Isotopen

Isotopen, or isotopes in English, are variants of a chemical element that share the same number of protons but have different numbers of neutrons. This means they have the same atomic number Z and occupy the same place in the periodic table, but different mass numbers A. Because they have identical electron configurations, isotopes of a given element typically show very similar chemical behavior, while their nuclear properties and physical characteristics such as mass, density, and stability differ.

Isotopes are usually denoted by a superscript mass number before the element symbol, e.g., 12C, 13C, 14C. Common notation also uses the atomic number as a subscript. The concept includes both stable isotopes and radioactive isotopes (radioisotopes). Many elements have several stable isotopes; some have only unstable ones. Hydrogen has three common isotopes (1H, 2H, 3H); carbon has 12C, 13C, 14C.

The concept was introduced in the early 20th century by Frederick Soddy and others, after researchers observed that elements could occur with different atomic masses while behaving chemically as the same element. Isotopes are identified by their mass, decay properties, and, increasingly, by mass spectrometry, which can separate nuclei with slightly different masses.

Applications of isotopes include dating, tracing, and medicine: radiocarbon dating with 14C; medical imaging with technetium-99m; and therapy with iodine-131. Natural isotopic abundances are used in climate and ecological research, while isotopes produced in reactors or accelerators expand the materials available for science and industry.