• Even for the lightest elements, whose ratio of neutron number to atomic number varies the most between isotopes, it usually has only a small effect although it matters in
  some circumstances (for hydrogen, the lightest element, the isotope effect is large enough to affect biology strongly).

• 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.

• 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.

• Thus different isotopes of a given element all have the same number of electrons and share a similar electronic structure.

• 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.

• [17][18][19][20] 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.

• Nuclear properties and stability[edit] 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.

• 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.

• 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 ).

• They have the same atomic number (number of protons in their nuclei) and position in the periodic table (and hence belong to the same chemical element), but differ in nucleon
  numbers (mass numbers) due to different numbers of neutrons in their nuclei.

• Numbers of isotopes per element[edit] Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin).

• 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.

• 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.

• While all isotopes of a given element have almost the same chemical properties, they have different atomic masses and physical properties.

• The extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five (5 2He , 5 3Li ) or eight (8 4Be ) nucleons from
  existing long enough to serve as platforms for the buildup of heavier elements via nuclear fusion in stars (see triple alpha process).

• [15][26] Stable isotopes[edit] The first evidence for multiple isotopes of a stable (non-radioactive) element was found by J. J. Thomson in 1912 as part of his exploration
  into the composition of canal rays (positive ions).

• 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.

• 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.

• [16] 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.

• Similarly, two molecules that differ only in the isotopes of their atoms (isotopologues) have identical electronic structures, and therefore almost indistinguishable physical
  and chemical properties (again with deuterium and tritium being the primary exceptions).

• 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, for heavier elements, the relative mass difference between isotopes is much less so that the mass-difference effects on chemistry are usually negligible.

• 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 emission), electron capture, or other less common decay modes such as spontaneous fission and cluster decay.

• Occurrence in nature Elements are composed either of one nuclide (mononuclidic elements), or of more than one naturally occurring isotopes.

• • Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes, both radioactive and stable.

• Even and odd nucleon numbers[edit] Main article: Even and odd atomic nuclei The proton:neutron ratio is not the only factor affecting nuclear stability.

• Because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior.

• 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.

• Five 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).

• Use of nuclear properties[edit] • 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 atomic masses of naturally occurring isotopes of an element determine the standard atomic weight of the element.

• 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.

• History Radioactive isotopes[edit] 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.

• Among the 41 even-Z elements that have a stable nuclide, only two elements (argon and cerium) have no even-odd stable nuclides.

• (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.)

• 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).

• 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.

• 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.

• Odd atomic number[edit] 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.

• 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 separation of hydrogen and deuterium is unusual because it is based on chemical rather than physical properties, for example in the Girdler sulfide process.

• The only other entirely “stable” odd-odd nuclide, 180m 73Ta (spin 9), is thought to be the rarest of the 251 stable nuclides, and is the only primordial nuclear isomer, which
  has not yet been observed to decay despite experimental attempts.

• 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.

• However, isotope is the older term and so is better known than nuclide and is still sometimes used in contexts in which nuclide might be more appropriate, such as nuclear
  technology and nuclear medicine.

• For example, carbon-12, carbon-13, and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13, and 14, respectively.

• 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.

• However, by using isotopes of different masses, even different nonradioactive stable isotopes can be distinguished by mass spectrometry or infrared spectroscopy.

• 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.

• For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 251/80 ≈ 3.14 isotopes per element.

• Applications of isotopes Purification of isotopes[edit] Main article: isotope separation Several applications exist that capitalize on the properties of the various isotopes
  of a given element.

• Odd neutron number[edit] 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.

• Even atomic number[edit] The 146 even-proton, even-neutron (EE) nuclides comprise ~58% of all stable nuclides and all have spin 0 because of pairing.

• Half of these even-numbered elements have six or more stable isotopes.

• 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.

• Many short-lived nuclides not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae.

• • 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.

• Thomson observed two separate parabolic patches of light on the photographic plate (see image), which suggested two species of nuclei with different mass to charge ratios.

• [3] The number of protons within the atom’s nucleus is called its atomic number and is equal to the number of electrons in the neutral (non-ionized) atom.

• Most stable nuclides are even-proton-even-neutron, where all numbers Z, N, and A are even.

 

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o ^ Soddy, Frederick (12 December 1922). “The origins of the conceptions of isotopes” (PDF). Nobelprize.org. p. 393. Retrieved 9 January 2019. Thus the chemically identical
elements – or isotopes, as I called them for the first time in this letter to Nature, because they occupy the same place in the Periodic Table …
o ^ “isotope—Origin and meaning”. www.etymonline.com. Retrieved 21 October 2021.
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(1913). “Intra-atomic charge”. Nature. 92 (2301): 399–400. Bibcode:1913Natur..92..399S. doi:10.1038/092399c0. S2CID 3965303.
o ^ “IUPAP Red Book” (PDF). Archived from the original (PDF) on 2015-03-18. Retrieved 2018-01-06.
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o ^ 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).
o ^ Jump up to:a b “Radioactives Missing From The Earth”.[dead link]
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o ^ Choppin, G.; Liljenzin, J. O. and Rydberg,
J.　（1995） Radiochemistry and Nuclear Chemistry (2nd ed.) Butterworth-Heinemann, pp. 3–5
o ^ 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 law.)
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and the electrochemical behavior of the relevant radioactive elements), Physikalische Zeitschrift, 14: 131–136.
o ^ 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. S2CID 3975657..
o ^ Soddy elaborated his displacement law in: Soddy, Frederick (1913) “Radioactivity,” Chemical Society Annual Report, 10: 262–288.
o ^ 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.
o ^ 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. S2CID 3965303.
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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.
o ^ Budzikiewicz
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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.
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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
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Photo credit: https://www.flickr.com/photos/ell-r-brown/5970300143/’]