Isotope

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The three naturally-occurring isotopes of hydrogen. The fact that each isotope has one proton makes them all variants of hydrogen: the identity of the isotope is given by the number of neutrons. From left to right, the isotopes are protium ( H) with zero neutrons, deuterium ( H) with one neutron, and tritium ( H) with two neutrons. Hydrogen Deuterium Tritium Nuclei Schmatic-en.svg
The three naturally-occurring isotopes of hydrogen. The fact that each isotope has one proton makes them all variants of hydrogen: the identity of the isotope is given by the number of neutrons. From left to right, the isotopes are protium ( H) with zero neutrons, deuterium ( H) with one neutron, and tritium ( H) with two neutrons.

Isotopes are variants of a particular chemical element which differ in neutron number, and consequently in nucleon number. All isotopes of a given element have the same number of protons but different numbers of neutrons in each atom. [1]

Chemical element a species of atoms having the same number of protons in the atomic nucleus

A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons.

Neutron number number of neutrons in a nuclide

The neutron number, symbol N, is the number of neutrons in a nuclide.

Proton nucleon (constituent of the nucleus of the atom) that has positive electric charge; symbol p

A proton is a subatomic particle, symbol
p
or
p+
, with a positive electric charge of +1e elementary charge and a mass slightly less than that of a neutron. Protons and neutrons, each with masses of approximately one atomic mass unit, are collectively referred to as "nucleons".

Contents

The term isotope is formed from the Greek roots isos (ἴσος "equal") and topos (τόπος "place"), meaning "the same place"; thus, the meaning behind the name is that different isotopes of a single element occupy the same position on the periodic table. [2] It was coined by a Scottish doctor and writer Margaret Todd in 1913 in a suggestion to chemist Frederick Soddy.

Periodic table Tabular arrangement of the chemical elements ordered by atomic number

The periodic table, also known as the periodic table of elements, is a tabular display of the chemical elements, which are arranged by atomic number, electron configuration, and recurring chemical properties. The structure of the table shows periodic trends. The seven rows of the table, called periods, generally have metals on the left and non-metals on the right. The columns, called groups, contain elements with similar chemical behaviours. Six groups have accepted names as well as assigned numbers: for example, group 17 elements are the halogens; and group 18 are the noble gases. Also displayed are four simple rectangular areas or blocks associated with the filling of different atomic orbitals.

Margaret Todd (doctor) Scottish medical doctor, schoolteacher and writer

Margaret Georgina Todd was a Scottish doctor and writer. She coined the term isotope in 1913 in a suggestion to chemist Frederick Soddy.

Frederick Soddy chemist and physicist from England

Frederick Soddy FRS 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.

The number of protons within the atom's nucleus is called atomic number and is equal to the number of electrons in the neutral (non-ionized) atom. Each atomic number identifies a specific element, but not the isotope; an atom of a given element may have a wide range in its number of neutrons. The number of nucleons (both protons and neutrons) in the nucleus is the atom's mass number, and each isotope of a given element has a different mass number.

Atomic nucleus core of the atom; composed of bound nucleons (protons and neutrons)

The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment. After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. An atom is composed of a positively-charged nucleus, with a cloud of negatively-charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud. Protons and neutrons are bound together to form a nucleus by the nuclear force.

Atomic number number of protons found in the nucleus of an atom

The atomic number or proton number of a chemical element is the number of protons found in the nucleus of every atom of that element. The atomic number uniquely identifies a chemical element. It is identical to the charge number of the nucleus. In an uncharged atom, the atomic number is also equal to the number of electrons.

Electron subatomic particle with negative electric charge

The electron is a subatomic particle, symbol
e
or
β
, whose electric charge is negative one elementary charge. Electrons belong to the first generation of the lepton particle family, and are generally thought to be elementary particles because they have no known components or substructure. The electron has a mass that is approximately 1/1836 that of the proton. Quantum mechanical properties of the electron include an intrinsic angular momentum (spin) of a half-integer value, expressed in units of the reduced Planck constant, ħ. Being fermions, no two electrons can occupy the same quantum state, in accordance with the Pauli exclusion principle. Like all elementary particles, electrons exhibit properties of both particles and waves: they can collide with other particles and can be diffracted like light. The wave properties of electrons are easier to observe with experiments than those of other particles like neutrons and protons because electrons have a lower mass and hence a longer de Broglie wavelength for a given energy.

For example, carbon-12, carbon-13, and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, and 14, respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7, and 8 respectively.

Carbon-12 (12C) is the more abundant of the two stable isotopes of carbon, amounting to 98.93% of the element carbon; its abundance is due to the triple-alpha process by which it is created in stars. Carbon-12 is of particular importance in its use as the standard from which atomic masses of all nuclides are measured, thus, its atomic mass is exactly 12 daltons by definition. Carbon-12 is composed of 6 protons, 6 neutrons, and 6 electrons.

Carbon-13 (13C) is a natural, stable isotope of carbon with a nucleus containing six protons and seven neutrons. As one of the environmental isotopes, it makes up about 1.1% of all natural carbon on Earth.

Carbon-14 (14C), or radiocarbon, is a radioactive isotope of carbon with an atomic nucleus containing 6 protons and 8 neutrons. Its presence in organic materials is the basis of the radiocarbon dating method pioneered by Willard Libby and colleagues (1949) to date archaeological, geological and hydrogeological samples. Carbon-14 was discovered on February 27, 1940, by Martin Kamen and Sam Ruben at the University of California Radiation Laboratory in Berkeley, California. Its existence had been suggested by Franz Kurie in 1934.

Isotope vs. nuclide

A nuclide is a species of an atom with a specific number of protons and neutrons in the nucleus, for example carbon-13 with 6 protons and 7 neutrons. The nuclide concept (referring to individual nuclear species) emphasizes nuclear properties over chemical properties, whereas the isotope concept (grouping all atoms of each element) emphasizes chemical over nuclear. The neutron number has large effects on nuclear properties, but its effect on chemical properties is negligible for most elements. Even in the case of the lightest elements where the ratio of neutron number to atomic number varies the most between isotopes it usually has only a small effect, although it does matter in some circumstances (for hydrogen, the lightest element, the isotope effect is large enough to strongly affect biology). The term isotopes (originally also isotopic elements [3] , now sometimes isotopic nuclides [4] ) is intended to imply comparison (like synonyms or isomers ), for example: the nuclides 12
6
C
, 13
6
C
, 14
6
C
are isotopes (nuclides with the same atomic number but different mass numbers [5] ), but 40
18
Ar
, 40
19
K
, 40
20
Ca
are isobars (nuclides with the same mass number [6] ). However, because isotope is the older term, it is better known than nuclide, and is still sometimes used in contexts where nuclide might be more appropriate, such as nuclear technology and nuclear medicine.

Nuclide atomic species characterized by the specific constitution of its nucleus: mass number, atomic number, and, where long-enough-lived to observe, the nuclear energy state

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 nuclear energy state.

Chemistry scientific discipline

Chemistry is the scientific discipline involved with elements and compounds composed of atoms, molecules and ions: their composition, structure, properties, behavior and the changes they undergo during a reaction with other substances.

Synonym Words or phrases having the same meaning

A synonym is a word or phrase that means exactly or nearly the same as another lexeme in the same language. Words that are synonyms are said to be synonymous, and the state of being a synonym is called synonymy. For example, the words begin, start, commence, and initiate are all synonyms of one another. Words are typically synonymous in one particular sense: for example, long and extended in the context long time or extended time are synonymous, but long cannot be used in the phrase extended family. Synonyms with exactly the same meaning share a seme or denotational sememe, whereas those with inexactly similar meanings share a broader denotational or connotational sememe and thus overlap within a semantic field. The former are sometimes called cognitive synonyms and the latter, near-synonyms, plesionyms or poecilonyms.

Notation

An isotope and/or nuclide is specified by the name of the particular element (this indicates the atomic number) followed by a hyphen and the mass number (e.g. helium-3, helium-4, carbon-12, carbon-14, uranium-235 and uranium-239). [7] When a chemical symbol is used, e.g. "C" for carbon, standard notation (now known as "AZE notation" because A is the mass number, Z the atomic number, and E for element) is to indicate the mass number (number of nucleons) with a superscript at the upper left of the chemical symbol and to indicate the atomic number with a subscript at the lower left (e.g. 3
2
He
, 4
2
He
, 12
6
C
, 14
6
C
, 235
92
U
, and 239
92
U
). [8] Because the atomic number is given by the element symbol, it is common to state only the mass number in the superscript and leave out the atomic number subscript (e.g. 3
He
, 4
He
, 12
C
, 14
C
, 235
U
, and 239
U
). The letter m is sometimes appended after the mass number to indicate a nuclear isomer, a metastable or energetically-excited nuclear state (as opposed to the lowest-energy ground state), for example 180m
73
Ta
(tantalum-180m).

Helium-3 light, non-radioactive isotope of helium with two protons and one neutron (common helium having two neutrons)

Helium-3 is a light, non-radioactive isotope of helium with two protons and one neutron. Other than protium, helium-3 is the only stable isotope of any element with more protons than neutrons. Helium-3 was discovered in 1939.

Helium-4 isotope of helium

Helium-4 is a stable isotope of the element helium. It is by far the more abundant of the two naturally occurring isotopes of helium, making up about 99.99986% of the helium on Earth. Its nucleus is identical to an alpha particle, and consists of two protons and two neutrons.

Uranium-235 isotope of uranium

Uranium-235 (235U) is an isotope of uranium making up about 0.72% of natural uranium. Unlike the predominant isotope uranium-238, it is fissile, i.e., it can sustain a fission chain reaction. It is the only fissile isotope that is primordial and found in relatively significant quantities in nature.

The common pronunciation of the AZE notation is different from how it is written: 4
2
He
is commonly pronounced as helium-four instead of four-two-helium, and 235
92
U
as uranium two-thirty-five (American English) or uranium-two-three-five (British) instead of 235-92-uranium.

Radioactive, primordial, and stable isotopes

Some isotopes/nuclides are radioactive, and are therefore referred to as radioisotopes or radionuclides, whereas others have never been observed to decay radioactively and are referred to as stable isotopes or stable nuclides. For example, 14
C
is a radioactive form of carbon, whereas 12
C
and 13
C
are stable isotopes. There are about 339 naturally occurring nuclides on Earth, [9] of which 286 are primordial nuclides, meaning that they have existed since the Solar System's formation.

Primordial nuclides include 34 nuclides with very long half-lives (over 100 million years) and 252 that are formally considered as "stable nuclides", [9] because they have not been observed to decay. In most cases, for obvious reasons, if an element has stable isotopes, those isotopes predominate in the elemental abundance found on Earth and in the Solar System. However, in the cases of three elements (tellurium, indium, and rhenium) the most abundant isotope found in nature is actually one (or two) extremely long-lived radioisotope(s) of the element, despite these elements having one or more stable isotopes.

Theory predicts that many apparently "stable" isotopes/nuclides are radioactive, with extremely long half-lives (discounting the possibility of proton decay, which would make all nuclides ultimately unstable). Some stable nuclides are in theory energetically susceptible to other known forms of decay, such as alpha decay or double beta decay, but no decay products have yet been observed, and so these isotopes are said to be "observationally stable". The predicted half-lives for these nuclides often greatly exceed the estimated age of the universe, and in fact there are also 31 known radionuclides (see primordial nuclide) with half-lives longer than the age of the universe.

Adding in the radioactive nuclides that have been created artificially, there are 3,339 currently known nuclides. [10] These include 905 nuclides that are either stable or have half-lives longer than 60 minutes. See list of nuclides for details.

History

Radioactive isotopes

The existence of isotopes was first suggested in 1913 by the radiochemist Frederick Soddy, based on studies of radioactive decay chains that indicated about 40 different species referred to as radioelements (i.e. radioactive elements) between uranium and lead, although the periodic table only allowed for 11 elements from uranium to lead. [11] [12] [13]

Several attempts to separate these new radioelements chemically had failed. [14] For example, Soddy had shown in 1910 that mesothorium (later shown to be 228Ra), radium (226Ra, the longest-lived isotope), and thorium X (224Ra) are impossible to separate. [15] Attempts to place the radioelements in the periodic table led Soddy and Kazimierz Fajans independently to propose their radioactive displacement law in 1913, to the effect that alpha decay produced an element two places to the left in the periodic table, whereas beta decay emission produced an element one place to the right. [16] Soddy recognized that emission of an alpha particle followed by two beta particles led to the formation of an element chemically identical to the initial element but with a mass four units lighter and with different radioactive properties.

Soddy proposed that several types of atoms (differing in radioactive properties) could occupy the same place in the table. [13] For example, the alpha-decay of uranium-235 forms thorium-231, whereas the beta decay of actinium-230 forms thorium-230. [14] The term "isotope", Greek for "at the same place", [13] was suggested to Soddy by Margaret Todd, a Scottish physician and family friend, during a conversation in which he explained his ideas to her. [15] [17] [18] [19] [20] [21] He won the 1921 Nobel Prize in Chemistry in part for his work on isotopes. [22]

In the bottom right corner of J. J. Thomson's photographic plate are the separate impact marks for the two isotopes of neon: neon-20 and neon-22. Discovery of neon isotopes.JPG
In the bottom right corner of J. J. Thomson's photographic plate are the separate impact marks for the two isotopes of neon: neon-20 and neon-22.

In 1914 T. W. Richards found variations between the atomic weight of lead from different mineral sources, attributable to variations in isotopic composition due to different radioactive origins. [14] [22]

Stable isotopes

The first evidence for multiple isotopes of a stable (non-radioactive) element was found by J. J. Thomson in 1913 as part of his exploration into the composition of canal rays (positive ions). [23] [24] Thomson channeled streams of neon 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 subsequently discovered multiple stable isotopes for numerous elements using a mass spectrograph. In 1919 Aston studied neon with sufficient resolution to show that the two isotopic masses are very close to the integers 20 and 22, and that neither is equal to the known molar mass (20.2) of neon gas. This is an example of Aston's whole number rule for isotopic masses, which states that large deviations of elemental molar masses from integers are primarily due to the fact that the element is a mixture of isotopes. Aston similarly showed[ when? ] that the molar mass of chlorine (35.45) is a weighted average of the almost integral masses for the two isotopes 35Cl and 37Cl. [25]

Variation in properties between isotopes

Chemical and molecular 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 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 by far for protium (1
H
), deuterium (2
H
), and tritium (3
H
), because deuterium has twice the mass of protium and tritium has three times the mass of protium. These mass differences also affect the behavior of their respective chemical bonds, by changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements the relative mass difference between isotopes is much less, so that the mass-difference effects on chemistry are usually negligible. (Heavy elements also have relatively more neutrons than lighter elements, so the ratio of the nuclear mass to the collective electronic mass is slightly greater.)

Isotope half-lives. The plot for stable isotopes diverges from the line Z = N as the element number Z becomes larger Isotopes and half-life.svg
Isotope half-lives. The plot for stable isotopes diverges from the line Z = N as the element number Z becomes larger

Similarly, two molecules that differ only in the isotopes of their atoms (isotopologues) have identical electronic structure, and therefore almost indistinguishable physical and chemical properties (again with deuterium and tritium being the primary exceptions). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms; so different isotopologues have different sets of vibrational modes. Because vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

Nuclear properties and stability

Atomic nuclei consist of protons and neutrons bound together by the residual strong force. 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 bind 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 3
2
He
is 1:2, the neutron:proton ratio of 238
92
U
is greater than 3:2. A number of lighter elements have stable nuclides with the ratio 1:1 (Z = N). The nuclide 40
20
Ca
(calcium-40) is observationally the heaviest stable nuclide with the same number of neutrons and protons; (theoretically, the heaviest stable one is sulfur-32). All stable nuclides heavier than calcium-40 contain more neutrons than protons.

Numbers of 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). No element has nine stable isotopes. Xenon is the only element with eight stable isotopes. Four elements have seven stable isotopes, eight have six stable isotopes, ten have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes (counting 180m
73
Ta
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). [26] In total, there are 252 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 252/80 = 3.15 isotopes per element.

Even and odd nucleon numbers

Even/odd Z, N (Hydrogen-1 included as OE)
p,nEEOOEOOETotal
Stable14655348252
Long-lived2243534
All primordial16895653286

The proton:neutron ratio is not the only factor affecting nuclear stability. It depends also on evenness or oddness of its atomic number Z, neutron number N and, consequently, of their sum, the mass number A. Oddness of both Z and N tends to lower the nuclear binding energy, making odd nuclei, generally, less stable. This remarkable difference of nuclear binding energy between neighbouring nuclei, especially of odd-A isobars, has important consequences: unstable isotopes with a nonoptimal number of neutrons or protons decay by beta decay (including positron decay), electron capture or other exotic means, such as spontaneous fission and cluster decay.

The majority of stable nuclides are even-proton-even-neutron, where all numbers Z, N, and A are even. The odd-A stable nuclides are divided (roughly evenly) into odd-proton-even-neutron, and even-proton-odd-neutron nuclides. Odd-proton-odd-neutron nuclei are the least common.

Even atomic number

The 146 even-proton, even-neutron (EE) nuclides comprise ~ 58% of all stable nuclides and all have spin 0 because of pairing. There are also 24 primordial long-lived even-even nuclides. As a result, each of the 41 even-numbered elements from 2 to 82 has at least one stable isotope, and most of these elements have several primordial isotopes. Half of these even-numbered elements have six or more stable isotopes. The extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five ( 5
2
He
, 5
3
Li
) or eight ( 8
4
Be
) nucleons from existing for long enough to serve as platforms for the buildup of heavier elements via nuclear fusion in stars (see triple alpha process).

Even-odd long-lived
Decay Half-life
113
48
Cd
beta 7.7×1015 a
147
62
Sm
alpha 1.06×1011 a
235
92
U
alpha 7.04×108 a

53 stable nuclides have an even number of protons and an odd number of neutrons. They are a minority in comparison to the even-even isotopes, which are about 3 times as numerous. Among the 41 even-Z elements that have a stable nuclide, only two elements (argon and cerium) have no even-odd stable nuclides. One element (tin) has three. There are 24 elements that have one even-odd nuclide and 13 that have two odd-even nuclides. Of 35 primordial radionuclides there exist four even-odd nuclides (see table at right), including the fissile 235
92
U
. Because of their odd neutron numbers, the even-odd nuclides tend to have large neutron capture cross sections, due to the energy that results from neutron-pairing effects. These stable even-proton odd-neutron nuclides tend to be uncommon by abundance in nature, generally because, to form and enter into primordial abundance, they must have escaped capturing neutrons to form yet other stable even-even isotopes, during both the s-process and r-process of neutron capture, during nucleosynthesis in stars. For this reason, only 195
78
Pt
and 9
4
Be
are the most naturally abundant isotopes of their element.

Odd atomic number

Forty-eight stable odd-proton-even-neutron nuclides, stabilized by their paired neutrons, form most of the stable isotopes of the odd-numbered elements; the very few odd-proton-odd-neutron nuclides comprise the others. There are 41 odd-numbered elements with Z = 1 through 81, of which 39 have stable isotopes (the elements technetium (
43
Tc
) and promethium (
61
Pm
) have no stable isotopes). Of these 39 odd Z elements, 30 elements (including hydrogen-1 where 0 neutrons is even) have one stable odd-even isotope, and nine elements: chlorine (
17
Cl
), potassium (
19
K
), copper (
29
Cu
), gallium (
31
Ga
), bromine (
35
Br
), silver (
47
Ag
), antimony (
51
Sb
), iridium (
77
Ir
), and thallium (
81
Tl
), have two odd-even stable isotopes each. This makes a total 30 + 2(9) = 48 stable odd-even isotopes.

There are also five primordial long-lived radioactive odd-even isotopes, 87
37
Rb
, 115
49
In
, 187
75
Re
, 151
63
Eu
, and 209
83
Bi
. The last two were only recently found to decay, with half-lives greater than 1018 years.

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons. The first four "odd-odd" nuclides occur in low mass nuclides, for which changing a proton to a neutron or vice versa would lead to a very lopsided proton-neutron ratio ( 2
1
H
, 6
3
Li
, 10
5
B
, and 14
7
N
; spins 1, 1, 3, 1). The only other entirely "stable" odd-odd nuclide, 180m
73
Ta
(spin 9), is thought to be the rarest of the 252 stable isotopes, and is the only primordial nuclear isomer, which has not yet been observed to decay despite experimental attempts. [27]

Many odd-odd radionuclides (like tantalum-180) with comparatively short half lives are known. Usually, they beta-decay to their nearby even-even isobars that have paired protons and paired neutrons. Of the nine primordial odd-odd nuclides (five stable and four radioactive with long half lives), only 14
7
N
is the most common isotope of a common element. This is the case because it is a part of the CNO cycle. The nuclides 6
3
Li
and 10
5
B
are minority isotopes of elements that are themselves rare compared to other light elements, whereas the other six isotopes make up only a tiny percentage of the natural abundance of their elements.

Odd neutron number

Neutron number parity (1H with 0 neutrons included as even)
NEvenOdd
Stable19458
Long-lived277
All primordial22165

Actinides with odd neutron number are generally fissile (with thermal neutrons), whereas those with even neutron number are generally not, though they are fissionable with fast neutrons. All observationally stable odd-odd nuclides have nonzero integer spin. This is because the single unpaired neutron and unpaired proton have a larger nuclear force attraction to each other if their spins are aligned (producing a total spin of at least 1 unit), instead of anti-aligned. See deuterium for the simplest case of this nuclear behavior.

Only 195
78
Pt
, 9
4
Be
and 14
7
N
have odd neutron number and are the most naturally abundant isotope of their element.

Occurrence in nature

Elements are composed of one nuclide (mononuclidic elements) or of more naturally occurring isotopes. The unstable (radioactive) isotopes are either primordial or postprimordial. Primordial isotopes were a product of stellar nucleosynthesis or another type of nucleosynthesis such as cosmic ray spallation, and have persisted down to the present because their rate of decay is so slow (e.g. uranium-238 and potassium-40). Post-primordial isotopes were created by cosmic ray bombardment as cosmogenic nuclides (e.g., tritium, carbon-14), or by the decay of a radioactive primordial isotope to a radioactive radiogenic nuclide daughter (e.g. uranium to radium). A few isotopes are naturally synthesized as nucleogenic nuclides, by some other natural nuclear reaction, such as when neutrons from natural nuclear fission are absorbed by another atom.

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 (
50
Sn
). There are about 94 elements found naturally on Earth (up to plutonium 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. [28] Only 252 of these naturally occurring nuclides are stable in the sense of never having been observed to decay as of the present time. An additional 34 primordial nuclides (to a total of 286 primordial nuclides), are radioactive with known half-lives, but have half-lives longer than 100 million years, allowing them to exist from the beginning of the Solar System. See list of nuclides for details.

All the known stable nuclides occur naturally on Earth; the other naturally occurring nuclides are radioactive but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. These include the afore-mentioned cosmogenic nuclides, the nucleogenic nuclides, and any radiogenic nuclides formed by ongoing decay of a primordial radioactive nuclide, such as radon and radium from uranium.

An additional ~3000 radioactive nuclides not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived nuclides not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminium-26, which is not naturally found on Earth, but 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% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units.

According to generally accepted cosmology theory, only isotopes of hydrogen and helium, traces of some isotopes of lithium and beryllium, and perhaps some boron, were created at the Big Bang, while all other nuclides were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously produced nuclides. (See nucleosynthesis for details of the various processes thought 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 (mr) of an isotope (nuclide) 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-12 atom. It is denoted with symbols "u" (for unified atomic mass unit) or "Da" (for dalton).

The atomic masses of naturally occurring isotopes of an element determine the atomic mass of the element. When the element contains N isotopes, the expression below is applied for the average atomic mass :

where m1, m2, …, mN are the atomic masses of each individual isotope, and x1, …, xN are the relative abundances of these isotopes.

Applications of isotopes

Purification 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 because 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 spectrometry.

Use of chemical and biological properties

Use of nuclear properties

See also

Notes

Related Research Articles

Stable nuclide nucleus of this isotope does not undergo radioactive decay

Stable nuclides are nuclides that are not radioactive and so do not spontaneously undergo radioactive decay. When such nuclides are referred to in relation to specific elements, they are usually termed stable isotopes.

Nucleosynthesis is the process that creates new atomic nuclei from pre-existing nucleons. The first nuclei were formed a few minutes after the Big Bang, through the process called Big Bang nucleosynthesis. After about 20 minutes, the universe had cooled to a point at which these processes ended, so only the fastest and simplest reactions occurred, leaving our universe containing about 75% hydrogen, 24% helium by mass. The rest is traces of other elements such as lithium and the hydrogen isotope deuterium. The universe still has approximately the same composition.

Decay chain series of elements in radioactive decay

In nuclear science, the decay chain refers to a series of radioactive decays of different radioactive decay products as a sequential series of transformations. It is also known as a "radioactive cascade". Most radioisotopes do not decay directly to a stable state, but rather undergo a series of decays until eventually a stable isotope is reached.

Mass number total number of heavy particles (protons and neutrons jointly called nucleons) in the atomic nucleus

The mass number, also called atomic mass number or nucleon number, is the total number of protons and neutrons in an atomic nucleus. It is approximately equal to the atomic mass of the atom expressed in atomic mass units. 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 each different isotope of a chemical element. Hence, the difference between the mass number and the atomic number Z gives the number of neutrons (N) in a given nucleus: .

Technetium (43Tc) is the first of the two elements lighter than bismuth that have no stable isotopes; the other such element is promethium. It is primarily artificial, with only trace quantities existing in nature produced by spontaneous fission or neutron capture by molybdenum. The first isotopes to be synthesized were 97Tc and 99Tc in 1936, the first artificial element to be produced. The most stable radioisotopes are 97Tc, 98Tc and 99Tc.

Nuclear binding energy energy required to split a nucleus of an atom into its component parts.

Nuclear binding energy is the minimum energy that would be required to disassemble the nucleus of an atom into its component parts. These component parts are neutrons and protons, which are collectively called nucleons. The binding energy is always a positive number, as we need to spend energy in moving these nucleons, attracted to each other by the strong nuclear force, away from each other. The mass of an atomic nucleus is less than the sum of the individual masses of the free constituent protons and neutrons, according to Einstein's equation E=mc2. This 'missing mass' is known as the mass defect, and represents the energy that was released when the nucleus was formed.

Table of nuclides graph of neutrons vs. protons in nuclides

A table of nuclides or chart of nuclides is a two-dimensional graph in which one axis represents the number of neutrons and the other represents the number of protons in an atomic nucleus. Each point plotted on the graph thus represents the nuclide of a real or hypothetical chemical element. This system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements instead of each of their isotopes.

Valley of stability Region of the Chart of the Nuclides with few radioactive isotopes

In nuclear physics, the valley of stability is a characterization of the stability of nuclides to radioactivity based on their binding energy. Nuclides are composed of protons and neutrons. The shape of the valley refers to the profile of binding energy as a function of the numbers of neutrons and protons, with the lowest part of the valley corresponding to the region of most stable nuclei. The line of stable nuclides down the center of the valley of stability is known as the line of beta stability. The sides of the valley correspond to increasing instability to beta decay. The decay of a nuclide becomes more energetically favorable the further it is from the line of beta stability. The boundaries of the valley correspond to the nuclear drip lines, where nuclides become so unstable they emit single protons or single neutrons. Regions of instability within the valley at high atomic number also include radioactive decay by alpha radiation or spontaneous fission. The shape of the valley is roughly an elongated paraboloid corresponding to the nuclide binding energies as a function of neutron and atomic numbers.

Atomic mass mass of an atom in unified atomic mass units.

The atomic mass (ma) is the mass of an atom. Its unit is the unified atomic mass unit where 1 unified atomic mass unit is defined as ​112 of the mass of a single carbon-12 atom, at rest. The protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. Thus, the atomic mass measured in u has nearly the same value as the mass number.

Primordial nuclide nuclides predating the Earths formation (found on Earth)

In geochemistry, geophysics and geonuclear physics, primordial nuclides, also known as primordial isotopes, are nuclides found on Earth that have existed in their current form since before Earth was formed. Primordial nuclides were present in the interstellar medium from which the solar system was formed, and were formed in, or after, the Big Bang, by nucleosynthesis in stars and supernovae followed by mass ejection, by cosmic ray spallation, and potentially from other processes. They are the stable nuclides plus the long-lived fraction of radionuclides surviving in the primordial solar nebula through planet accretion until the present. Only 286 such nuclides are known.

Isobar (nuclide) class of nuclide

Isobars are atoms (nuclides) of different chemical elements that have the same number of nucleons. Correspondingly, isobars differ in atomic number but have the same mass number. An example of a series of isobars would be 40S, 40Cl, 40Ar, 40K, and 40Ca. The nuclei of these nuclides all contain 40 nucleons; however, they contain varying numbers of protons and neutrons.

Monoisotopic element one of 26 chemical elements which have only a single stable isotope (nuclide)

A monoisotopic element is one of 26 chemical elements which have only a single stable isotope (nuclide). A list is given in a following section.

Nuclear transmutation conversion of an atom from one element to another

Nuclear transmutation is the conversion of one chemical element or an isotope into another chemical element. Because any element is defined by its number of protons in its atoms, i.e. in the atomic nucleus, nuclear transmutation occurs in any process where the number of protons or neutrons in the nucleus is changed.

Even and odd atomic nuclei

In nuclear physics, properties of a nucleus depend on evenness or oddness of its atomic number Z, neutron number N and, consequently, of their sum, the mass number A. Most notably, oddness of both Z and N tends to lower the nuclear binding energy, making odd nuclei, generally, less stable. This effect is not only experimentally observed, but is included in the semi-empirical mass formula and explained by some other nuclear models, such as the nuclear shell model. This remarkable difference of nuclear binding energy between neighbouring nuclei, especially of odd-A isobars, has important consequences for beta decay.

References

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  6. IUPAC Gold Book
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  8. This notation seems to have been introduced in the second half of the 1930s. Before that, various notations were used, such as Ne(22) for neon-22 (1934), Ne22 for neon-22 (1935), or even Pb210 for lead-210 (1933).
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    • Strömholm, Daniel and Svedberg, Theodor (1909) "Untersuchungen über die Chemie der radioactiven Grundstoffe II." (Investigations into the chemistry of the radioactive elements, part 2), Zeitschrift für anorganischen Chemie, 63: 197–206; see especially page 206.
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  16. See:
    • Kasimir Fajans (1913) "Über eine Beziehung zwischen der Art einer radioaktiven Umwandlung und dem elektrochemischen Verhalten der betreffenden Radioelemente" (On a relation between the type of radioactive transformation and the electrochemical behavior of the relevant radioactive elements), Physikalische Zeitschrift, 14: 131–136.
    • Soddy announced his "displacement law" in: Soddy, Frederick (1913). "The Radio-Elements and the Periodic Law". Nature. 91 (2264): 57–58. Bibcode:1913Natur..91...57S. doi:10.1038/091057a0..
    • Soddy elaborated his displacement law in: Soddy, Frederick (1913) "Radioactivity," Chemical Society Annual Report, 10: 262–288.
    • Alexander Smith Russell (1888–1972) also published a displacement law: Russell, Alexander S. (1913) "The periodic system and the radio-elements," Chemical News and Journal of Industrial Science, 107: 49–52.
  17. Soddy first used the word "isotope" in: Soddy, Frederick (1913). "Intra-atomic charge". Nature. 92 (2301): 399–400. Bibcode:1913Natur..92..399S. doi:10.1038/092399c0.
  18. Fleck, Alexander (1957). "Frederick Soddy". Biographical Memoirs of Fellows of the Royal Society. 3: 203–216. doi:10.1098/rsbm.1957.0014. p. 208: Up to 1913 we used the phrase 'radio elements chemically non-separable' and at that time the word isotope was suggested in a drawing-room discussion with Dr. Margaret Todd in the home of Soddy's father-in-law, Sir George Beilby.
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  21. In his 1893 book, William T. Preyer also used the word "isotope" to denote similarities among elements. From p. 9 of William T. Preyer, Das genetische System der chemischen Elemente [The genetic system of the chemical elements] (Berlin, Germany: R. Friedländer & Sohn, 1893): "Die ersteren habe ich der Kürze wegen isotope Elemente genannt, weil sie in jedem der sieben Stämmme der gleichen Ort, nämlich dieselbe Stuffe, einnehmen." (For the sake of brevity, I have named the former "isotopic" elements, because they occupy the same place in each of the seven families [i.e., columns of the periodic table], namely the same step [i.e., row of the periodic table].)
  22. 1 2 The origins of the conceptions of isotopes Frederick Soddy, Nobel prize lecture
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