Isotopes

Isotopes are atoms of the same element that have identical atomic numbers (the number of protons) but different numbers of neutrons, giving them different mass numbers. Because chemical behavior is determined largely by electron structure and proton count, isotopes of a given element behave very similarly in reactions, but their physical and nuclear properties can differ markedly.

Stable isotopes do not decay over time; radioactive isotopes (radioisotopes) are unstable and decay at characteristic rates, described by half-lives. The fraction of a given isotope in nature is called its isotopic abundance; many elements have several stable isotopes, while others have only one stable isotope or none.

Isotopes are usually written by a mass number attached to the element symbol, for example carbon-12 (12C) and carbon-14 (14C).

Natural sources produce isotopes through cosmic-ray spallation, neutron capture, and radioactive decay chains; additional isotopes are synthesized in nuclear reactors or particle accelerators for research, medicine, industry, and archaeology.

Applications include radiometric dating (carbon-14 dating, uranium-lead dating), tracing chemical and biological processes, medical imaging and therapy (such as technetium-99m and fluorine-18), and industrial uses e.g. in dating artifacts, studying environmental processes, or checking material age.

Isotopes provide insights into nuclear physics, stellar nucleosynthesis, and Earth history, while their practical use depends on stability, half-life, and production feasibility.