A nuclide is any particular atomic nucleus with a specific number of a atom 'Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons. Collectively, all the isotopes of all the elements form the set of nuclides. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. Isotope is better used when referring to several different nuclides of the same element; nuclide'' is more generic and is used when referencing only one nucleus or several nuclei of different elements. For example, it is more correct to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope.
In IUPAC nomenclature, isotopes and nuclides are specified by the name of the particular element, implicitly giving the atomic number, followed by a hyphen and the mass number (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g. 3He, 12C, 13C, 131I and 238U).
About 339 nuclides occur naturally on Earth, of which 250 (about 74%) are stable. Counting the radioactive nuclides not found in nature that have been created artificially, more than 3100 nuclides are currently known.
The term isotope was coined in 1919 by Margaret Todd, a Scottish doctor, during a conversation with Frederick Soddy (to whom she was distantly related by marriage). Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if several elements occupied each position in the periodic table. Hence Todd suggested the Greek term for "at the same place" as a suitable name. Soddy adopted the term and went on to win the Nobel Prize for Chemistry in 1921 for his work on radioactive substances.
Soddy's use of the word isotope was initially with regard to radioactive (unstable) atoms. However, in 1913, as part of his exploration into the composition of canal rays, J. J. Thomson channeled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. This was the first observation of different stable isotopes for an element. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.
Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.
Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). Xenon is the only element which has nine stable isotopes. There is no element with exactly eight stable isotopes. See list of elements by nuclear stability for a complete list. Five elements have seven stable isotopes, three have six stable isotopes, nine elements have five stable isotopes, eight have four stable isotopes, 11 have three stable isotopes, 15 have two stable isotopes, and 27 mononuclidic elements have only a single stable isotope. For the 80 elements which have stable isotopes, the average number of stable isotopes is 250/80 = 3.1.
Other effects besides the bulk ratio of protons and neutrons affect nuclear stability. For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five nucleons from existing for long enough to serve as platforms for building up of heavier elements during fusion formation in stars (see triple alpha process). A similar pairing pattern shows in the fact that the 250 known stable nuclides contain only four that have both an odd number of protons and an odd number of neutrons: 2H, 6Li, 10B, 14N. Also, a few long-lived radioactive odd-odd nuclides (40K, 50V, 138La, 180mTa) occur naturally. Most odd-odd nuclides are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.
Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture cross-sections and gamma spectroscopy and nuclear magnetic resonance properties.
If too many or too few neutrons are present with regard to the optimum ratio, the nucleus becomes unstable and subject to certain types of nuclear decay. Unstable isotopes with a non-optimal number of neutrons decay by alpha decay, beta decay, or other exotic means, such as spontaneous fission and cluster decay.
Elements are composed of one or more naturally occurring isotopes, which are normally stable. Some elements have unstable (radioactive) isotopes, either because their decay is so slow that a fraction still remains since they were created (examples: uranium, potassium), or because they are continually created through cosmic radiation (tritium, carbon-14) or by decay from an isotope in the first category (radium, radon).
Only 80 elements have stable isotopes, and 16 of these have only one stable isotope. Most elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (element number 50). There are about 94 elements found naturally on Earth (up to plutonium, element 94, inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists estimates that the elements which occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 250 of these naturally-occurring isotopes are stable (all known stable isotopes occur naturally on Earth); the other 89 naturally-occurring isotopes are radioactive, but occur on Earth due to their relatively long half-lives, or from other means of natural production. An additional ~ 2700 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.
The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. A good example is chlorine, having the composition 35Cl, 75.8%, and 37Cl, 24.2%, giving an atomic mass of 35.5. Values like this confounded scientists before the discovery of isotopes, as most light element atomic masses are close to integer multiples of hydrogen.
According to generally accepted cosmology only isotopes of hydrogen and helium, and traces of some isotopes of lithium, beryllium and boron were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously-produced isotopes. The most common isotope of hydrogen has no neutrons at all. (See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.
The molecular mass (Mr) of an element is determined by its nucleons. For example, Carbon-12 (12C) has 6 Protons and 6 Neutrons. When a sample contains two isotopes the equation below is applied where Mr(1) and Mr(2) are the molecular masses of each individual isotope, and % abundance is the percentage abundance of that isotope in the sample.
Several applications exist that capitalize on properties of the various isotopes of a given element.