Isotopes are variants of a particular chemical element which differ in
neutron number. All isotopes of a given element have the same number
of protons in each atom. 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
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.
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
Isotope vs. nuclide
2 Definition of isotopes
4 Radioactive, primordial, and stable isotopes
5.2 Stable isotopes
6 Variation in properties between isotopes
6.1 Chemical and molecular properties
6.2 Nuclear properties and stability
6.3 Numbers of isotopes per element
6.4 Even and odd nucleon numbers
6.4.1 Even atomic number
6.4.2 Odd atomic number
6.4.3 Odd neutron number
7 Occurrence in nature
Atomic mass of isotopes
9 Applications of isotopes
9.1 Purification of isotopes
9.2 Use of chemical and biological properties
9.3 Use of nuclear properties
10 See also
13 External links
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, now sometimes
isotopic nuclides) is intended to imply comparison (like synonyms
or isomers), for example: the nuclides 12
6C are isotopes (nuclides with the same atomic number but different
mass numbers), but 40
20Ca are isobars (nuclides with the same mass number). 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.
Definition of isotopes
Isotopes are atoms of the same element having the same numbers of
protons (atomic number), but different numbers of neutrons. They have
same chemical properties due to the same electronic configuration but
different physical properties.
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). 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
92U, and 239
92U). 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. 3He, 4He, 12C, 14C, 235U, and
239U). 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
The common pronunciation of the AZE notation is different from how it
is written: 4
2He is commonly pronounced as helium-four instead of four-two-helium,
92U 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, 14C is a radioactive form of carbon,
whereas 12C and 13C are stable isotopes. There are about 339 naturally
occurring nuclides on Earth, of which 286 are primordial nuclides,
meaning that they have existed since the Solar System's formation.
Primordial nuclides include 32 nuclides with very long half-lives
(over 100 million years) and 253 that are formally considered as
"stable nuclides", 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
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). Of the 253 nuclides never observed to decay, only 90 of
these (all from the first 40 elements) are theoretically stable to all
known forms of decay. Element 41 (niobium) is theoretically unstable
via spontaneous fission, but this has never been detected. Many other
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 27 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. These
include 905 nuclides that are either stable or have half-lives longer
than 60 minutes. See list of nuclides for details.
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
Several attempts to separate these new radioelements chemically had
failed. 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. 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. 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
Soddy proposed that several types of atoms (differing in radioactive
properties) could occupy the same place in the table. For example,
the alpha-decay of uranium-235 forms thorium-231, whereas the beta
decay of actinium-230 forms thorium-230. The term "isotope", Greek
for "at the same place", was suggested to Soddy by Margaret Todd,
a Scottish physician and family friend, during a conversation in which
he explained his ideas to her. He won the 1921
Nobel Prize in
Chemistry in part for his work on isotopes.
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
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.
The first evidence for multiple isotopes of a stable (non-radioactive)
element was found by
J. J. Thomson
J. J. Thomson in 1913 as part of his exploration
into the composition of canal rays (positive ions). 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
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
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 (1H), deuterium (2H), and tritium (3H), 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
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
See also: Stable nuclide,
Stable isotope ratio, List of nuclides, and
List of elements by stability of isotopes
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
2He is 1:2, the neutron:proton ratio of 238
92U is greater than 3:2. A number of lighter elements have stable
nuclides with the ratio 1:1 (Z = N). The nuclide 40
20Ca (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
73Ta 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). In total, there are 253 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 253/80
= 3.1625 isotopes per element.
Even and odd nucleon numbers
Main article: Even and odd atomic nuclei
Even/odd Z, N (
Hydrogen-1 included as OE)
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 148 even-proton, even-neutron (EE) nuclides comprise ~ 58% of all
stable nuclides and all have spin 0 because of pairing. There are also
22 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 or eight nucleons
from existing for long enough to serve as platforms for the buildup of
heavier elements via nuclear fusion in stars (see triple alpha
These 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
92U. 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
78Pt and 9
4Be are the most naturally abundant isotopes of their element.
Odd atomic number
Forty-eight stable odd-proton-even-neutron nuclides, stabilized by
their even numbers of 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 (
43Tc) and promethium (
61Pm) 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 (
17Cl), potassium (
19K), copper (
29Cu), gallium (
31Ga), bromine (
35Br), silver (
47Ag), antimony (
51Sb), iridium (
77Ir), and thallium (
81Tl), 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
63Eu, and 209
83Bi. 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
5B, and 14
7N; spins 1, 1, 3, 1). The only other entirely "stable" odd-odd
nuclide is 180m
73Ta (spin 9) is thought to be the rarest of the 253 stable isotopes,
and is the only primordial nuclear isomer, which has not yet been
observed to decay despite experimental attempts.
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
7N 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
3Li and 10
5B 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)
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.
4Be and 14
7N have odd neutron number and are the most naturally abundant isotope
of their element.
Occurrence in nature
See also: Abundance of the chemical elements
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 (
50Sn). 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. Only 253 of these naturally
occurring nuclides are stable in the sense of never having been
observed to decay as of the present time. An additional 35 primordial
nuclides (to a total of 289 primordial nuclides), are radioactive with
known half-lives, but have half-lives longer than 80 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
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
displaystyle overline m _ a
displaystyle overline m _ a =m_ 1 x_ 1 +m_ 2 x_ 2 +...+m_ N x_
where m1, m2, …, mN are the atomic masses of each individual
isotope, and x1, …, xN are the relative abundances of these
Applications of isotopes
Purification of isotopes
Main article: isotope separation
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
Main articles: isotope geochemistry, cosmochemistry, and
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 in
food products or the geographic origins of products using
isoscapes. The identification of certain meteorites as having
Mars is based in part upon the isotopic signature of
trace gases contained in them.
Isotopic substitution can be used to determine the mechanism of a
chemical reaction via the kinetic isotope effect.
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, even different nonradioactive
stable isotopes 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).
Isotopes are commonly used to determine the concentration of various
elements or substances using the isotope dilution method, whereby
known amounts of isotopically-substituted compounds are mixed with the
samples and the isotopic signatures of the resulting mixtures are
determined with mass spectrometry.
Use of nuclear properties
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 concentration of
isotope existed. The most widely known example is radiocarbon dating
used to determine the age of carbonaceous materials.
Several forms of spectroscopy rely on the unique nuclear properties of
specific isotopes, both radioactive and stable. For example, nuclear
magnetic resonance (NMR) spectroscopy can be used only for isotopes
with a nonzero nuclear spin. The most common nuclides 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
Nuclear medicine and radiation oncology utilize
radioisotopes respectively for medical diagnosis and treatment.
Abundance of the chemical elements
List of elements by stability of isotopes
List of nuclides
List of particles
Nuclear medicine (includes medical isotopes)
Radionuclide (or radioisotope)
Table of nuclides
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
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. Isomers are by definition both isotopic and isobaric. (Not
to be confused with chemical isomers.)
Isodiaphers are nuclides having the same neutron excess, i.e. number
of neutrons minus number of protons.
Bainbridge mass spectrometer
^ Soddy, Frederick (1913). "Intra-atomic charge". Nature. 92 (2301):
399–400. Bibcode:1913Natur..92..399S. doi:10.1038/092399c0.
^ IUPAP Red Book
IUPAC Gold Book
IUPAC Gold Book
IUPAC (Connelly, N. G.; Damhus, T.; Hartshorn, R. M.; and Hutton, A.
T.), Nomenclature of Inorganic
2005, The Royal Society of Chemistry, 2005;
IUPAC (McCleverty, J. A.;
and Connelly, N. G.), Nomenclature of Inorganic
Recommendations 2000, The Royal Society of Chemistry, 2001; IUPAC
(Leigh, G. J.), Nomenclature of Inorganic
1990), Blackwell Science, 1990; IUPAC, Nomenclature of Inorganic
Chemistry, Second Edition, 1970; probably in the 1958 first edition as
^ 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
^ a b "Radioactives Missing From The Earth".
^ "NuDat 2 Description". Retrieved 2 January 2016.
^ Choppin, G.; Liljenzin, J. O. and Rydberg, J. （1995）
Radiochemistry and Nuclear
Chemistry (2nd ed.) Butterworth-Heinemann,
^ Others had also suggested the possibility of isotopes; for example:
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.
Alexander Thomas Cameron,
Radiochemistry (London, England: J. M. Dent
& Sons, 1910), p. 141. (Cameron also anticipated the displacement
^ a b c Ley, Willy (October 1966). "The Delayed Discovery". For Your
Information. Galaxy Science Fiction. pp. 116–127.
^ a b c Scerri, Eric R. (2007) The Periodic Table Oxford University
Press, pp. 176–179 ISBN 0-19-530573-6
^ a b Nagel, Miriam C. (1982). "Frederick Soddy: From Alchemy to
Isotopes". Journal of Chemical Education. 59 (9): 739–740.
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:
Soddy announced his "displacement law" in: Soddy, Frederick (1913).
"The Radio-Elements and the Periodic Law". Nature. 91 (2264): 57.
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:
^ Soddy first used the word "isotope" in: Soddy, Frederick (1913).
"Intra-atomic charge". Nature. 92 (2301): 399–400.
^ 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.
^ Budzikiewicz H, Grigsby RD (2006). "
Mass spectrometry and isotopes:
a century of research and discussion".
Mass spectrometry reviews. 25
(1): 146–57. Bibcode:2006MSRv...25..146B. doi:10.1002/mas.20061.
^ Scerri, Eric R. (2007) The Periodic Table, Oxford University Press,
ISBN 0-19-530573-6, Ch. 6, note 44 (p. 312) citing Alexander
Fleck, described as a former student of Soddy's.
^ 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].)
^ a b The origins of the conceptions of isotopes Frederick Soddy,
Nobel prize lecture
^ Thomson, J. J. (1912). "XIX. Further experiments on positive rays".
Philosophical Magazine. Series 6. 24 (140): 209.
^ Thomson, J. J. (1910). "LXXXIII. Rays of positive electricity".
Philosophical Magazine. Series 6. 20 (118): 752.
^ Mass spectra and isotopes Francis W. Aston, Nobel prize lecture 1922
^ Sonzogni, Alejandro (2008). "Interactive Chart of Nuclides".
National Nuclear Data Center: Brookhaven National Laboratory.
^ Hult, Mikael; Wieslander, J. S.; Marissens, Gerd; Gasparro, Joël;
Wätjen, Uwe; Misiaszek, Marcin (2009). "Search for the radioactivity
of 180mTa using an underground HPGe sandwich spectrometer". Applied
Radiation and Isotopes. 67 (5): 918–21.
doi:10.1016/j.apradiso.2009.01.057. PMID 19246206.
^ "Radioactives Missing From The Earth". Don-lindsay-archive.org.
^ Jamin, Eric; Guérin, Régis; Rétif, Mélinda; Lees, Michèle;
Martin, Gérard J. (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 (18): 5202. doi:10.1021/jf030167m.
^ Treiman, A. H.; Gleason, J. D.; Bogard, D. D. (2000). "The SNC
meteorites are from Mars".
Planet. Space Sci. 48 (12–14): 1213.
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