Isotope

An isotope is one of two or more forms of the same chemical element that contain the same number of protons but a different number of neutrons. Because the proton count defines the element’s identity, all isotopes of an element share the same chemical properties, but their nuclear properties and masses differ. The total number of nucleons is the mass number A, which equals the sum of protons Z and neutrons N. Some isotopes are stable, while others are radioactive and decay over time into different elements or isotopes. Natural elements usually include several isotopes in characteristic proportions, and additional isotopes can be produced by cosmic processes or human activities.

Isotopes are commonly written by giving the element name followed by the mass number, for example carbon-12 or carbon-14. In notation using symbols, one writes the element symbol with a superscript mass number, e.g., 12C or 14C. Although isotopes behave chemically very similarly, their differing masses can produce subtle physical differences and influence reaction rates in a phenomenon known as the isotope effect. Mass spectrometry and other analytical techniques distinguish isotopes by their mass-to-charge ratios.

Examples include hydrogen, which has protium (1H), deuterium (2H or D), and tritium (3H or T); carbon-12, carbon-13, and carbon-14; oxygen-16, oxygen-17, and oxygen-18; and uranium-235 and uranium-238. Natural abundances vary by element, and many isotopes are either radioactive or artificially produced for research, medical imaging and therapy (such as technetium-99m or fluorine-18 in PET), industrial tracing, or as fuel and materials in nuclear reactors. The concept of isotopes was established in the early 20th century by Frederick Soddy.