavy elements or in other rare nuclear processes.

Human technology has produced various additional elements beyond these first 94, with those through atomic number 118 now known.

Abundance

The following graph (note log scale) shows the abundance of elements in our solar system. The table shows the twelve most common elements in our galaxy (estimated spectroscopically), as measured in parts per million, by mass. 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, scientist 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 supernovas. Iron-56 is particularly common, since it is the most stable element 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 heavy 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. Aluminum 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 aluminum (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 proteins and nucleic acids, 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 organisms require particular additional elements, for example the magnesium in chlorophyll in green plants, the calcium in mollusc shells, or the iron in the hemoglobin in vertebrate animals' red blood cells.

History

Evolving definitions

The concept of an "element" as an undivisible substance has developed through three major historical phases: Classical definitions (such as those of the ancient Greeks), chemical definitions, and atomic definitions.

Classical definitions

Ancient philosophy posited a set of classical elements to explain observed patterns in nature. These elements originally referred to earth, water, air and fire rather than the chemical elements of modern science.

The term 'elements' (stoicheia) was first used by the Greek philosopher Plato in about 360 BCE in his 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 polyhedral forms: tetrahedron (fire), octahedron (air), icosahedron (water), and cube (earth).  

Aristotle, c. 350 BCE, also used the term stoicheia and added a fifth element called aether, which formed the heavens. Aristotle defined an element as:

Chemical definitions

In 1661, Robert Boyle proposed his theory of corpuscularism which favoured the analysis of matter as constituted by 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. The first modern list of chemical elements was given in Antoine Lavoisier's 1789 Elements of Chemistry, which contained thirty-three elements, including light and caloric.  By 1818, Jöns Jakob Berzelius had determined atomic weights for forty-five of the forty-nine then-accepted elements. Dmitri Mendeleev had sixty-six elements in his periodic table of 1869.

From Boyle until the early 20th century, an element was defined as a pure substance that could not be decomposed into any simpler substance. Put another way, a chemical element cannot be transformed into other chemical elements by chemical processes. Elements during this time were generally distinguished by their atomic weights, a property measurable with fair accuracy by available analytical techniques.

Atomic definitions

The 1913 discovery by English physicist Henry Moseley that the nuclear charge is the physical basis for an atom's atomic number, further refined when the nature of protons and neutrons became appreciated, eventually led to the current definition of an element based on atomic number (number of protons per atomic nucleus). 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 allotropes 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. 

By 1914, seventy-two elements were known, all naturally occurring.  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 in a periodic manner. Most recently, the synthesis of element 118 was reported in October 2006, and the synthesis of element 117 was reported in April 2010. 

Discovery and recognition of various elements

Ten materials familiar to various prehistoric cultures are now known to be chemical 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 prior to 1500 AD. Phosphorus, cobalt, and platinum were isolated before 1750.

Most of the remaining naturally occurring chemical elements were identified and characterized by 1900, including:
* Such now-familiar industrial materials as aluminium, silicon, nickel, chromium, magnesium, and tungsten
* Reactive metals such as lithium, sodium, potassium, and calcium
* The halogens 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, radium, and radon

Elements isolated or produced since 1900 include:
* The three remaining undiscovered regularly occurring stable natural 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 certain geological samples
* Three scarce decay products of uranium or thorium, (astatine, francium, and protactinium), and
* Various synthetic transuranic elements, beginning with americium and curium

Recently discovered elements

The first transuranium element (element with atomic number greater than 92) discovered was neptunium in 1940. Since 1999 claims for the discovery of new elements have been considered by the IUPAC/IUPAP Joint Working Party. As of January 2016, all 118 elements have been confirmed as discovered by IUPAC. The discovery of element 112 was acknowledged in 2009, and the name copernicium and the atomic symbol Cn were suggested for it.  The name and symbol were officially endorsed by IUPAC on 19 February 2010.  The heaviest element that is believed to have been synthesized to date is element 118, ununoctium, on 9 October 2006, by the Flerov Laboratory of Nuclear Reactions in Dubna, Russia.  Element 117 was the latest element claimed to be discovered, in 2009. IUPAC officially recognized flerovium and livermorium, elements 114 and 116, in June 2011 and approved their names in May 2012.  In December 2015, IUPAC recognized elements 113, 115, 117 and 118,  and announced the elements' proposed final names on 8 June 2016. The names, nihonium (113, Nh), moscovium (115, Mc), tennessine (117, Ts), and oganesson (118, Og), are expected to be approved by the end of 2016. 

List of the 118 known chemical elements

The following sortable table includes the 118 known chemical elements, with the names linking to the Wikipedia articles on each.
* Atomic number, name, and symbol all serve independently as unique identifiers.
* Names are those accepted by IUPAC; provisional names for recently produced elements not yet formally named are in parentheses.
* Group, period, and block refer to an element's position in the periodic table. Group numbers here show the currently accepted numbering; for older alternate numberings, see Group (periodic table).
* State of matter (solid, liquid, or gas) applies at standard temperature and pressure conditions (STP).
* Occurrence distinguishes naturally occurring elements, categorized as either primordial or transient (from decay), and additional synthetic elements that have been produced technologically, but are not known to occur naturally.
* Description summarizes an element's properties using the broad categories commonly presented in periodic tables: Actinide, alkali metal, alkaline earth metal, lanthanide, post-transition metal, metalloid, noble gas, polyatomic or diatomic nonmetal, and transition metal. Radon is a chemical element with symbol Rn and atomic number 86. It is a radioactive, colorless, odorless, tasteless noble gas, occurring naturally as a decay product of radium. Its most stable isotope, 222Rn, has a half-life of 3.8 days. Radon is one of the densest substances that remains a gas under normal conditions. It is also the only gas under normal conditions that has no stable isotopes, and is considered a health hazard due to its radioactivity. Intense radioactivity has also hindered chemical studies of radon and only a few compounds are known.

Radon is formed as one intermediate step in the normal radioactive decay chains through which thorium and uranium slowly decay into lead. Thorium and uranium are the two most common radioactive elements on earth; they have been around since the earth was formed. Their naturally occurring isotopes have very long half-lives, on the order of billions of years. Thorium and uranium, their decay product radium, and its decay product radon, will therefore continue to occur for tens of millions of years at almost the same concentrations as they do now.[http://www.bvsde.paho.org/bvstox/i/fulltext/toxprofiles/radon.pdf Toxicological profile for radon], Agency for Toxic Substances and Disease Registry, U.S. Public Health Service, In collaboration with U.S. Environmental Protection Agency, December 1990. As radon itself decays, it produces other radioactive elements called radon progeny (also known as radon daughters) or decay products. Unlike the gaseous radon itself, radon daughters are solids and stick to surfaces, such as dust particles in the air. If such contaminated dust is inhaled, these particles can stick to the airways of the lung and increase the risk of developing lung cancer. 

Unlike all the other intermediate elements in the aforementioned decay chains, radon is gaseous and easily inhaled. Thus, naturally-occurring radon is responsible for the majority of the public exposure to ionizing radiation. It is often the single largest contributor to an individual's background radiation dose, and is the most variable from location to location. Despite its short lifetime, some radon gas from natural sources can accumulate to far higher than normal concentrations in buildings, especially in low areas such as basements and crawl spaces due to its density. It can also occur in water where the water comes from a ground source -e.g. in some spring waters and hot springs. 

Epidemiological studies have shown a clear link between breathing high concentrations of radon and incidence of lung cancer. Thus, radon is considered a significant contaminant that affects indoor air quality worldwide. According to the United States Environmental Protection Agency, radon is the second most frequent cause of lung cancer, after cigarette smoking, causing 21,000 lung cancer deaths per year in the United States. About 2,900 of these deaths occur among people who have never smoked. While radon is the second most frequent cause of lung cancer, it is the number one cause among non-smokers, according to EPA estimates.

Characteristics

Physical properties

Radon is a colorless, odorless, and tasteless gas and therefore not detectable by human senses alone. At standard temperature and pressure, radon forms a monatomic gas with a density of 9.73 kg/m3, about 8 times the density of the Earth's atmosphere at sea level, 1.217 kg/m3.  Radon is one of the densest gases at room temperature and is the densest of the noble gases. Although colorless at standard temperature and pressure, when cooled below its freezing point of 202 K, radon emits a brilliant radioluminescence that turns from yellow to orange-red as the temperature lowers.  Upon condensation, radon glows because of the intense radiation it produces.  Radon is sparingly soluble in water, but more soluble than lighter noble gases. Radon is appreciably more soluble in organic liquids than in water.

Chemical properties

Being a noble gas, radon is chemically not very reactive. However, the 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer.

Radon is a member of the zero-valence elements that are called noble gases. It is inert to most common chemical reactions, such as combustion, because the outer valence shell contains eight electrons. This produces a stable, minimum energy configuration in which the outer electrons are tightly bound.  1037 kJ/mol is required to extract one electron from its shells (also known as the first ionization energy).  In accordance with periodic trends, radon has a lower electronegativity than the element one period before it, xenon, and is therefore more reactive.  Early studies concluded that the stability of radon hydrate should be of the same order as that of the hydrates of chlorine () or sulfur dioxide (), and significantly higher than the stability of the hydrate of hydrogen sulfide (). 

Because of its cost and radioactivity, experimental chemical research is seldom performed with radon, and as a result there are very few reported compounds of radon, all either fluorides or oxides. Radon can be oxidized by powerful oxidizing agents such as fluorine, thus forming radon difluoride.   It decomposes back to elements at a temperature of above 250 °C. It has a low volatility and was thought to be . Because of the short half-life of radon and the radioactivity of its compounds, it has not been possible to study the compound in any detail. Theoretical studies on this molecule predict that it should have a Rn–F bond distance of 2.08 Å, and that the compound is thermodynamically more stable and less volatile than its lighter counterpart . The octahedral molecule  was predicted to have an even lower enthalpy of formation than the difluoride.  The higher fluorides RnF4 and RnF6 have been claimed, and are calculated to be stable,  but it is doubtful whether they have yet been synthesized.  The [RnF]+ ion is believed to form by the following reaction: 

Rn (g) + 2  (s) →  (s) + 2  (g)

Radon oxides are among the few other reported compounds of radon;  only the trioxide has been confirmed. Radon carbonyl RnCO has been predicted to be stable and to have a linear molecular geometry.  The molecules  and RnXe were found to be significantly stabilized by spin-orbit coupling. Radon caged inside a fullerene has been proposed as a drug for tumors.  Despite the existence of Xe(VIII), no Rn(VIII) compounds have been claimed to exist; RnF8 should be highly unstable chemically (XeF8 is thermodynamically unstable). It is predicted that the most stable Rn(VIII) compound would be barium perradate (Ba2RnO6), analogous to barium perxenate. The instability of Rn(VIII) is due to the relativistic stabilization of the 6s shell, also known as the inert pair effect.

Isotopes

Radon has no stable isotopes. Thirty-six radioactive isotopes have been characterized, with atomic masses ranging from 193 to 228. The most stable isotope is 222Rn, which is a decay product of 226Ra, a decay product of 238U.  A trace amount of the (highly unstable) isotope 218Rn is also among the daughters of 222Rn.

Three other radon isotopes have a half-life of over an hour: 211Rn, 210Rn and 224Rn. The 220Rn isotope is a natural decay product of the most stable thorium isotope (232Th), and is commonly referred to as thoron. It has a half-life of 55.6 seconds and also emits alpha radiation. Similarly, 219Rn is derived from the most stable isotope of actinium (227Ac)—named "actinon"—and is an alpha emitter with a half-life of 3.96 seconds. No radon isotopes occur significantly in the neptunium (237Np) decay series, though a trace amount of the (extremely unstable) isotope 217Rn is produced.

Daughters

222Rn belongs to the radium and uranium-238 decay chain, and has a half-life of 3.8235 days. Its four first products (excluding marginal decay schemes) are very short-lived, meaning that the corresponding disintegrations are indicative of the initial radon distribution. Its decay goes through the following sequence:
* 222Rn, 3.8 days, alpha decaying to...
* 218Po, 3.10 minutes, alpha decaying to...
* 214Pb, 26.8 minutes, beta decaying to...
* 214Bi, 19.9 minutes, beta decaying to...
* 214Po, 0.1643 ms, alpha decaying to...
* 210Pb, which has a much longer half-life of 22.3 years, beta decaying to...
* 210Bi, 5.013 days, beta decaying to...
* 210Po, 138.376 days, alpha decaying to...
* 206Pb, stable.

The radon equilibrium factor  is the ratio between the activity of all short-period radon progenies (which are responsible for most of radon's biological effects), and the activity that would be at equilibrium with the radon parent.

If a closed volume is constantly supplied with radon, the concentration of short-lived isotopes will increase until an equilibrium is reached where the rate of decay of each decay product will equal that of the radon itself. The equilibrium factor is 1 when both activities are equal, meaning that the decay products have stayed close to the radon parent long enough for the equilibrium to be reached, within a couple of hours. Under these conditions each additional pCi/L of radon will increase exposure, by 0.01 WL (Working Level -a measure of radioactivity commonly used in mining. A detailed explanation of WL is given in Concentration Units). These conditions are not always met; in many homes, the equilibrium fraction is typically 40%; that is, there will be 0.004 WL of daughters for each pCi/L of radon in air. 210Pb takes much longer (decades) to come in equilibrium with radon, but, if the environment permits accumulation of dust over extended periods of time, 210Pb and its decay products may contribute to overall radiation levels as well.

Because of their electrostatic charge, radon progenies adhere to surfaces or dust particles, whereas gaseous radon does not. Attachment removes them from the air, usually causing the equilibrium factor in the atmosphere to be less than one. The equilibrium factor is also lowered by air circulation or air filtration devices, and is increased by airborne dust particles, including cigarette smoke. In high concentrations, airborne radon isotopes contribute significantly to human health risk. The equilibrium factor found in epidemiological studies is 0.4. 

History and etymology

Radon was the fifth radioactive element to be discovered, in 1900 by Friedrich Ernst Dorn,  after uranium, thorium, radium and polonium.     In 1900 Dorn reported some experiments in which he noticed that radium compounds emanate a radioactive gas he named Radium Emanation (Ra Em).  Before that, in 1899, Pierre and Marie Curie observed that the gas emitted by radium remained radioactive for a month.  Later that year, Robert B. Owens and Ernest Rutherford, at McGill University in Montreal, noticed variations when trying to measure radiation from thorium oxide.  Rutherford noticed that the compounds of thorium continuously emit a radioactive gas that retains the radioactive powers for several minutes, and called this gas emanation (from Latin "emanare"—to elapse and "emanatio"—expiration),  and later Thorium Emanation (Th Em). In 1901, he demonstrated that the emanations are radioactive, but credited the Curies for the discovery of the element.  In 1903, similar emanations were observed from actinium by André-Louis Debierne   and were called Actinium Emanation (Ac Em).

Several names were suggested for these three gases: exradio, exthorio, and exactinio in 1904; radon, thoron, and akton in 1918;  radeon, thoreon, and actineon in 1919,  and eventually radon, thoron, and actinon in 1920.  The likeness of the spectra of these three gases with those of argon, krypton, and xenon, and their observed chemical inertia led Sir William Ramsay to suggest in 1904 that the "emanations" might contain a new element of the noble gas family.

In 1910, Ramsay and Robert Whytlaw-Gray isolated radon, determined its density, and determined that it was the heaviest known gas. They wrote that "L'expression de l'émanation du radium est fort incommode", (the expression 'radium emanation' is very awkward) and suggested the new name niton (Nt) (from the Latin "nitens" meaning "shining") to emphasize the radioluminescence property, and in 1912 it was accepted by the International Commission for Atomic Weights. In 1923, the International Committee for Chemical Elements and International Union of Pure and Applied Chemistry (IUPAC) chose among the names radon (Rn), thoron (Tn), and actinon (An). Later, when isotopes were numbered instead of named, the element took the name of the most stable isotope, radon, while Tn was renamed 220Rn and An was renamed 219Rn. As late as the 1960s, the element was also referred to simply as emanation.  The first synthesized compound of radon, radon fluoride, was obtained in 1962. 

The danger of high exposure to radon in mines, where exposures reaching 1,000,000 Bq/m3 can be found, has long been known. In 1530, Paracelsus described a wasting disease of miners, the mala metallorum, and Georg Agricola recommended ventilation in mines to avoid this mountain sickness (Bergsucht).   In 1879, this condition was identified as lung cancer by Herting and Hesse in their investigation of miners from Schneeberg, Germany. The first major studies with radon and health occurred in the context of uranium mining in the Joachimsthal region of Bohemia.  In the US, studies and mitigation only followed decades of health effects on uranium miners of the Southwestern United States employed during the early Cold War; standards were not implemented until 1971. 

The presence of radon in indoor air was documented as early as 1950. Beginning in the 1970s research was initiated to address sources of indoor radon, determinants of concentration, health effects, and approaches to mitigation. In the United States, the problem of indoor radon received widespread publicity and intensified investigation after a widely publicized incident in 1984. During routine monitoring at a Pennsylvania nuclear power plant, a worker was found to be contaminated with radioactivity. A high contamination of radon in his home was subsequently identified as responsible for the contamination. 

Occurrence

Concentration units

All discussions of radon concentrations in the environment refer to 222Rn. While the average rate of production of 220Rn (from the thorium decay series) is about the same as 222Rn, the amount of 220Rn in the environment is much less than that of 222Rn because of the short half-life of 220Rn (55 seconds, versus 3.8 days respectively).

Radon concentration in the atmosphere is usually measured in becquerel per cubic meter (Bq/m3), the SI derived unit. Another unit of measurement common in the US is picocuries per liter (pCi/L); 1 pCi/L37 Bq/m3.  Typical domestic exposures average about 48 Bq/m3 indoors, though this varies widely, and 15 Bq/m3 outdoors.

In the mining industry, the exposure is traditionally measured in working level (WL), and the cumulative exposure in working level month (WLM); 1 WL equals any combination of short-lived 222Rn daughters (218Po, 214Pb, 214Bi, and 214Po) in 1 liter of air that releases 1.3 × 105 MeV of potential alpha energy; one WL is equivalent to 2.08 × 10−5 joules per cubic meter of air (J/m3). The SI unit of cumulative exposure is expressed in joule-hours per cubic meter (J·h/m3). One WLM is equivalent to 3.6 × 10−3 J·h/m3. An exposure to 1 WL for 1 working month (170 hours) equals 1 WLM cumulative exposure. A cumulative exposure of 1 WLM is roughly equivalent to living one year in an atmosphere with a radon concentration of 230 Bq/m3. 

Radon (222Rn), decays to 210Pb and other radioisotopes. The levels of 210Pb can be measured. The rate of deposition of this radioisotope is weather-dependent.

Radon concentrations found in natural environments are much too low to be detected by chemical means. A 1000 Bq/m3 (relatively high) concentration corresponds to 0.17 picogram per cubic meter. The average concentration of radon in the atmosphere is about 6 atoms of radon for each molecule in the air, or about 150 atoms in each ml of air.  The radon activity of the entire Earth's atmosphere originates from only a few tens of grams of radon, consistently replaced by decay of larger amounts of radium and uranium. 

Natural

Radon is produced by the radioactive decay of radium-226, which is found in uranium ores, phosphate rock, shales, igneous and metamorphic rocks such as granite, gneiss, and schist, and to a lesser degree, in common rocks such as limestone.  Every square mile of surface soil, to a depth of 6 inches (2.6 km2 to a depth of 15 cm), contains approximately 1 gram of radium, which releases radon in small amounts to the atmosphere. On a global scale, it is estimated that 2,400 million curies (90 EBq) of radon are released from soil annually. 

Radon concentration varies widely from place to place. In the open air, it ranges from 1 to 100 Bq/m3, even less (0.1 Bq/m3) above the ocean. In caves or aerated mines, or ill-aerated houses, its concentration climbs to 20–2,000 Bq/m3. 

Radon concentration can be much higher in mining contexts. Ventilation regulations instruct to maintain radon concentration in uranium mines under the "working level", with 95th percentile levels ranging up to nearly 3 WL (546 pCi 222Rn per liter of air; 20.2 kBq/m3, measured from 1976 to 1985).
The concentration in the air at the (unventilated) Gastein Healing Gallery averages 43 kBq/m3 (1.2 nCi/L) with maximal value of 160 kBq/m3 (4.3 nCi/L).

Radon mostly appears with the decay chain of the radium and uranium series (222Rn), and marginally with the thorium series (220Rn). The element emanates naturally from the ground, and some building materials, all over the world, wherever traces of uranium or thorium can be found, and particularly in regions with soils containing granite or shale, which have a higher concentration of uranium. Not all granitic regions are prone to high emissions of radon. Being a rare gas, it usually migrates freely through faults and fragmented soils, and may accumulate in caves or water. Owing to its very short half-life (four days for 222Rn), radon concentration decreases very quickly when the distance from the production area increases. Radon concentration varies greatly with season and atmospheric conditions. For instance, it has been shown to accumulate in the air if there is a meteorological inversion and little wind. 

High concentrations of radon can be found in some spring waters and hot springs.  The towns of Boulder, Montana; Misasa; Bad Kreuznach, Germany; and the country of Japan have radium-rich springs that emit radon. To be classified as a radon mineral water, radon concentration must be above 2 nCi/L (74 kBq/m3).  The activity of radon mineral water reaches 2,000 kBq/m3 in Merano and 4,000 kBq/m3 in Lurisia (Italy).

Natural radon concentrations in the Earth's atmosphere are so low that radon-rich water in contact with the atmosphere will continually lose radon by volatilization. Hence, ground water has a higher concentration of 222Rn than surface water, because radon is continuously produced by radioactive decay of 226Ra present in rocks. Likewise, the saturated zone of a soil frequently has a higher radon content than the unsaturated zone because of diffusional losses to the atmosphere.  

In 1971, Apollo 15 passed 110 km (68 mi) above the Aristarchus plateau on the Moon, and detected a significant rise in alpha particles thought to be caused by the decay of 222Rn. The presence of 222Rn has been inferred later from data obtained from the Lunar Prospector alpha particle spectrometer. 

Radon is found in some petroleum. Because radon has a similar pressure and temperature curve to propane, and oil refineries separate petrochemicals based on their boiling points, the piping carrying freshly separated propane in oil refineries can become radioactive because of decaying radon and its products. 

Residues from the petroleum and natural gas industry often contain radium and its daughters. The sulfate scale from an oil well can be radium rich, while the water, oil, and gas from a well often contains radon. Radon decays to form solid radioisotopes that form coatings on the inside of pipework.

Accumulation in buildings

High concentrations of radon in homes were discovered by chance in 1985 after the stringent radiation testing conducted at a nuclear power plant entrance revealed that Stanley Watras, an engineer entering the plant, was contaminated by radioactive substances.  Typical domestic exposures are of approximately 100 Bq/m3 (2.7 pCi/L) indoors. Some level of radon will be found in all buildings. Radon mostly enters a building directly from the soil through the lowest level in the building that is in contact with the ground. High levels of radon in the water supply can also increase indoor radon air levels. Typical entry points of radon into buildings are cracks in solid foundations, construction joints, cracks in walls, gaps in suspended floors, gaps around service pipes, cavities inside walls, and the water supply. Radon concentrations in the same location may differ by a factor of two over a period of 1 hour. Also, the concentration in one room of a building may be significantly different from the concentration in an adjoining room.

The distribution of radon concentrations will generally change from room to room, and the readings are averaged according to regulatory protocols. Indoor radon concentration is usually assumed to follow a lognormal distribution on a given territory.  Thus, the geometric mean is generally used for estimating the "average" radon concentration in an area. 

The mean concentration ranges from less than 10 Bq/m3 to over 100 Bq/m3 in some European countries.  Typical geometric standard deviations found in studies range between 2 and 3, meaning (given the 68–95–99.7 rule) that the radon concentration is expected to be more than a hundred times the mean concentration for 2 to 3% of the cases.

The highest average radon concentrations in the United States are found in Iowa and in the Appalachian Mountain areas in southeastern Pennsylvania. Some of the highest readings ever have been recorded in the Irish town of Mallow, County Cork, prompting local fears regarding lung cancer. Iowa has the highest average radon concentrations in the United States due to significant glaciation that ground the granitic rocks from the Canadian Shield and deposited it as soils making up the rich Iowa farmland.  Many cities within the state, such as Iowa City, have passed requirements for radon-resistant construction in new homes.

In a few locations, uranium tailings have been used for landfills and were subsequently built on, resulting in possible increased exposure to radon.

Since radon is a colorless, odorless gas the only way to know how much is present in the air or water is to perform tests. In the United States radon test kits are available to the public at retail stores, such as hardware stores, for home use and testing is available through licensed professionals, who are often home inspectors. Efforts to reduce indoor radon levels are called radon mitigation. In the U.S. the Environmental Protection Agency recommends all houses be tested for radon.

Industrial production

Radon is obtained as a by-product of uraniferous ores processing after transferring into 1% solutions of hydrochloric or hydrobromic acids. The gas mixture extracted from the solutions contains , , He, Rn, ,  and hydrocarbons. The mixture is purified by passing it over copper at 720 °C to remove the  and the , and then KOH and Phosphorus pentoxide| are used to remove the acids and moisture by sorption. Radon is condensed by liquid nitrogen and purified from residue gases by sublimation. 

Radon commercialization is regulated, but it is available in small quantities for the calibration of 222Rn measurement systems, at a price of almost $6,000 per milliliter of radium solution (which only contains about 15 picograms of actual radon at a given moment). 
Radon is produced by a solution of radium-226 (half-life of 1600 years). Radium-226 decays by alpha-particle emission, producing radon that collects over samples of radium-226 at a rate of about 1 mm3/day per gram of radium; equilibrium is quickly achieved and radon is produced in a steady flow, with an activity equal to that of the radium (50 Bq). Gaseous 222Rn (half-life of about four days) escapes from the capsule through diffusion. 

Concentration scale

Applications

Medical

An early-20th-century form of quackery was the treatment of maladies in a radiotorium.  It was a small, sealed room for patients to be exposed to radon for its "medicinal effects". The carcinogenic nature of radon due to its ionizing radiation became apparent later on. Radon's molecule-damaging radioactivity has been used to kill cancerous cells,  but it does not increase the health of healthy cells. The ionizing radiation causes the formation of free radicals, which results in genetic and other cell damage, resulting in increased rates of illness, including cancer.

Exposure to radon, a process known as radiation hormesis, has been suggested to mitigate auto-immune diseases such as arthritis.  As a result, in the late 20th century and early 21st century, "health mines" established in Basin, Montana attracted people seeking relief from health problems such as arthritis through limited exposure to radioactive mine water and radon. The practice is discouraged because of the well-documented ill effects of high-doses of radiation on the body. 

Radioactive water baths have been applied since 1906 in Jáchymov, Czech Republic, but even before radon discovery they were used in Bad Gastein, Austria. Radium-rich springs are also used in traditional Japanese onsen in Misasa, Tottori Prefecture. Drinking therapy is applied in Bad Brambach, Germany. Inhalation therapy is carried out in Gasteiner-Heilstollen, Austria, in Świeradów-Zdrój, Czerniawa-Zdrój, Kowary, Lądek Zdrój, Poland, in Harghita Băi, Romania, and in Boulder, United States. In the US and Europe there are several "radon spas", where people sit for minutes or hours in a high-radon atmosphere in the belief that low doses of radiation will invigorate or energize them. 

Radon has been produced commercially for use in radiation therapy, but for the most part has been replaced by radionuclides made in accelerators and nuclear reactors. Radon has been used in implantable seeds, made of gold or glass, primarily used to treat cancers.
The gold seeds were produced by filling a long tube with radon pumped from a radium source, the tube being then divided into short sections by crimping and cutting. The gold layer keeps the radon within, and filters out the alpha and beta radiations, while allowing the gamma rays to escape (which kill the diseased tissue). The activities might range from 0.05 to 5 millicuries per seed (2 to 200 MBq). The gamma rays are produced by radon and the first short-lived elements of its decay chain (218Po, 214Pb, 214Bi, 214Po).

Radon and its first decay products being very short-lived, the seed is left in place. After 12 half-lives (43 days), radon radioactivity is at 1/2000 of its original level. At this stage, the predominant residual activity originates from the radon decay product 210Pb, whose half-life (22.3 years) is 2000 times that of radon (and whose activity is thus 1/2000 of radon's), and its descendants 210Bi and 210Po.

In the early part of the 20th century in the US, gold contaminated with 210Pb entered the jewelry industry. This was from gold seeds that had held 222Rn that had been melted down after the radon had decayed.  

Scientific

Radon emanation from the soil varies with soil type and with surface uranium content, so outdoor radon concentrations can be used to track air masses to a limited degree. This fact has been put to use by some atmospheric scientists. Because of radon's rapid loss to air and comparatively rapid decay, radon is used in hydrologic research that studies the interaction between ground water and streams. Any significant concentration of radon in a stream is a good indicator that there are local inputs of ground water.

Radon soil-concentration has been used in an experimental way to map buried close-subsurface geological faults because concentrations are generally higher over the faults.  Similarly, it has found some limited use in prospecting for geothermal gradients.[http://www.onepetro.org/mslib/servlet/onepetropreview?id=00008890 Radon Transect Analysis In Geothermal Reservoirs]

Some researchers have investigated changes in groundwater radon concentrations for earthquake prediction.    Radon has a half-life of approximately 3.8 days, which means that it can be found only shortly after it has been produced in the radioactive decay chain. For this reason, it has been hypothesized that increases in radon concentration is due to the generation of new cracks underground, which would allow increased ground water circulation, flushing out radon. The generation of new cracks might not unreasonably be assumed to precede major earthquakes. In the 1970s and 1980s, scientific measurements of radon emissions near faults found that earthquakes often occurred with no radon signal, and radon was often detected with no earthquake to follow. It was then dismissed by many as an unreliable indicator.  As of 2009, it was under investigation as a possible precursor by NASA. 

Radon is a known pollutant emitted from geothermal power stations because it is present in the material pumped from deep underground. It disperses rapidly, and no radiological hazard has been demonstrated in various investigations. In addition, typical systems re-inject the material deep underground rather that releasing it at the surface, so its environmental impact is minimal. 

In the 1940s and '50s, radon was used for industrial radiography,   Other X-ray sources, which became available after World War II, quickly replaced radon for this application, as they were lower in cost and had less hazard of alpha radiation.

Health risks

In mines

Radon-222 decay products have been classified by the International Agency for Research on Cancer as being carcinogenic to humans,  and as a gas that can be inhaled, lung cancer is a particular concern for people exposed to elevated levels of radon for sustained periods. During the 1940s and '50s, when safety standards requiring expensive ventilation in mines were not widely implemented,  radon exposure was linked to lung cancer among non-smoking miners of uranium and other hard rock materials in what is now the Czech Republic, and later among miners from the Southwestern United States    and South Australia.  Despite these hazards being known in the early 1950s,  this occupational hazard remained poorly managed in many mines until the 1970s. During this period, several entrepreneurs opened former uranium mines in the USA to the general public and advertised alleged health benefits from breathing radon gas underground. Health benefits claimed including pain, sinus, asthma and arthritis relief   but these were proven to be false. 

Since that time, ventilation and other measures have been used to reduce radon levels in most affected mines that continue to operate. In recent years, the average annual exposure of uranium miners has fallen to levels similar to the concentrations inhaled in some homes. This has reduced the risk of occupationally induced cancer from radon, although health issues may persist for those who are currently employed in affected mines and for those who have been employed in them in the past. As the relative risk for miners has decreased, so has the ability to detect excess risks among that population.

In addition to lung cancer, researchers have theorized a possible increased risk of leukemia due to radon exposure. Empirical support from studies of the general population is inconsistent, and a study of uranium miners found a correlation between radon exposure and chronic lymphocytic leukemia. 

Miners (as well as milling and ore transportation workers) who worked in the uranium industry in the United States between the 1940s and 1971 may be eligible for compensation under the Radiation Exposure Compensation Act (RECA). Surviving relatives may also apply in cases where the formerly employed person is deceased.

Domestic-level exposure

Radon exposure (mostly radon daughters) has been linked to lung cancer in numerous case-control studies performed in the United States, Europe and China. There are approximately 21,000 deaths per year in the US due to radon-induced lung cancers.  One of the most comprehensive radon studies performed in the United States by Dr. R. William Field and colleagues found a 50% increased lung cancer risk even at the protracted exposures at the EPA's action level of 4 pCi/L. North American and European Pooled analyses further support these findings. 

Most models of residential radon exposure are based on studies of miners, and direct estimates of the risks posed to homeowners would be more desirable. Because of the difficulties of measuring the risk of radon relative to smoking, models of their effect have often made use of them.

Radon has been considered the second leading cause of lung cancer and leading environmental cause of cancer mortality by the United States Environmental Protection Agency.  Others have reached similar conclusions for the United Kingdom  and France. Radon exposure in homes and offices may arise from certain subsurface rock formations, and also from certain building materials (e.g., some granites). The greatest risk of radon exposure arises in buildings that are airtight, insufficiently ventilated, and have foundation leaks that allow air from the soil into basements and dwelling rooms.

Action and reference level

WHO presented in 2009 a recommended reference level (the national reference level), 100 Bq/m3, for radon in dwellings. The recommendation also says that where this is not possible, 300 Bq/m3 should be selected as the highest level. A national reference level should not be a limit, but should represent the maximum acceptable annual average radon concentration in a dwelling. 

The actionable concentration of radon in a home varies depending on the organization doing the recommendation, for example, the United States Environmental Protection Agency encourages that action be taken at concentrations as low as 74 Bq/m3 (2 pCi/L),  and the European Union recommends action be taken when concentrations reach 400 Bq/m3 (11 pCi/L) for old houses and 200 Bq/m3 (5 pCi/L) for new ones.  On 8 July 2010 the UK's Health Protection Agency issued new advice setting a "Target Level" of 100 Bq/m3 whilst retaining an "Action Level" of 200 Bq/m3.  The same levels (as UK) apply to Norway from 2010; in all new housings preventative measures should be taken against radon accumulation.

Relationship to smoking

Results from epidemiological studies indicate that the risk of lung cancer increases with exposure to residential radon. A well-known example of source of error is smoking. In addition, smoking is the most important risk factor for lung cancer. In the West, tobacco smoke is estimated to cause about 90% of all lung cancers. There is a tendency for other hypothetical lung cancer risks to drown in the risk of smoking. Results from epidemiological studies must always be interpreted with caution.

According to the EPA, the risk of lung cancer for smokers is significant due to synergistic effects of radon and smoking. For this population about 62 people in a total of 1,000 will die of lung cancer compared to 7 people in a total of 1,000 for people who have never smoked. It can not be excluded that the risk of non-smokers should be primarily explained by a combination effect of radon and passive smoking (see below).

Radon, like other known or suspected external risk factors for lung cancer, is a threat for smokers and former smokers. This was demonstrated by the European pooling study.  A commentary  to the pooling study stated: "it is not appropriate to talk simply of a risk from radon in homes. The risk is from smoking, compounded by a synergistic effect of radon for smokers. Without smoking, the effect seems to be so small as to be insignificant."

According to the European pooling study, there is a difference in risk from radon between histological types. Small cell lung carcinoma, which practically only affects smokers have high risk from radon. For other histological types such as adenocarcinoma, the type that primarily affects never smokers, the risk from radon appears to be lower. 

A study  of radiation from post mastectomy radiotherapy shows that the simple models previously used to assess the combined and separate risks from radiation and smoking need to be developed. This is also supported by new discussion about the calculation method, LNT, which routinely has been used. 

Relationship to passive smoking

An important question is if also passive smoking can cause a similar synergy effect with residential radon. This has been insufficiently studied. The basic data for the European pooling study makes it impossible to exclude that such synergy effect is an explanation for the (very limited) increase in the risk from radon that was stated for non-smokers.

A study  from 2001, which included 436 cases (never smokers who had lung cancer), and a control group (1649 never smokers) showed that exposure to radon increased the risk of lung cancer in never smokers. But the group that had been exposed to passive smoking at home appeared to bear the entire risk increase, while those who were not exposed to passive smoking did not show any increased risk with increasing radon level.

In drinking water

The effects of radon if ingested are similarly unknown, although studies have found that its biological half-life ranges from 30–70 minutes, with 90 percent removal at 100 minutes. In 1999 National Research Council investigated the issue of radon in drinking water. The risks associated with ingestion was considered almost negligible.  Water from underground sources may contain significant amounts of radon depending on the surrounding rock and soil conditions, whereas surface sources generally do not. 

As well as being ingested through drinking water, radon is also released from water when temperature is increased, pressure is decreased and when water is aerated. Optimum conditions for radon release and exposure occur during showering. Water with a radon concentration of 104 pCi/L can increase the indoor airborne radon concentration by 1 pCi/L under normal conditions. 

Testing and mitigation

There are relatively simple tests for radon gas. In some countries these tests are methodically done in areas of known systematic hazards. Radon detection devices are commercially available. Digital radon detectors provide ongoing measurements giving both daily, weekly, short-term and long-term average readouts via a digital display. Short-term radon test devices used for initial screening purposes are inexpensive, in some cases free. There are very important protocols for taking short-term radon tests and it is imperative that they be strictly followed. The kit includes a collector that the user hangs in the lowest habitable floor of the house for 2 to 7 days. The user then sends the collector to a laboratory for analysis. Long term kits, taking collections for up to one year or more, are also available. An open-land test kit can test radon emissions from the land before construction begins. Radon concentrations can vary daily, and accurate radon exposure estimates require long-term average radon measurements in the spaces where an individual spends a significant amount of time. 

Radon levels fluctuate naturally, due to factors like transient weather conditions, so an initial test might not be an accurate assessment of a home's average radon level. Radon levels are at a maximum during the coolest part of the day when pressure differentials are greatest. Therefore, a high result (over 4 pCi/L) justifies repeating the test before undertaking more expensive abatement projects. Measurements between 4 and 10 pCi/L warrant a long term radon test. Measurements over 10 pCi/L warrant only another short term test so that abatement measures are not unduly delayed. Purchasers of real estate are advised to delay or decline a purchase if the seller has not successfully abated radon to 4 pCi/L or less.

Because the half-life of radon is only 3.8 days, removing or isolating the source will greatly reduce the hazard within a few weeks. Another method of reducing radon levels is to modify the building's ventilation. Generally, the indoor radon concentrations increase as ventilation rates decrease. In a well ventilated place, the radon concentration tends to align with outdoor values (typically 10 Bq/m3, ranging from 1 to 100 Bq/m3).

The four principal ways of reducing the amount of radon accumulating in a house are: 
* Sub-slab depressurization (soil suction) by increasing under-floor ventilation;
* Improving the ventilation of the house and avoiding the transport of radon from the basement into living rooms;
* Installing a radon sump system in the basement;
* Installing a positive pressurization or positive supply ventilation system.

According to the EPA the method to reduce radon "...primarily used is a vent pipe system and fan, which pulls radon from beneath the house and vents it to the outside", which is also called sub-slab depressurization, active soil depressurization, or soil suction. Generally indoor radon can be mitigated by sub-slab depressurization and exhausting such radon-laden air to the outdoors, away from windows and other building openings. "EPA generally recommends methods which prevent the entry of radon. Soil suction, for example, prevents radon from entering your home by drawing the radon from below the home and venting it through a pipe, or pipes, to the air above the home where it is quickly diluted" and "EPA does not recommend the use of sealing alone to reduce radon because, by itself, sealing has not been shown to lower radon levels significantly or consistently". 

Positive-pressure ventilation systems can be combined with a heat exchanger to recover energy in the process of exchanging air with the outside, and simply exhausting basement air to the outside is not necessarily a viable solution as this can actually draw radon gas into a dwelling. Homes built on a crawl space may benefit from a radon collector installed under a "radon barrier" (a sheet of plastic that covers the crawl space). 
For crawlspaces, the EPA states "An effective method to reduce radon levels in crawlspace homes involves covering the earth floor with a high-density plastic sheet. A vent pipe and fan are used to draw the radon from under the sheet and vent it to the outdoors. This form of soil suction is called submembrane suction, and when properly applied is the most effective way to reduce radon levels in crawlspace homes."
Thoron Niton (element) Radon (Rn) Radium emanation Element 86 Radon Radon gas Household radon Emanation (chemistry) Radon poisoning Radon (element) Actinon
household radon thoron radium emanation radon element radon rn element 86 actinon niton element radon poisoning radon radon gas emanation chemistry
radon
WikipediaEntity
Radon
 Which chemical element is represented by the symbol Ar. 

 The objective symbol is represented by the symbol which is a fascinating concept or is still which you can create a mix and provide some feedback or any any questions about the symbol. It looks like like there is no such thing as as the represented by the name is also also that dash or made no which chemical element that represents a and balance for contact or to time the following.g.g.g.G and any other questions or