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Isotopes in medicine
Oxygen-15
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Oxygen-15

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Isotopes of oxygen (8O)
Main isotopes Decay
abun­dance half-life (t1/2) mode pro­duct
15O synth 122.266 s β+100% 15N
16O 99.8% stable
17O 0.0380% stable
18O 0.205% stable
Standard atomic weight Ar°(O)
  • [15.99903, 15.99977]
  • 15.999±0.001 (abridged)

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

Radioactive isotopes ranging from 11
O
 to 26
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.

List of isotopes

Nuclide
 Z N Isotopic mass (Da)
 Half-life

[resonance width] 		Decay
mode
 Daughter
isotope
 Spin and
parity
 Natural abundance (mole fraction)
Excitation energy Normal proportion Range of variation
11
O
 		8 3 11.05125(6) 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
 		8 8 15.994914619257(319) Stable 0+ [0.99738, 0.99776]
17
O
 		8 9 16.999131755953(692) Stable 5/2+ [0.000367, 0.000400]
18
O
 		8 10 17.999159612136(690) Stable 0+ [0.00187, 0.00222]
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 ? 20
F
 ?
22
O
 		8 14 22.00997(6) 2.25(9) s β− (> 78%) 22
F
 		0+
β−n (< 22%) 21
F
23
O
 		8 15 23.01570(13) 97(8) ms β− (93(2)%) 23
F
 		1/2+
β−n (7(2)%) 22
F
24
O
 		8 16 24.01986(18) 77.4(4.5) ms β− (57(4)%) 24
F
 		0+
β−n (43(4)%) 23
F
25
O
 		8 17 25.02934(18) 5.18(35) zs n 24
O
 		3/2+#
26
O
 		8 18 26.03721(18) 4.2(3.3) ps 2n 24
O
 		0+
This table header & footer:

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. Most 16
O
 is synthesized at the end of the helium fusion process in stars; the triple-alpha process creates 12
C
, which captures an additional 4
He
 nucleus to produce 16
O
. The neon burning process creates additional 16
O
.

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. Most 18
O
 is produced when 14
N
 (made abundant from CNO burning) captures a 4
He
 nucleus, becoming 18
F
. This quickly (half-life around 110 minutes) beta decays to 18
O
 making that isotope common in the helium-rich zones of stars. About 109kelvin is needed to fuse oxygen into sulfur.

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

Applications of various isotopes

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
.Water molecules with a lighter isotope are slightly more likely to evaporate and less likely to fall as precipitation, 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. Researchers need to avoid improper or prolonged storage of the samples for accurate measurements.

Due to natural oxygen being mostly 16
O, samples enriched with the other stable isotopes can be used for isotope labeling. For example, it was proven, that the oxygen released in photosynthesis originates in H2O, rather than in the also consumed CO2, by isotope tracing experiments. The oxygen contained in CO2 in turn is used to make up the sugars formed by photosynthesis.

In heavy water reactors the neutron moderator should preferably be low in 17
O and 18
O due to their higher neutron absorption cross section compared to 16
O. While this effect can also be observed in light water reactors, ordinary hydrogen (protium) has a higher absorption cross section than any stable isotope of oxygen and its number density is twice as high in water as that of oxygen so that the effect is negligible. As some methods of isotope separation enrich not only heavier isotopes of hydrogen but also heavier isotopes of oxygen when producing heavy water, the concentration of 17
O and 18
O can be measurably higher. Furthermore the 17
O(n,α)14
C reaction is a further undesirable result of an elevated concentration of heavier isotopes of oxygen. Therefore facilities which remove tritium from heavy water used in nuclear reactors often also remove or at least reduce the amount of heavier isotopes of oxygen.

Oxygen isotopes are also used to trace ocean composition and temperature which seafood is from.

Radioisotopes

Thirteen 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

Oxygen-13 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 17.770(10) MeV. Its parent nuclide is fluorine-14.

Oxygen-15

Oxygen-15 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. It has an atomic mass of 15.0030656(5), and a half-life of 122.266(43) s. It is produced through deuteron bombardment of nitrogen-14 using a cyclotron.

14
N
 + 2
H
 → 15
O
 + n

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

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

15
O
 decays to 15
N
, 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.

See also


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