Everything about Relative Abundance totally explained
In
chemistry,
natural abundance (NA) refers to the prevalence of
isotopes of a
chemical element as naturally found on a planet. The relative atomic mass (a weighted average) 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 but remains relatively constant in time.
As an example,
uranium has three naturally occurring isotopes: U-238, U-235 and U-234. Their respective NA is 99.2745%, 0.72% and 0.0055%. For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,275 U-238 atoms, 720 U-235 atoms, and no more than 5 or 6 U-234 atoms. This is because U-238 is much more stable than U-235 or U-234, as the
half-life of each isotope reveals: 4.468×10
9 years for U-238 compared to 7.038×10
8 years for U-235 and 245,500 years for U-234.
Deviations from natural abundance
We now know from study of the sun and primitive meteorites that our solar system was initially almost homogeneous in isotopic composition. Deviations from the (evolving) galactic average, locally-sampled around the time that the sun's nuclear burning began, can generally be accounted for by mass fractionation plus a limited number of nuclear decay and transmutation processes. There is also evidence for injection of short-lived (now extinct) isotopes from a nearby supernova explosion that may have triggered solar nebula collapse. Hence deviations from natural abundance on earth are often measured in parts per thousand (per mil or ‰) because they're less than one percent (%).
The single exception to this lies with the
presolar grains found in primitive meteorites. These bypassed the homogenization, and often carry the nuclear signature of specific nucleosynthesis processes in which their elements were made. In these materials, deviations from "natural abundance" are sometimes measured in
factors of 100.
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