isotopers

Isotopes are variants of a chemical element that share the same atomic number (number of protons) but differ in the number of neutrons. This difference in neutron count gives isotopes different atomic masses and can affect nuclear stability while leaving most chemical properties largely unchanged.

Some isotopes are stable, remaining unchanged for long periods, while others are radioactive and decay, emitting radiation as they transform into different elements. Decay modes include alpha, beta, and gamma emissions, and the rate is described by a half-life. The presence of both stable and unstable isotopes explains a wide range of natural and artificial phenomena.

Natural occurrence and production: Many elements have several naturally occurring isotopes with varying abundances. Others have only a single stable isotope or none at all in nature. Isotopes are produced in nuclear reactors, particle accelerators, or through stellar processes, and can be separated or measured using techniques such as mass spectrometry.

Applications and examples: Isotopes are used in radiometric dating (for example, carbon-14 dating), medical diagnostics and therapy (such as fluorine-18 in PET imaging), archaeology, hydrology, and environmental tracing. Hydrogen has three light isotopes: protium, deuterium, and tritium, each with distinct nuclear properties and applications.

History and terminology: The concept emerged in the early 20th century through the work of scientists including Frederick Soddy, who coined the term isotope to describe variants of an element with identical chemical behavior but different masses. The term isotoper is not a standard scientific term; if encountered, it is likely a misspelling or a fictional usage, and should be treated as synonymous with isotope only if explicitly defined in a given context.