Isotopes of oxygen

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Late in a massive star's life, ¹⁶O concentrates in the O-shell, ¹⁷O in the H-shell and ¹⁸O in the He-shell

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.

[edit] Stable isotopes

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]

Relative and absolute abundance of ¹⁶O is due to it being a principal product of stellar evolution and the fact that 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 nuclei, 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]

[edit] Radioisotopes

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).

[edit] Table

• The precision of the isotope abundances and atomic mass is limited through variations. The given ranges should be applicable to any normal terrestrial material.
• Values marked # are not purely derived from experimental data, but at least partly from systematic trends. Spins with weak assignment arguments are enclosed in parentheses.
• Uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC which use expanded uncertainties.

[edit] See also

• Dole effect

[edit] Notes

1. ^ ^{a b c d e} Oxygen Nuclides / Isotopes. EnvironmentalChemistry.com. Retrieved on 2007-12-17.
2. ^ ^{a b c d} Meyer, B.S. (September 19-21, 2005). "NUCLEOSYNTHESIS AND GALACTIC CHEMICAL EVOLUTION OF THE ISOTOPES OF OXYGEN" in Workgroup on Oxygen in the Earliest Solar System. Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute. 9022. Retrieved on 2007-12-23. 
3. ^ Emsley 2001, p.297
4. ^ Mellor 1939, Chapter VI, Section 7
5. ^ Cook 1968, p.500
6. ^ Dansgaard, W (1964) Stable isotopes in precipitation. Tellus 16, 436-468

[edit] References

For the table

• Isotope masses from Ame2003 Atomic Mass Evaluation by G. Audi, A.H. Wapstra, C. Thibault, J. Blachot and O. Bersillon in Nuclear Physics A729 (2003).
• Isotopic compositions and standard atomic masses from Atomic weights of the elements. Review 2000 (IUPAC Technical Report). Pure Appl. Chem. Vol. 75, No. 6, pp. 683-800, (2003) and Atomic Weights Revised (2005).
• Half-life, spin, and isomer data selected from these sources. Editing notes on this article's talk page.
  • Audi, Bersillon, Blachot, Wapstra. The Nubase2003 evaluation of nuclear and decay properties, Nuc. Phys. A 729, pp. 3-128 (2003).
  • National Nuclear Data Center, Brookhaven National Laboratory. Information extracted from the NuDat 2.1 database (retrieved Sept. 2005).
  • David R. Lide (ed.), Norman E. Holden in CRC Handbook of Chemistry and Physics, 85th Edition, online version. CRC Press. Boca Raton, Florida (2005). Section 11, Table of the Isotopes.

For the prose

• Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen", in Clifford A. Hampel: The Encyclopedia of the Chemical Elements. New York: Reinhold Book Corporation, 499-512. LCCN 68-29938. 
• Emsley, John (2001). "Oxygen", Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press, 297-304. ISBN 0198503407. 
• Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry, 6th edition, London: Longmans, Green and Co.