Inphysics,natural abundance (NA) refers to the abundance ofisotopes of achemical element as naturally found on aplanet. The relative atomic mass (a weighted average, weighted bymole-fraction abundance figures) of these isotopes is theatomic weight listed for the element in theperiodic 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 hasthree naturally occurring isotopes:238U,235U, and234U. Their respective natural mole-fraction abundances are 99.2739–99.2752%, 0.7198–0.7202%, and 0.0050–0.0059%.[1] For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,274238U atoms, approximately 720235U atoms, and very few (most likely 5 or 6)234U atoms. This is because238U is much more stable than235U or234U, as thehalf-life of each isotope reveals: 4.468 billion years for238U compared with 7.038 × 108 years for235U and 245,500 years for234U.
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 billion years ago the NA of235U was 3.1% compared with today's 0.7%, and that allowed anatural 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 innucleosynthesis (as in the case ofsamarium; radioactive147Sm and148Sm are much more abundant than stable144Sm) and by production of a given isotope as a daughter of natural radioactive isotopes (as in the case of radiogenicisotopes of lead).
It is now known from study of the Sun and primitive meteorites that theSolar 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 (see the article onmass-independent fractionation) plus a limited number of nuclear decay and transmutation processes.[2] There is also evidence for injection of short-lived (now-extinct) isotopes from a nearby supernova explosion that may have triggered solar nebula collapse.[3] Hence deviations from natural abundance on Earth are often measured in parts per thousand (per mille or ‰) because they are less than one percent (%).
An exception to this lies with thepresolar grains found in primitive meteorites. These small grains condensed in the outflows of evolved ("dying") stars and escaped the mixing and homogenization processes in the interstellar medium and the solar accretion disk (also known as the solar nebula or protoplanetary disk).[4][clarification needed] As stellar condensates ("stardust"), these grains carry the isotopic signatures of specific nucleosynthesis processes in which their elements were made.[5] In these materials, deviations from "natural abundance" are sometimes measured in factors of 100.[citation needed][4]
The next table gives theterrestrial isotope distributions for some elements. Some elements, such asphosphorus andfluorine, only exist as a single isotope, with a natural abundance of 100%.
| Isotope | % nat. abundance | atomic mass |
|---|---|---|
| 1H | 99.985 | 1.007825 |
| 2H | 0.015 | 2.0140 |
| 12C | 98.89 | 12 (formerly by definition) |
| 13C | 1.11 | 13.00335 |
| 14N | 99.64 | 14.00307 |
| 15N | 0.36 | 15.00011 |
| 16O | 99.76 | 15.99491 |
| 17O | 0.04 | 16.99913 |
| 18O | 0.2 | 17.99916 |
| 28Si | 92.23 | 27.97693 |
| 29Si | 4.67 | 28.97649 |
| 30Si | 3.10 | 29.97376 |
| 32S | 95.0 | 31.97207 |
| 33S | 0.76 | 32.97146 |
| 34S | 4.22 | 33.96786 |
| 35Cl | 75.77 | 34.96885 |
| 37Cl | 24.23 | 36.96590 |
| 79Br | 50.69 | 78.9183 |
| 81Br | 49.31 | 80.9163 |
