Stable isotopes are chemical isotopes that are not radioactive (they have not been observed to decay, though a few of them may be theoretically unstable with exceedingly long half-lives). By this definition, there are 256 known stable isotopes of the 80 elements which have one or more stable isotopes. A list of these is given at the end of this article. About two thirds of the elements have more than one stable isotope. One element (tin) has ten stable isotopes.
Different isotopes of the same element (whether stable or unstable) have nearly the same chemical characteristics and therefore behave almost identically in biology (a notable exception is the isotopes of hydrogen - see heavy water). The mass differences, due to a difference in the number of neutrons, will result in partial separation of the light isotopes from the heavy isotopes during chemical reactions and during physical processes such as diffusion and vaporization. This process is called (isotope fractionation). For example, the difference in mass between the two stable isotopes of hydrogen, 1H (1 proton, no neutron, also known as protium) and 2H (1 proton, 1 neutron, also known as deuterium) is almost 100%. Therefore, a significant fractionation will occur.
Commonly analysed stable isotopes include oxygen, carbon, nitrogen, hydrogen and sulfur. These isotope systems have been under investigation for many years in order to study processes of isotope fractionation in natural systems because they are relatively simple to measure. Recent advances in mass spectrometry (i.e. multiple-collector inductively coupled plasma mass spectrometry) now enable the measurement of heavier stable isotopes, such as iron, copper, zinc, molybdenum, etc.
Stable isotopes have been used in botanical and plant biological investigations for many years, and more and more ecological and biological studies are finding stable isotopes (mostly carbon, nitrogen and oxygen) to be extremely useful. Other workers have used oxygen isotopes to reconstruct historical atmospheric temperatures, making them important tools for climate research.
Most naturally occurring isotopes are stable; however, a few tens of them are radioactive with very long half-lives. If the half-life of a nuclide is comparable to or greater than the Earth's age (4.5 billion years), a significant amount will have survived since the formation of the Solar System (it will be primordial), and will contribute in that way to the natural isotopic composition of a chemical element. The shortest half-lives of easily detectable primordially present radioisotopes are around 700 million years (e.g., 235U), with a lower present limit on detection of primordial isotopes of 80 million years (e.g., 244Pu). Many radioisotopes are known in nature with still shorter half-lives, but they are made freshly by decay processes or ongoing energetic reactions, such as those produced by present bombardment of Earth by cosmic rays.
Many isotopes that are presumed to be stable (i.e. no radioactivity has been observed for them) are predicted to be radioactive with extremely long half-lives (sometimes as high as 1018 years or more). If the predicted half-life falls into an experimentally accessible range, such isotopes have a chance to move from the list of stable nuclides to the radioactive category, once their activity is observed. Good examples are bismuth-209 and tungsten-180 which were formerly classed as stable, but have been recently (2003) found to be alpha-active.
Most stable isotopes in the earth are believed to have been formed in processes of nucleosynthesis, either in the 'Big Bang', or in generations of stars that preceded the formation of the solar system. However, some stable isotopes also show abundance variations in the earth as a result of decay from long-lived radioactive nuclides. These decay-products are termed radiogenic isotopes, in order to distinguish them from the much larger group of 'non-radiogenic' isotopes. They play an important role in radiometric dating and isotope geochemistry.