Isotopes are forms of a chemical element whose nuclei have the same atomic number, Z, but different atomic masses, A. The word isotope, meaning at the same place, comes from the fact that all isotopes of an element are located at the same place on the periodic table.

The atomic number corresponds to the number of protons in an atom. Thus, isotopes of a particular element contain the same number of protons. The difference in atomic masses results from differences in the number of neutrons in the atomic nuclei.

Collectively, the isotopes of the elements form the set of nuclides. A nuclide is a particular type of atomic nucleus, or more generally an agglomeration of protons and neutrons. Strictly speaking, 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 scientific nomenclature, isotopes (nuclides) are specified by the name of the particular element by a hyphen and the number of nucleons (protons and neutrons) in the atomic nucleus (e.g., helium-3, carbon-12, carbon-14, iron-57, uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g., 3He, 12C, 14C, 57Fe, 238U).


Variation in properties between isotopes

In a neutral atom, the number of electrons equals the number of protons. Thus, different isotopes of a given element also have the same number of electrons and the same electronic structure. Because the chemical behavior of an atom is largely determined by its electronic structure, isotopes exhibit nearly identical chemical behavior. The primary exception is that, due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. (This phenomenon is termed the kinetic isotope effect).

This "mass effect" is most pronounced for protium (1H) vis--vis deuterium (2H), because deuterium has twice the mass of protium. For heavier elements the relative mass difference between isotopes is much less, and the mass effect is usually negligible.

Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have nearly identical electronic structure, and therefore have similar physical and chemical properties. 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.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Atomic nuclei consist of protons and neutrons bound together by the strong nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion and stabilizing the nucleus. For this reason neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, additional neutrons are needed to form a stable nucleus; for example, although the neutron/proton ratio of 3He is 1/2, the neutron/proton ratio of 238U is greater than 3/2. If too many neutrons or too few neutrons are present, the nucleus becomes unstable and subject to nuclear decay.

Occurrence in nature

Several isotopes of each element can be found in nature. The relative abundance of an isotope is strongly correlated with its tendency toward nuclear decay; short-lived nuclides quickly decay away, while their long-lived counterparts endure. However, this does not mean that short-lived species disappear entirely; many are continually produced through the decay of longer-lived nuclides. The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses.

According to generally accepted cosmology, virtually all nuclides other than isotopes of hydrogen and helium were built in stars and supernovae. Their respective abundances here result from the quantities formed by these processes, their spread through the galaxy, and their rates of decay. After the initial coalescence of the solar system, isotopes were redistributed according to mass (see also Origin of the solar system). The isotopic composition of elements is different on different planets, making it possible to determine the origin of meteorites.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element.

Use of chemical properties

  • One of the most common applications is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy (see "Properties"). If radioactive isotopes are used, they can be detected by the radiation they emit (this is radioisotopic labelling).
  • A technique similar to radioisotopic labelling is radiometric dating (most famously radiocarbon dating). It can be used to study chemical processes that the experimenter does not witness, by using naturally-occurring isotopic tracers.

Use of nuclear properties

  • Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are 1H, 2D, 13C, and 31P.

See also

External links

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