Isotopes of oxygen

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 (¹⁶O, ¹⁷O, and ¹⁸O). Radioactive isotopes with mass numbers from ¹²O to ²⁴O have also been characterized, all short-lived, with the longest-lived being ¹⁵O with a half-life of 122.24 seconds. The shortest-lived is ¹²O with a half-life of 580(30)×10⁻²⁴ second.

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); thus oxygen (O) has a relative atomic mass of 15.9994(3). Known oxygen isotopes range in mass number from 12 to 24.

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.^[1] 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.^[1]

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.^[1] 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.^[1] Approximately a billion degrees Celsius is required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.^[2]

Measurements of the ratio of oxygen-18 to oxygen-16 are often used to interpret changes in paleoclimate. The isotopic composition of oxygen atoms in the Earth's atmosphere is 99.759% ¹⁶O, 0.037% ¹⁷O and 0.204% ¹⁸O.^[3] Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation,^[4] fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope ¹⁸O than air (0.204%) or seawater (0.1995%). This disparity allows analysis of temperature patterns via historic ice cores.

An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon ¹²C.^[5] Since physicists referred to ¹⁶O only, while chemists meant the naturally-abundant mixture of isotopes, this led to slightly different mass scales between the two disciplines.

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.^[6] 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 (ms).^[6] For example, ²⁴O has a half-life of 61 ms.^[7] The most common decay mode for isotopes lighter than the stable isotopes is β⁺ decay (to nitrogen)^[8][9][10] and the most common mode after is β⁻ decay (to fluorine).

Oxygen-13

Oxygen-13 is an unstable isotope of oxygen. It consists of 8 protons and electrons, and 5 neutrons. It has a spin of 3/2-, and a half-life of 8.58 ms. Its atomic mass is 13.0248 Da. It decays to nitrogen-13 by electron capture, and has a decay energy of 17.765 MeV.^[11] Its parent nuclide is fluorine-14.^[12]

Oxygen-15

Oxygen-15 is an isotope of oxygen, frequently used in positron emission tomography, or PET imaging. It has 8 protons, 7 neutrons, and 8 electrons. The total atomic mass is 15.0030654 amu. It has a half-life of 122.24 seconds.^[13] Oxygen-15 is synthesized through deuteron bombardment of nitrogen-14 using a cyclotron.^[14]

Table

nuclide symbol	Z(p)	N(n)	isotopic mass (u)	half-life	decay mode(s)[15]	daughter isotope(s)[n 1]	nuclear spin	representative isotopic composition (mole fraction)	range of natural variation (mole fraction)
^12O	8	4	12.034405(20)	580(30)×10^−24 s [0.40(25) MeV]	2p (60.0%)	^10C	0+
					p (40.0%)	^11N
^13O	8	5	13.024812(10)	8.58(5) ms	β^+ (89.1%)	^13N	(3/2−)
					β^+, p (10.9%)	^12C
^14O	8	6	14.00859625(12)	70.598(18) s	β^+	^14N	0+
^15O	8	7	15.0030656(5)	122.24(16) s	β^+	^15N	1/2−
^16O[n 2]	8	8	15.99491461956(16)	Stable			0+	0.99757(16)	0.99738–0.99776
^17O[n 3]	8	9	16.99913170(12)	Stable			5/2+	3.8(1)×10^−4	3.7×10^−4–4.0×10^−4
^18O[n 2][n 4]	8	10	17.9991610(7)	Stable			0+	2.05(14)×10^−3	1.88×10^−3–2.22×10^−3
^19O	8	11	19.003580(3)	26.464(9) s	β^−	^19F	5/2+
^20O	8	12	20.0040767(12)	13.51(5) s	β^−	^20F	0+
^21O	8	13	21.008656(13)	3.42(10) s	β^−	^21F	(1/2,3/2,5/2)+
^22O	8	14	22.00997(6)	2.25(15) s	β^− (78.0%)	^22F	0+
					β^−, n (22.0%)	^21F
^23O	8	15	23.01569(13)	82(37) ms	β^−, n (57.99%)	^22F	1/2+#
					β^− (42.0%)	^23F
^24O	8	16	24.02047(25)	65(5) ms	β^−, n (57.99%)	^23F	0+
					β^− (42.01%)	^24F

1. Bold for stable isotopes
2. The ratio between ¹⁶O and ¹⁸O is used to deduce ancient temperatures
3. Can be used in NMR studies of metabolic pathways
4. Can be used in studying certain metabolic pathways

Oxygen isotopes with 25–28 nucleons exist for less than 1 microsecond, since they are beyond the neutron drip line.^[16]

Notes

• 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.
• Nuclide masses are given by IUPAP Commission on Symbols, Units, Nomenclature, Atomic Masses and Fundamental Constants (SUNAMCO).
• Isotope abundances are given by IUPAC Commission on Isotopic Abundances and Atomic Weights.

See also

• Dole effect

Notes and references

1. B. S. Meyer (September 19–21, 2005). "Nucleosynthesis and galactic chemical evolution of the isotopes of oxygen" (PDF). Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute. Workgroup on Oxygen in the Earliest Solar System. Gatlinburg, Tennessee. 9022.
2. Emsley 2001, p. 297.
3. Cook 1968, p. 500.
4. Dansgaard, W (1964). "Stable isotopes in precipitation" (PDF). Tellus 16: 436–468. doi:10.1111/j.2153-3490.1964.tb00181.x.
5. Parks & Mellor 1939, Chapter VI, Section 7.
6. K. L. Barbalace. "Periodic Table of Elements: O - Oxygen". EnvironmentalChemistry.com. Retrieved 2007-12-17.
7. Ekström, L. P.; Firestone, R. B. (28 February 1999). "Oxygen-24". WWW Table of Radioactive Isotopes. LUNDS Universitet, LBNL Isotopes Project. Retrieved 2009-06-08.
8. "NUDAT". Retrieved 2009-07-06.
9. "NUDAT". Retrieved 2009-07-06.
10. "NUDAT". Retrieved 2009-07-06.
11. "Periodic Table of Elements: O - Oxygen". EnvironmentalChemistry.com. 1995-10-22. Retrieved 2014-12-02.
12. "Periodic Table of Elements: F - Fluorine". EnvironmentalChemistry.com. 1995-10-22. Retrieved 2014-12-02.
13. "oxygen 15 - definition of oxygen 15 by Medical dictionary". Medical-dictionary.thefreedictionary.com. Retrieved 2014-12-02.
14. "Production of PET Radionuclides". Austin Hospital, Austin Health. Retrieved 6 December 2012.
15. "Universal Nuclide Chart". Nucleonica. Retrieved 2014-12-02. (registration required (help)).
16. "Nuclear Physicists Examine Oxygen's Limits". Science Daily. 2007-09-18. Retrieved 2010-09-16.

References

For the table

• Isotope masses from the area:
  • G. Audi; A. H. Wapstra; C. Thibault; J. Blachot; O. Bersillon (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A 729: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001.
• Isotopic compositions and standard atomic masses from:
  • J. R. de Laeter; J. K. Böhlke; P. De Bièvre; H. Hidaka; H. S. Peiser; K. J. R. Rosman; P. D. P. Taylor (2003). "Atomic weights of the elements. Review 2000 (IUPAC Technical Report)" (PDF). Pure and Applied Chemistry 75 (6): 683–800. doi:10.1351/pac200375060683.
  • M. E. Wieser (2006). "Atomic weights of the elements 2005 (IUPAC Technical Report)". Pure and Applied Chemistry 78 (11): 2051–2066. doi:10.1351/pac200678112051. Lay summary.
• Half-life, spin, and isomer data selected from the following sources. See editing notes on this article's talk page.
  • G. Audi; A. H. Wapstra; C. Thibault; J. Blachot; O. Bersillon (2003). "The NUBASE evaluation of nuclear and decay properties" (PDF). Nuclear Physics A 729: 3–128. Bibcode:2003NuPhA.729....3A. doi:10.1016/j.nuclphysa.2003.11.001.
  • National Nuclear Data Center. "NuDat 2.1 database". Brookhaven National Laboratory. Retrieved September 2005.
  • N. E. Holden (2004). "Table of the Isotopes". In D. R. Lide. CRC Handbook of Chemistry and Physics (85th ed.). CRC Press. Section 11. ISBN 978-0-8493-0485-9.

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. pp. 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. pp. 297–304. ISBN 0-19-850340-7.
• Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry (6th ed.). London: Longmans, Green and Co.