There are three stable isotopes of oxygen that lead to oxygen (O) having a standard atomic mass of 15.9994(3) u. There are also 14 other isotopes that have unstable nuclei.
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Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O being the most abundant (99.762% natural abundance).[1] Oxygen isotopes range in mass number from 12 to 28.[1]
The relative and absolute abundance of 16O is high because it is a principal product of stellar evolution and because it is a primary isotope, meaning it can be made by stars that were initially made exclusively of hydrogen.[2] Most 16O is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates 12C, which captures an additional 4He to make 16O. The neon burning process creates additional 16O.[2]
Both 17O and 18O are secondary isotopes, meaning that their nucleosynthesis requires seed nuclei. 17O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[2] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of stars.[2] Approximately a billion degrees Celsius is required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.[3]
Fourteen radioisotopes have been characterized, with the most stable being 15O with a half-life of 122.24 s and 14O with a half-life of 70.606 s.[1] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds.[1] The most common decay mode before the stable isotopes is electron capture and the most common mode after is beta decay. The decay products before the stable isotopes are element 7 (nitrogen) isotopes and the products after are element 9 (fluorine) isotopes.[1]
An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C.[4] Since physicists referred to 16O only, while chemists meant the naturally-abundant mixture of isotopes, this led to slightly different atomic mass scales.
The isotopic composition of oxygen atoms in the earth's atmosphere is 99.759% 16O, 0.037% 17O and 0.204% 18O.[5] Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation[6], fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope 18O than air (0.204%) or seawater containing (0.1995%). For this reason, tracking this ratio is used by scientists to estimate past climate change (see section on scientific uses).
| nuclide symbol |
Z(p) | N(n) | isotopic mass (u) |
half-life | nuclear spin |
representative isotopic composition (mole fraction) |
range of natural variation (mole fraction) |
|---|---|---|---|---|---|---|---|
| excitation energy | |||||||
| 12O | 8 | 4 | 12.034405(20) | 580(30)E-24 s [0.40(25) MeV] | 0+ | ||
| 13O | 8 | 5 | 13.024812(10) | 8.58(5) ms | (3/2-) | ||
| 14O | 8 | 6 | 14.00859625(12) | 70.598(18) s | 0+ | ||
| 15O | 8 | 7 | 15.0030656(5) | 122.24(16) s | 1/2- | ||
| 16O | 8 | 8 | 15.99491461956(16) | STABLE | 0+ | 0.99757(16) | 0.99738-0.99776 |
| 17O | 8 | 9 | 16.99913170(12) | STABLE | 5/2+ | 0.00038(1) | 0.00037-0.00040 |
| 18O | 8 | 10 | 17.9991610(7) | STABLE | 0+ | 0.00205(14) | 0.00188-0.00222 |
| 19O | 8 | 11 | 19.003580(3) | 26.464(9) s | 5/2+ | ||
| 20O | 8 | 12 | 20.0040767(12) | 13.51(5) s | 0+ | ||
| 21O | 8 | 13 | 21.008656(13) | 3.42(10) s | (1/2,3/2,5/2)+ | ||
| 22O | 8 | 14 | 22.00997(6) | 2.25(15) s | 0+ | ||
| 23O | 8 | 15 | 23.01569(13) | 82(37) ms | 1/2+# | ||
| 24O | 8 | 16 | 24.02047(25) | 65(5) ms | 0+ | ||
| 25O | 8 | 17 | 25.02946(28)# | <50 ns | (3/2+)# | ||
| 26O | 8 | 18 | 26.03834(28)# | <40 ns | 0+ | ||
| 27O | 8 | 19 | 27.04826(54)# | <260 ns | 3/2+# | ||
| 28O | 8 | 20 | 28.05781(64)# | <100 ns | 0+ | ||
| Isotopes of nitrogen | Isotopes of oxygen | Isotopes of fluorine |
| Index to isotope pages · Table of nuclides | ||