Isotopes of oxygen

There are three known stable isotopes of oxygen (8O): 16
O
, 17
O
, and 18
O
.

Main isotopes of oxygen (8O)
Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
16O [99.738%, 99.776%] stable
17O [0.0367%, 0.0400%] stable
18O [0.187%, 0.222%] stable
Standard atomic weight Ar, standard(O)[15.99903, 15.99977] conventional: 15.999[1][2]

Radioactive isotopes ranging from 11
O
 to 28
O
 have also been characterized, all short-lived. The longest-lived radioisotope is 15
O
 with a half-life of 122.266(43) s, while the shortest-lived isotope is 11
O
 with a half-life of 198(12) yoctoseconds (though the half-lives of the neutron-unbound 27
O
 and 28
O
 are still unknown).

List of isotopes
Nuclide[3]
[n 1]		 Z N Isotopic mass (Da)[4]
[n 2]		 Half-life

[resonance width]		 Decay
mode
[n 3]		 Daughter
isotope
[n 4]		 Spin and
parity
[n 5][n 6]		 Natural abundance (mole fraction)
Excitation energy Normal proportion Range of variation
11
O
[5]		 8 3 11.051250(60) 198(12) ys
[2.31(14) MeV]		 2p 9
C
 (3/2−)
12
O
 8 4 12.034368(13) 8.9(3.3) zs 2p 10
C
 0+
13
O
 8 5 13.024815(10) 8.58(5) ms β+ (89.1(2)%) 13
N
 (3/2−)
β+p (10.9(2)%) 12
C
14
O
 8 6 14.008596706(27) 70.621(11) s β+ 14
N
 0+
15
O
 8 7 15.0030656(5) 122.266(43) s β+ 15
N
 1/2−
16
O
[n 7]		 8 8 15.994914619257(319) Stable 0+ [0.99738, 0.99776][6]
17
O
[n 8]		 8 9 16.999131755953(692) Stable 5/2+ [0.000367, 0.000400][6]
18
O
[n 7][n 9]		 8 10 17.999159612136(690) Stable 0+ [0.00187, 0.00222][6]
19
O
 8 11 19.0035780(28) 26.470(6) s β− 19
F
 5/2+
20
O
 8 12 20.0040754(9) 13.51(5) s β− 20
F
 0+
21
O
 8 13 21.008655(13) 3.42(10) s β− 21
F
 (5/2+)
β−n ?[n 10] 20
F
 ?
22
O
 8 14 22.009970(60) 2.25(9) s β− (> 78%) 22
F
 0+
β−n (< 22%) 21
F
23
O
 8 15 23.015700(130) 97(8) ms β− (93(2)%) 23
F
 1/2+
β−n (7(2)%) 22
F
24
O
 8 16 24.019860(180) 77.4(4.5) ms β− (57(4)%) 24
F
 0+
β−n (43(4)%) 23
F
25
O
 8 17 25.029340(180) 5.18(35) zs n 24
O
 3/2+#
26
O
 8 18 26.037210(180) 4.2(3.3) ps 2n 24
O
 0+
27
O
 8 19 27.047960(540)# < 260 ns n ?[n 10] 26
O
 ?		 3/2+#
2n ?[n 10] 25
O
 ?
28
O
 8 20 28.055910(750)# < 100 ns 2n ?[n 10] 26
O
 ?		 0+
β− (0%) 28
F
This table header & footer:
  1. mO – Excited nuclear isomer.
  2. ( ) – Uncertainty (1σ) is given in concise form in parentheses after the corresponding last digits.
  3. Modes of decay:
    n:Neutron emission
    p:Proton emission
  4. Bold symbol as daughter – Daughter product is stable.
  5. ( ) spin value – Indicates spin with weak assignment arguments.
  6. # – Values marked # are not purely derived from experimental data, but at least partly from trends of neighboring nuclides (TNN).
  7. The ratio between 16
    O
     and 18
    O
     is used to deduce ancient temperatures.
  8. Can be used in NMR studies of metabolic pathways.
  9. Can be used in studying certain metabolic pathways.
  10. Decay mode shown is energetically allowed, but has not been experimentally observed to occur in this nuclide.

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

Natural oxygen is made of three stable isotopes, 16
O
, 17
O
, and 18
O
, with 16
O
 being the most abundant (99.762% natural abundance). Depending on the terrestrial source, the standard atomic weight varies within the range of [15.99903, 15.99977] (the conventional value is 15.999).

16
O
 has high relative and absolute abundance 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 hydrogen only.[7] Most 16
O
 is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates 12
C
, which captures an additional 4
He
 nucleus to produce 16
O
. The neon burning process creates additional 16
O
.[7]

Both 17
O
 and 18
O
 are secondary isotopes, meaning their synthesis requires seed nuclei. 17
O
 is primarily made by burning hydrogen into helium in the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[7] Most 18
O
 is produced when 14
N
 (made abundant from CNO burning) captures a 4
He
 nucleus, becoming 18
F
. This quickly decays to 18
O
 making that isotope common in the helium-rich zones of stars.[7] About 109 kelvin is needed to fuse oxygen into sulfur.[8]

Measurements of 18O/16O ratio are often used to interpret changes in paleoclimate. Oxygen in Earth's air is 99.759% 16
O
, 0.037% 17
O
 and 0.204% 18
O
.[9] Water molecules with a lighter isotope are slightly more likely to evaporate and less likely to fall as precipitation,[10] so Earth's freshwater and polar ice have slightly less (0.1981%) 18
O
 than air (0.204%) or seawater (0.1995%). This disparity allows analysis of temperature patterns via historic ice cores.

Solid samples (organic and inorganic) for oxygen isotopic ratios are usually stored in silver cups and measured with pyrolysis and mass spectrometry.[11] Researchers need to avoid improper or prolonged storage of the samples for accurate measurements.[11]

An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based on 12
C
.[12] Since physicists referred to 16
O
 only, while chemists meant the natural mix of isotopes, this led to slightly different mass scales.

Radioisotopes

13 radioisotopes have been characterized; the most stable are 15
O
 with half-life 122.266(43) s and 14
O
 with half-life 70.621(11) s. All remaining radioisotopes have half-lives less than 27 s and most have half-lives less than 0.1 s. 24
O
 has half-life 77.4(4.5) ms. The most common decay mode for isotopes lighter than the stable isotopes is β+ decay to nitrogen, and the most common mode after is β− decay to fluorine.

Oxygen-13

13O is an unstable isotope, with 8 protons and 5 neutrons. It has spin 3/2-, and half-life 8.58(5) ms. Its atomic mass is 13.024815(10) Da. It decays to nitrogen-13 by electron capture, with a decay energy of 17.765 MeV. Its parent nuclide is fluorine-14.

Oxygen-15

15O is a radioisotope, often used in positron emission tomography (PET). It can be used in, among other things, water for PET myocardial perfusion imaging and for brain imaging.[13][14] It has an atomic mass of 15.0030656(5), and a half-life of 122 s. It is made through deuteron bombardment of nitrogen-14 using a cyclotron.[15]

Oxygen-15 and nitrogen-13 are produced in air when gamma rays (for example from lightning) knock neutrons out of 16O and 14N:[16]

16
O
 + γ → 15
O
 + n
14
N
 + γ → 13
N
 + n

15O decays to 15N, emitting a positron. The positron quickly annihilates with an electron, producing two gamma rays of about 511 keV. After a lightning bolt, this gamma radiation dies down with half-life 2 min, but these low-energy gamma rays go on average only about 90 metres through the air. Together with rays produced from positrons from nitrogen-13 they may only be detected for a minute or so as the "cloud" of 15
O
 and 13
O
 floats by, carried by the wind.[17]

See also

References
  1. "Standard Atomic Weights: Oxygen". CIAAW. 2009.
  2. Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
  3. Half-life, decay mode, nuclear spin, and isotopic composition is sourced in:
    Kondev, F.G.; Wang, M.; Huang, W.J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
  4. Wang, Meng; Huang, W.J.; Kondev, F.G.; Audi, G.; Naimi, S. (2021). "The AME 2020 atomic mass evaluation (II). Tables, graphs and references*". Chinese Physics C. 45 (3): 030003. doi:10.1088/1674-1137/abddaf.
  5. Webb, T. B.; et al. (2019). "First Observation of Unbound 11O, the Mirror of the Halo Nucleus 11Li". Physical Review Letters. 122 (12): 122501–1–122501–7. arXiv:1812.08880. Bibcode:2019PhRvL.122l2501W. doi:10.1103/PhysRevLett.122.122501. PMID 30978039. S2CID 84841752.
  6. "Atomic Weight of Oxygen | Commission on Isotopic Abundances and Atomic Weights". ciaaw.org. Retrieved 2022-03-15.
  7. 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.
  8. Emsley 2001, p. 297.
  9. Cook & Lauer 1968, p. 500.
  10. Dansgaard, W (1964). "Stable isotopes in precipitation" (PDF). Tellus. 16 (4): 436–468. Bibcode:1964TellA..16..436D. doi:10.1111/j.2153-3490.1964.tb00181.x.
  11. Tsang, Man-Yin; Yao, Weiqi; Tse, Kevin (2020). Kim, Il-Nam (ed.). "Oxidized silver cups can skew oxygen isotope results of small samples". Experimental Results. 1: e12. doi:10.1017/exp.2020.15. ISSN 2516-712X.
  12. Parks & Mellor 1939, Chapter VI, Section 7.
  13. Rischpler, Christoph; Higuchi, Takahiro; Nekolla, Stephan G. (22 November 2014). "Current and Future Status of PET Myocardial Perfusion Tracers". Current Cardiovascular Imaging Reports. 8 (1): 333–343. doi:10.1007/s12410-014-9303-z. S2CID 72703962.
  14. Kim, E. Edmund; Lee, Myung-Chul; Inoue, Tomio; Wong, Wai-Hoi (2012). Clinical PET and PET/CT: Principles and Applications. Springer. p. 182. ISBN 9781441908025.
  15. "Production of PET Radionuclides". Austin Hospital, Austin Health. Archived from the original on 15 January 2013. Retrieved 6 December 2012.
  16. Timmer, John (25 November 2017). "Lightning strikes leave behind a radioactive cloud". Ars Technica.
  17. Teruaki Enoto; et al. (Nov 23, 2017). "Photonuclear reactions triggered by lightning discharge". Nature. 551 (7681): 481–484. arXiv:1711.08044. Bibcode:2017Natur.551..481E. doi:10.1038/nature24630. PMID 29168803. S2CID 4388159.
  • Cook, Gerhard A.; Lauer, Carol M. (1968). "Oxygen". In Clifford A. Hampel (ed.). 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 978-0-19-850340-8.
  • Parks, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry (6th ed.). London: Longmans, Green and Co.
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