In physics, natural abundance (NA) refers to the abundance of isotopes of a chemical element as naturally found on a planet. The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet, and even from place to place on the Earth, but remains relatively constant in time (on a short-term scale).
As an example, uranium has three naturally occurring isotopes: U, U, and U. Their respective natural mole-fraction abundances are 99.2739–99.2752%, 0.7198–0.7202%, and 0.0050–0.0059%. For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,274 U atoms, approximately 720 U atoms, and very few (most likely 5 or 6) U atoms. This is because U is much more stable than U or U, as the half-life of each isotope reveals: 4.468 × 10 years for U compared with 7.038 × 10 years for U and 245,500 years for U.
Exactly because the different uranium isotopes have different half-lives, when the Earth was younger, the isotopic composition of uranium was different. As an example, 1.7×10 years ago the NA of U was 3.1% compared with today's 0.7%, and that allowed a natural nuclear fission reactor to form, something that cannot happen today.
However, the natural abundance of a given isotope is also affected by the probability of its creation in nucleosynthesis (as in the case of samarium; radioactive Sm and Sm are much more abundant than stable Sm) and by production of a given isotope as a daughter of natural radioactive isotopes (as in the case of radiogenic isotopes of lead).