Isotope
One of two or more alternative forms of an element that have the same number of protons in their nucleus, but have different numbers of neutrons.
One member of a (chemical-element) family of atomic species which has two or more nuclides with the same number of protons (Z) but a different number of neutrons (N). Because the atomic mass is determined by the sum of the number of protons and neutrons contained in the nucleus, isotopes differ in mass. Since they contain the same number of protons (and hence electrons), isotopes have the same chemical properties. However, the nuclear and atomic properties of isotopes can be different. The electronic energy levels of an atom depend upon the nuclear mass. Thus, corresponding atomic levels of isotopes are slightly shifted relative to each other. A nucleus can have a magnetic moment which can interact with the magnetic field generated by the electrons and lead to a splitting of the electronic levels. The number of resulting states of nearly the same energy depends upon the spin of the nucleus and the characteristics of the specific electronic level.
Example: hydrogen has three isotopes, of atomic masses 1, 2, and 3, generally written as 1H, 2H (deuterium), and 3H (tritium). 1H is the most abundant isotope of hydrogen; 2H is stable, while 3H is radioactive. Radioactive isotopes are unstable, and decay to stable elements, emitting radiation in the process. This may be ?-radiation, ?-radiation (electrons), ?-radiation, or X-rays, depending on the isotope. The time taken for half the radioactivity to decay is the half-life of the isotope, and can vary from a fraction of a second, through several days to years (e.g. the half-life of 3H is 12½ years, that of 14C is 5200 years).
Stable isotopes can be detected only by their different atomic mass. Since they emit no radiation, they are safe for use in labelled compounds given to human beings. Examples of stable isotopes commonly used in nutrition research include 2H, 13C, 15N, and 18O.
