An isotope is one of two or more forms of the same chemical element that have the same number of protons but different numbers of neutrons in their atomic nuclei. Because the number of protons determines the atomic number and therefore the chemical element, all isotopes of an element occupy the same position in the periodic table. Their neutron numbers and consequently their mass numbers differ.
For example, hydrogen has three commonly discussed isotopes: hydrogen 1 contains one proton and no neutron, deuterium contains one proton and one neutron, and tritium contains one proton and two neutrons. Isotopes can be stable or radioactive. Stable isotopes do not undergo spontaneous radioactive decay under ordinary conditions, whereas radioactive isotopes, also called radioisotopes or radionuclides, transform through nuclear decay. Isotopic differences therefore primarily concern nuclear composition and mass, while chemical identity remains determined by proton number. The word isotope derives from Greek roots meaning same place, referring to isotopes occupying the same place in the periodic table.
The defining quantities for an isotope are atomic number and mass number. Atomic number is the number of protons in the nucleus, while mass number is the total number of protons and neutrons. The neutron number can therefore be calculated by subtracting atomic number from mass number. Isotopes are commonly written using the element name or symbol followed by the mass number, such as carbon 12, carbon 13 and carbon 14.
These three carbon isotopes all contain six protons but contain six, seven and eight neutrons respectively. Their chemical behaviour is broadly similar because chemical behaviour is primarily determined by electron configuration, which is governed by nuclear charge. However, isotopic mass differences can produce measurable differences in physical and chemical behaviour. These are called isotope effects and are particularly significant for light elements. Isotopes can also differ dramatically in nuclear stability, because changing the neutron to proton ratio can move a nucleus away from the region of stable nuclear configurations.
Isotopes occur naturally and can also be produced artificially. Natural isotopes arise from stellar nucleosynthesis, radioactive decay, cosmic ray interactions and other nuclear processes, while artificial radionuclides can be produced in nuclear reactors, particle accelerators or other nuclear facilities. Many elements exist naturally as mixtures of several stable isotopes. The relative abundance of these isotopes contributes to the standard atomic weight reported for the element. Some elements have only one naturally occurring stable isotope, while others have many.
Radioactive isotopes can occur naturally as members of decay chains or can be continuously generated by cosmic ray interactions and other natural processes. Human activities have also produced radionuclides that may enter the environment. Isotope analysis can distinguish materials according to their isotopic composition, making isotopes valuable in chemistry, geology, archaeology, environmental science, biology, medicine and nuclear engineering.
The stability of an isotope depends on nuclear structure rather than simply on its mass. The nucleus contains positively charged protons and electrically neutral neutrons. The strong nuclear interaction binds nucleons at short range, while electrostatic repulsion acts between protons. Stable nuclear configurations depend on the balance between these interactions and on quantum mechanical shell structure. Light stable nuclei often have approximately similar numbers of protons and neutrons, while heavier stable nuclei generally require proportionally more neutrons.
If the neutron to proton ratio is unsuitable, the nucleus may undergo beta decay, alpha decay, electron capture, spontaneous fission or another transformation. Radioactive isotopes therefore provide information about nuclear instability and can transform into isotopes of the same or different elements. A beta minus decay converts a neutron into a proton while emitting an electron and an antineutrino, changing the element but preserving the mass number. Alpha decay removes two protons and two neutrons, changing both the element and mass number.
Isotopes have many practical applications. Carbon 14 is used for radiocarbon dating because its production and radioactive decay provide a timescale for studying relatively recent organic material. Stable isotopes such as carbon 13 and nitrogen 15 are used to investigate biological pathways, diets and ecological relationships. Oxygen and hydrogen isotopes are used in hydrology and climate research to study water sources and past environmental conditions. Medical imaging and treatment use radionuclides selected for their radiation properties, half-lives and biological behaviour.
Technetium 99m is widely used in diagnostic nuclear medicine because its decay characteristics are suitable for imaging. Cobalt 60 has been used as a source of gamma radiation for medical treatment and industrial applications. Isotopes are also used as tracers because chemically similar isotopes can follow the same pathways while remaining distinguishable through their mass or radiation.
Isotopes are related to several other nuclear terms that are sometimes confused. Isobars are nuclei with the same mass number but different atomic numbers, whereas isotones are nuclei with the same neutron number but different proton numbers. Isomers are nuclei of the same isotope that occupy different nuclear energy states, with an excited metastable state potentially persisting for a measurable time before transitioning to a lower energy state.
An isotope is therefore defined specifically by its proton number and neutron number, while these other categories use different combinations of nuclear properties. The term nuclide refers more generally to a specific nuclear species characterised by a defined proton number, neutron number and nuclear state. Radioisotope and radionuclide are commonly used for unstable isotopes, although radionuclide can be technically preferable because radioactivity is a property of the nuclear state. These distinctions matter in technical writing because isotope is a broad category that includes both stable and radioactive forms.
Isotopic composition can reveal information that ordinary chemical analysis cannot provide. Because isotopes of the same element participate in many of the same chemical reactions, differences in isotope abundance can preserve records of physical, biological and geological processes. Stable isotope ratios are frequently expressed relative to reference standards and can be used to trace sources and transformations.
Mass spectrometry can separate isotopes according to differences in mass to charge ratio, while radiation detectors can identify radioactive isotopes through their characteristic emissions and energies. Nuclear magnetic resonance and related techniques also exploit nuclear properties that differ among isotopes. In nuclear power engineering, isotopic composition strongly affects reactor behaviour because different isotopes have different neutron absorption and fission probabilities. Enrichment processes therefore change the relative abundance of isotopes rather than changing the chemical identity of the element itself.
The scientific concept of isotopes developed after researchers recognised that atoms of the same element could have different atomic masses. Frederick Soddy introduced the term isotope in 1913 during the development of early nuclear and atomic theory. Mass spectrometry and subsequent nuclear discoveries established that isotopic differences arise from different neutron numbers. The discovery of the neutron in 1932 provided the physical explanation for why nuclei of the same element could have different masses while retaining the same proton number.
Modern nuclear science treats isotopes as specific configurations of nucleons governed by quantum mechanics and nuclear forces. An isotope is therefore not simply a heavier or lighter version of an atom in an everyday sense. Its mass, nuclear stability, magnetic properties, decay behaviour and reaction probabilities can differ significantly from those of other isotopes of the same element. Understanding isotopes is fundamental to nuclear physics because nearly every nuclear process depends on the specific nuclide involved rather than merely on the chemical element name.
