Isotopes are different types of atoms (nuclides A nuclide is an atomic species characterized by the specific constitution of its nucleus, i.e., by its number of protons Z, its number of neutrons N, and its energy state. Thus, all nuclides are atoms which have at least one electron (though certain ions may be included), but naked nuclei (such as occur in cosmic rays and sufficiently hot plasmas)) of the same chemical element A chemical element is a pure chemical substance consisting of one type of atom distinguished by its atomic number, which is the number of protons in its nucleus. The term is also used to refer to a pure chemical substance composed of atoms with the same number of protons. Common examples of elements are iron, copper, silver, gold, hydrogen, carbon,, each having a different number of neutrons The neutron is a subatomic particle with no net electric charge and a mass slightly larger than that of a proton. They are usually found in atomic nuclei. The nuclei of most atoms consist of protons and neutrons, which are therefore collectively referred to as nucleons. The number of protons in a nucleus is the atomic number and defines the type. In a corresponding manner, isotopes differ in mass number The mass number , also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. Because protons and neutrons both are baryons, the mass number A is identical with the baryon number B as of the nucleus as of the whole atom or ion. The mass number is different for (or number of nucleons A nucleon is a collective name for two baryons: the neutron and the proton in physics. They are constituents of the atomic nucleus and until the 1960s were thought to be elementary particles. In those days their interactions defined strong interactions. Now they are known to be composite particles, made of quarks. Understanding the properties of) but never in atomic number In chemistry and physics, the atomic number is the number of protons found in the nucleus of an atom and therefore identical to the charge number of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely identifies a chemical element. In an atom of neutral charge, the atomic number is also equal to the number of.^[1] The number of protons The proton is a subatomic particle with an electric charge of +1 elementary charge. It is found in the nucleus of each atom, along with neutrons, but is also stable by itself and has a second identity as the hydrogen ion, H+. It is composed of three fundamental particles: two up quarks and one down quark (the atomic number In chemistry and physics, the atomic number is the number of protons found in the nucleus of an atom and therefore identical to the charge number of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely identifies a chemical element. In an atom of neutral charge, the atomic number is also equal to the number of) is the same because that is what characterizes a chemical element. For example, carbon-12 Carbon-12 is the more abundant of the two stable isotopes of the element carbon, accounting for 98.89% of carbon; it contains 6 protons, 6 neutrons, and 6 electrons, carbon-13 Carbon-13 is a natural, stable isotope of carbon and one of the environmental isotopes. It makes up about 1.1% of all natural carbon on Earth and carbon-14 Carbon-14, 14C, or radiocarbon, is a radioactive isotope of carbon discovered on February 27, 1940, by Martin Kamen and Sam Ruben at the University of California Radiation Laboratory in Berkeley, though its existence had been suggested already in 1934 by Franz Kurie. Its nucleus contains 6 protons and 8 neutrons. Its presence in organic materials are three isotopes of the element carbon with mass numbers 12, 13 and 14, respectively. The atomic number In chemistry and physics, the atomic number is the number of protons found in the nucleus of an atom and therefore identical to the charge number of the nucleus. It is conventionally represented by the symbol Z. The atomic number uniquely identifies a chemical element. In an atom of neutral charge, the atomic number is also equal to the number of of carbon is 6, so the neutron numbers Atomic number plus neutron number equals mass number: Z+N=A in these isotopes of carbon are therefore 12−6 = 6, 13−6 = 7, and 14–6 = 8, respectively.

A nuclide A nuclide is an atomic species characterized by the specific constitution of its nucleus, i.e., by its number of protons Z, its number of neutrons N, and its energy state. Thus, all nuclides are atoms which have at least one electron (though certain ions may be included), but naked nuclei (such as occur in cosmic rays and sufficiently hot plasmas) is an atomic nucleus with a specified composition of protons and neutrons. The nuclide concept emphasizes nuclear properties over chemical properties, while the isotope concept emphasizes chemical over nuclear. The neutron number has drastic effects on nuclear properties, but negligible effects on chemical properties. Since isotope is the older term, it is better known, and is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology Nuclear technology is technology that involves the reactions of atomic nuclei. It has found applications from smoke detectors to nuclear reactors, and from gun sights to nuclear weapons.

An isotope and/or nuclide is specified by the name of the particular element (this indicates the atomic number implicitly) followed by a hyphen and the mass number (e.g. helium-3 Helium-3 is a light, non-radioactive isotope of helium with two protons and one neutron. It is rare on Earth, and is sought for use in nuclear fusion research. The abundance of helium-3 is thought to be greater on the Moon (embedded in the upper layer of regolith by the solar wind over billions of years) and the solar system's gas giants (left, carbon-12 Carbon-12 is the more abundant of the two stable isotopes of the element carbon, accounting for 98.89% of carbon; it contains 6 protons, 6 neutrons, and 6 electrons, carbon-13 Carbon-13 is a natural, stable isotope of carbon and one of the environmental isotopes. It makes up about 1.1% of all natural carbon on Earth, iodine-131 Iodine-131 , also called radioiodine, is a radioisotope of iodine which has medical and pharmaceutical uses. It is also a major radioactive hazard in nuclear fission products, and was a significant contributor to the health effects from open-air atomic bomb testing in the 1950's, and from the Chernobyl disaster and uranium-238 Uranium-238 is the most common isotope of uranium found in nature. It is not fissile, but is a fertile material: it can capture a slow neutron and after two beta decays become fissile plutonium-239. U-238 is fissionable by fast neutrons, but cannot support a chain reaction because inelastic scattering reduces neutron energy below the range where). When a chemical symbol Chemical symbols may also be modified by the use of superscripts or subscripts to show a specific isotope of an atom. Additionally superscripts may be used to indicate the ionization or oxidation state of an element is used, e.g., "C" for carbon, standard notation is to indicate the number of nucleons with a superscript A subscript or superscript is a number, figure, symbol, or indicator that appears smaller than the normal line of type and is set slightly below or above it – subscripts appear at or below the baseline, while superscripts are above. Subscripts and superscripts are perhaps best known for their use in formulas, mathematical expressions, and at the upper left of the chemical symbol and to indicate the atomic number with a subscript A subscript or superscript is a number, figure, symbol, or indicator that appears smaller than the normal line of type and is set slightly below or above it – subscripts appear at or below the baseline, while superscripts are above. Subscripts and superscripts are perhaps best known for their use in formulas, mathematical expressions, and at the lower left (e.g. 32He, 42He, 126C, 146C, 23592U, and 23992U).

Some isotopes are radioactive Radioactive decay is the process by which an unstable atomic nucleus loses energy by emitting ionizing particles or radiation. The emission is spontaneous in that the nucleus decays without collision with another particle. This decay, or loss of energy, results in an atom of one type, called the parent nuclide, transforming to an atom of a and are therefore described as radioisotopes A radionuclide is an atom with an unstable nucleus, which is a nucleus characterized by excess energy which is available to be imparted either to a newly-created radiation particle within the nucleus, or else to an atomic electron . The radionuclide, in this process, undergoes radioactive decay, and emits a gamma ray(s) and/or subatomic particles or radionuclides A radionuclide is an atom with an unstable nucleus, which is a nucleus characterized by excess energy which is available to be imparted either to a newly-created radiation particle within the nucleus, or else to an atomic electron . The radionuclide, in this process, undergoes radioactive decay, and emits a gamma ray(s) and/or subatomic particles, while others have never been observed to undergo radioactive decay and are described as stable isotopes Stable isotopes are chemical isotopes that are not radioactive . By this definition, there are 256 known stable isotopes of the 80 elements, which have one or more stable nuclides. A list of these is given at the end of this article. About two thirds of the elements have more than one stable isotope. One element (tin) has ten stable isotopes. For example, ¹⁴C is a radioactive form of carbon while ¹²C and ¹³C are stable isotopes. There are about 339 naturally occurring nuclides on Earth^[2], of which 288 are primordial nuclides and 259 are "stable Stable isotopes are chemical isotopes that are not radioactive . By this definition, there are 256 known stable isotopes of those 80 elements which have one or more stable nuclides. A list of these is given at the end of this article. Of these 80, twenty-six have only a single stable isotope, and are thus termed monoisotopic, and the rest have"^[2]. However, some apparently "stable" isotopes are predicted by theory to be radioactive with very long half-lives.^{[citation needed]} Adding in the radioactive nuclides that have been created artificially, there are more than 3100 currently known nuclides.^[3]

Contents 1 History of the term 2 Variation in properties between isotopes 2.1 Chemical and atomic properties 2.2 Nuclear properties and stability 2.3 Isotopes per element 2.4 Even and odd 2.4.1 Even mass number 2.4.1.1 Even proton-even neutron 2.4.1.2 Odd proton-odd neutron 2.4.2 Odd mass number 2.4.2.1 Odd proton-even neutron 2.4.2.2 Even proton-odd neutron 2.4.3 Odd neutron number 3 Occurrence in nature 4 Atomic mass of isotopes 5 Applications of isotopes 5.1 Use of chemical and biological properties 5.2 Use of nuclear properties 6 See also 7 References 8 External links

History of the term

In the bottom right corner of JJ Thomson's photographic plate are the separate impact marks for the two isotopes of neon Neon is the chemical element that has the symbol Ne and an atomic number of 10. Although a very common element in the universe, it is rare on Earth. A colorless, inert noble gas under standard conditions, neon gives a distinct reddish-orange glow when used in discharge tubes and neon lamps and advertising signs. It is commercially extracted from: neon-20 and neon-22.

The term isotope was coined in 1913 by Margaret Todd, a Scottish physician, during a conversation with Frederick Soddy Frederick Soddy was an English radiochemist who explained, with Ernest Rutherford, that radioactivity is due to the transmutation of elements, now known to involve nuclear reactions. He also proved the existence of isotopes of certain radioactive elements. He received the Nobel Prize for Chemistry in 1921, and has a crater named for him on the far (to whom she was distantly related by marriage).^[4] Soddy, a chemist at Glasgow University The University of Glasgow is the fourth-oldest university in the English-speaking world and one of Scotland's four ancient universities. Located in Glasgow, the university was founded in 1451 and is presently one of seventeen British higher education institutions ranked amongst the top 100 of the world, explained that it appeared from his investigations as if each position in the periodic table The periodic table of the chemical elements is a tabular display of the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table was occupied by multiple entities. Hence Todd made the suggestion, which Soddy adopted, that a suitable name for such an entity would be the Greek term for "at the same place".

Soddy's own studies were of radioactive (unstable) atoms. The first observation of different stable isotopes for an element was by J. J. Thomson Sir Joseph John “J. J.” Thomson, OM, FRS was a British physicist and Nobel laureate. He is credited for the discovery of the electron and of isotopes, and the invention of the mass spectrometer. Thomson was awarded the 1906 Nobel Prize in Physics for the discovery of the electron and for his work on the conduction of electricity in gases in 1913. As part of his exploration into the composition of canal rays, Thomson channeled streams of neon Neon is the chemical element that has the symbol Ne and an atomic number of 10. Although a very common element in the universe, it is rare on Earth. A colorless, inert noble gas under standard conditions, neon gives a distinct reddish-orange glow when used in discharge tubes and neon lamps and advertising signs. It is commercially extracted from ions through a magnetic and an electric field and measured their deflection by placing a photographic plate in their path. Each stream created a glowing patch on the plate at the point it struck. Thomson observed two separate patches of light on the photographic plate (see image), which suggested two different parabolas of deflection. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest. F.W. Aston Francis Aston was born in Harborne, now part of Greater Birmingham, on September 1, 1877. He was the third child and second son of William Aston and Fanny Charlotte Hollis. He was educated at the Harborne Vicarage School and later Malvern College in Worcestershire where he was a boarder. In 1893 Francis William Aston began his university studies subsequently discovered different stable isotopes for numerous elements using a mass spectrograph.

Variation in properties between isotopes

Chemical and atomic properties

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and share a similar electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium A hydrogen atom is an atom of the chemical element hydrogen. The electrically neutral atom contains a single positively-charged proton and a single negatively-charged electron bound to the nucleus by the Coulomb force. The most abundant isotope, hydrogen-1, protium, or light hydrogen, contains no neutrons; other isotopes of hydrogen, such as (¹H) and deuterium Deuterium, also called heavy hydrogen, is a stable isotope of hydrogen with a natural abundance in the oceans of Earth of approximately one atom in 6,500 of hydrogen . Deuterium thus accounts for approximately 0.0154% (alternately, on a mass basis: 0.0308%) of all naturally occurring hydrogen in the oceans on Earth (see VSMOW; the abundance (²H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass Reduced mass is the "effective" inertial mass appearing in the two-body problem of Newtonian mechanics. This is a quantity with the unit of mass, which allows the two-body problem to be solved as if it were a one-body problem. Note however that the mass determining the gravitational force is not reduced. In the computation one mass can) of the atomic systems. However, for heavier elements, which have more neutrons than lighter elements, the ratio of the nuclear mass to the collective electronic mass is far greater, and the relative mass difference between isotopes is much less. For these two reasons, the mass-difference effects on chemistry are usually negligible.

Isotope half lifes. Note that the plot for stable isotopes diverges from the line, protons Z = neutrons N as the element number Z becomes larger

In similar manner, two molecules A molecule is defined as an electrically neutral group of at least two atoms in a definite arrangement held together by very strong chemical bonds. Molecules are distinguished from polyatomic ions in this strict sense. In organic chemistry and biochemistry, the term molecule is used less strictly and also is applied to charged organic molecules that differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. As a consequence, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons In physics, a photon is an elementary particle, the quantum of the electromagnetic interaction and the basic unit of light and all other forms of electromagnetic radiation. It is also the force carrier for the electromagnetic force. The effects of this force are easily observable at both the microscopic and macroscopic level, because the photon of corresponding energies, isotopologues have different optical properties in the infrared Infrared light is electromagnetic radiation with a wavelength between 0.7 and 300 micrometres, which equates to a frequency range between approximately 1 and 430 THz range.

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Nuclear properties and stability

See also: Stable isotope Stable isotopes are chemical isotopes that are not radioactive . By this definition, there are 256 known stable isotopes of those 80 elements which have one or more stable nuclides. A list of these is given at the end of this article. Of these 80, twenty-six have only a single stable isotope, and are thus termed monoisotopic, and the rest have, List of elements by nuclear stability, and List of elements by stability of isotopes Atomic nuclei consist of protons and neutrons, which attract each other through the strong nuclear force, while protons repel each other via the electric force due to their positive charge. These two forces compete, leading to some combinations of neutrons and protons being more stable than others. Neutrons stabilize the nucleus, because they

Atomic nuclei consist of protons and neutrons bound together by the residual strong force The nuclear force is the force between two or more nucleons. It is responsible for binding of protons and neutrons into atomic nuclei. To a large extent, this force can be understood in terms of the exchange of virtual light mesons, such as the pions. Sometimes the nuclear force is called the residual strong force, in contrast to the strong. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their copresence pushes protons slightly apart, reducing the electrostatic repulsion between the protons, and they exert the attractive nuclear force on each other and on protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, so does the ratio of neutrons to protons necessary to ensure a stable nucleus (see graph at right). For example, although the neutron:proton ratio of 32He is 1:2, the neutron:proton ratio of 23892U is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 4020Ca (calcium-40) is the heaviest stable nuclide with the same number of neutrons and protons; all heavier stable nuclides contain more neutrons than protons.

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Isotopes per element

Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin Tin is a chemical element with the symbol Sn and atomic number 50. It is a main group metal in group 14 of the periodic table. Tin shows chemical similarity to both neighboring group 14 elements, germanium and lead, like the two possible oxidation states +2 and +4. Tin is the 49th most abundant element and has, with 10 stable isotopes, the largest). Xenon is the only element that has nine stable isotopes. Cadmium has eight stable isotopes. Five elements have seven stable isotopes, eight have six stable isotopes, nine have five stable isotopes, nine have four stable isotopes, nine have three stable isotopes, 16 have two stable isotopes (counting 180m73Ta as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic elements, having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well).^[5] In total, there are 256 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 256/80 = 3.20 isotopes per element.

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Even and odd

Even/odd N

Mass	E	O	All
Stable	145	101	246
Longlived	20	6	26
Primordial	165	107	272

The proton:neutron ratio is not the only factor affecting nuclear stability. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture Neutron capture is a kind of nuclear reaction in which an atomic nucleus collides with one or more neutrons and they merge to form a heavier nucleus. Since neutrons have no electric charge, they can enter a nucleus more easily than charged particles which are repelled by electrostatic repulsion cross-sections and gamma spectroscopy and nuclear magnetic resonance Nuclear magnetic resonance is a property that magnetic nuclei have in a magnetic field and applied electromagnetic (EM) pulse or pulses, which cause the nuclei to absorb energy from the EM pulse and radiate this energy back out. The energy radiated back out is at a specific resonance frequency which depends on the strength of the magnetic field properties.

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Even mass number

Beta decay of an even-even nucleus produces an odd-odd nucleus, and vice versa. An even number of protons or of neutrons are more stable (lower binding energy Binding energy is the mechanical energy required to disassemble a whole into separate parts. A bound system has typically a lower potential energy than its constituent parts; this is what keeps the system together. The usual convention is that this corresponds to a positive binding energy) because of pairing effects, so even-even nuclei are much more stable than odd-odd. One effect is that there are few stable odd-odd nuclei, but another effect is to prevent beta decay of many even-even nuclei into another even-even nucleus of the same mass number but lower energy, because decay proceeding one step at a time would have to pass through an odd-odd nucleus of higher energy. This makes for a larger number of stable even-even nuclei, up to three for some mass numbers, and up to seven for some atomic (proton) numbers Atomic nuclei consist of protons and neutrons, which attract each other through the strong nuclear force, while protons repel each other via the electric force due to their positive charge. These two forces compete, leading to some combinations of neutrons and protons being more stable than others. Neutrons stabilize the nucleus, because they. Double beta decay In double-beta decay, two neutrons in the nucleus are converted to protons, and two electrons and two electron antineutrinos are emitted. In the process of beta minus decay, unstable nuclei decay by converting a neutron in the nucleus to a proton and emitting an electron and an electronic antineutrino. In order for beta decay to be possible, the directly from even-even to even-even skipping over an odd-odd nuclide is only occasionally possible, and even then with a half-life greater than a billion times the age of the universe.

Even-mass-number nuclides have integer spin and are bosons.

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Even proton-even neutron

Even/odd Z, N

p,n	EE	OO	EO	OE
Stable	140	5	53	48
Longlived	16	4	2	4
Primordial	156	9	55	52

For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five or eight nucleons from existing for long enough to serve as platforms for the buildup of heavier elements during fusion formation in stars (see triple alpha process).

There are 140 stable even-even isotopes, forming 55% of the 256 stable isotopes. There are also 16 primordial longlived even-even isotopes. As a result, many of the 41 even-numbered elements from 2 to 82 have many primordial isotopes. Half have six or more.

All even-even nuclides have spin 0 in their ground state.

Odd proton-odd neutron

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons: the first four odd-odd nuclides 21H, 63Li, 105B, and 147N (where changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio) and 180m73Ta, which has not yet been observed to decay despite experimental attempts^[6]. Also, four long-lived radioactive odd-odd nuclides (4019K, 5023V, 13857La, 17671Lu) occur naturally.

Of these 9 primordial odd-odd nuclides, only 147N is the most common isotope of a common element, because it is a part of the CNO cycle; 63Li and 105B are minority isotopes of elements that are rare compared to other light elements, while the other six isotopes make up only a tiny percentage of their elements.

Few odd-odd nuclides (and none of the primordial ones) have spin 0 in the ground state.

Odd mass number

There is only one beta-stable nuclide per odd mass number because there is no difference in binding energy between even-odd and odd-even comparable to that between even-even and odd-odd, and other nuclides of the same mass are free to beta decay towards the lowest-energy one. For mass numbers 5, 147, 151, and 209 and up, the one beta-stable isobar is able to alpha decay, so that there are no stable isotopes with these mass numbers. This gives a total of 101 stable isotopes with odd mass numbers.

Odd-mass-number nuclides have half-integer spin and are fermions.

Odd proton-even neutron

These form most of the stable isotopes of the odd-numbered elements, but there is only one stable odd-even isotope for each of the 41 odd-numbered elements from 1 to 81, except for technetium (₄₃Tc) and promethium (₆₁Pm) that have no stable isotopes, and chlorine (₁₇Cl), potassium (₁₉K), copper (₂₉Cu), gallium (₃₁Ga), bromine (₃₅Br), silver (₄₇Ag), antimony (₅₁Sb), iridium (Ir), and thallium (₈₁Tl), each of which has two, making a total of 48 stable odd-even isotopes. There are also four primordial long-lived odd-even isotopes, 8737Rb, 11549In, 15163Eu, and 18775Re.

Even proton-odd neutron

There are 54 stable isotopes that have an even number of protons and an odd number of neutrons. There are also four primordial long lived even-odd isotopes, 11348Cd (beta decay, half-life is 7.7 × 10¹⁵ years); 14762Sm (1.06 × 10¹¹a); and 14962Sm (>2 × 10¹⁵a); and the fissile 23592U.

The only even-odd isotopes that are the most common one for their element are 19578Pt and 94Be. Beryllium-9 is the only stable beryllium isotope because the expected beryllium-8 has higher energy than two alpha particles and therefore decays to them.

Odd neutron number

Even/odd N

n	E	O
Stable	188	58
Longlived	20	6
Primordial	208	64

The only odd-neutron-number isotopes that are the most common isotope of their element are 19578Pt, 94Be and 147N.

Actinides with odd neutron number are generally fissile, while those with even neutron number are generally not, though they are split when bombarded with fast neutrons.

Occurrence in nature

Elements are composed of one or more naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial, in which case they have persisted down to the present because their rate of decay is so slow (e.g., uranium-238 and potassium-40), or they are postprimordial, created by cosmic ray bombardment (e.g., tritium, carbon-14) or by the decay of a primordial isotope (e.g., 22688Ra to 22286Rn).

As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (₅₀Sn). There are about 94 elements found naturally on Earth (up to plutonium (₉₄Pu, inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimate that the elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total.^[7] Only 256 of these naturally occurring isotopes are stable in the sense of never having been observed to decay as of the present time. All the known stable isotopes occur naturally on Earth); the other 85 naturally occurring isotopes are radioactive, but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. An additional ~2700 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.

The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. Before the discovery of isotopes, empirically determined noninteger values of atomic mass confounded scientists. For example, a sample of Chlorine contains 75.8% 3517Cl and 24.2% 3717Cl, giving an average atomic mass of 35.5.

According to generally accepted cosmology theory, only isotopes of hydrogen and helium, and traces of some isotopes of lithium and beryllium were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced isotopes. (See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.

Atomic mass of isotopes

The atomic mass (m_r) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see mass defect), the slight difference in mass between proton and neutron, and the mass of the electrons associated with the atom, the latter because the electron:nucleon ratio differs among isotopes.

The mass number is a dimensionless quantity. The atomic mass, on the other hand, is measured using the atomic mass unit based on the mass of the carbon atom. It is denoted with symbols "u" (for unit) or "Da" (for Dalton).

The atomic masses of naturally occurring isotopes of an element determine the atomic weight of the element. When the element contains N isotopes, the equation below is applied for the atomic weight M:

M = m₁x₁ + m₂x₂ + ... + m_Nx_N

where m₁, m₂, ..., m_N are the atomic masses of each individual isotope, and x₁, ... , x_N are the relative abundances of these isotopes.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element. Isotope separation is a significant technological challenge, particularly with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen, and oxygen are commonly separated by gas diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual since it is based on chemical rather than physical properties, for example in the Girdler sulfide process. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectroscopy.

Use of chemical and biological properties

• Isotope analysis is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration of food products.^[8] The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.^[9]
• Another common application is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy. For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling).
• A technique similar to radioisotopic labeling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.
• Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.

Use of nuclear properties

• Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are ¹H, ²D,¹⁵N, ¹³C, and ³¹P.
• Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as ⁵⁷Fe.
• Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes.

See also

• Atom
• Table of nuclides
• Radionuclide (or radioisotope)
• Nuclear medicine (includes medical isotopes)
• Isotopomer
• List of particles
• Isotopes are nuclides having the same number of protons; compare:
  • Isotones are nuclides having the same number of neutrons.
  • Isobars are nuclides having the same mass number, i.e. sum of protons plus neutrons.
  • Nuclear isomers are different excited states of the same type of nucleus. A transition from one isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. (Not to be confused with chemical isomers.)
• Bainbridge mass spectrometer

References

1. ^ IUPAC http://goldbook.iupac.org/I03331.html
2. ^ ^{a b} "Radioactives Missing From The Earth". http://www.don-lindsay-archive.org/creation/isotope_list.html.
3. ^ "NuDat 2 Description". http://www.nndc.bnl.gov/nudat2/help/index.jsp.
4. ^ Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews 25 (1): 146–57. doi:10.1002/mas.20061. PMID 16134128.
5. ^ Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brook haven National Laboratory. http://www.nndc.bnl.gov/chart/.
6. ^ hhttp://bryza.if.uj.edu.pl/zdfk/wp-includes/publications/misiaszek_180mTa_2009.pdf Search for the radioactivity of 180mTa using an underground HPGe sandwich spectrometer, 2009
7. ^ [1]
8. ^ E. Jamin et al. (2003). "Improved Detection of Added Water in Orange Juice by Simultaneous Determination of the Oxygen-18/Oxygen-16 Isotope Ratios of Water and Ethanol Derived from Sugars"". J. Agric. Food Chem. 51: 5202. doi:10.1021/jf030167 m. http://pubs.acs.org/cgi-bin/article.cgi/jafcau/2003/51/i18/pdf/jf030167 m.pdf.
9. ^ A. H. Treiman, J. D. Gleason and D. D. Bogard (2000). ""The SNC meteorites are from Mars"". Planet. Space. Sci. 48: 1213. doi:10.1016/S0032-0633(00)00105-7. http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6V6T-41WBDHD-8&_user=2400262&_coverDate=10%2F31%2F2000&_alid=678948366&_rdoc=3&_fmt=summary&_orig=search&_cdi=5823&_sort=r&_docanchor=&view=c&_ct=89&_acct=C000057185&_version=1&_urlVersion=0&_userid=2400262&md5=c5ae2aa8ea60dbd76c2870048730a299.

External links

• Nucleonica Nuclear Science Portal
• Nucleonica Nuclear Science Wiki
• International Atomic Energy Agency
• Atomic weights of all isotopes
• Atomgewichte, Zerfallsenergien und Halbwertszeiten aller Isotope
• Chart of the Nuclides produced by the Knolls Atomic Power Laboratory $25
• Exploring the Table of the Isotopes at the LBNL
• Current isotope research and information
• Radioactive Isotopes by the CDC
• Interacive Chart of the nuclides, isotopes and Periodic Table
• The LIVEChart of Nuclides - IAEA with isotope data, in Java or HTML

Categories: Isotopes | Nuclear chemistry | Nuclear physics

 

 

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