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Isotopes of oxygen

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Isotopes of oxygen (8O)
Main isotopes[1] Decay
Isotope abun­dance half-life (t1/2) mode pro­duct
15O trace 122.27 s β+ 15N
16O 99.8% stable
17O 0.0384% stable
18O 0.205% stable
Standard atomic weight Ar°(O)

There are three known stable isotopes of oxygen (8O): 16
O, 17
O, and 18
O. Radioisotopes are known from 11O to 28O (particle-bound from mass number 13 to 24), and the most stable are 15
O with half-life 122.27 seconds and 14
O with half-life 70.62 seconds. All remaining radioisotopes are even shorter in lifetime. The four heaviest known isotopes (up to 28
O) decay by neutron emission to 24
O, whose half-life is 77 milliseconds; 24O, along with 28Ne, have been used in the model of reactions in the crust of neutron stars.[4] 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.

List of isotopes

[edit]
Nuclide
[n 1] 		Z N Isotopic mass (Da)[5]
[n 2] 		Half-life[1]

[resonance width] 		Decay
mode[1]
[n 3] 		Daughter
isotope
[n 4] 		Spin and
parity[1]
[n 5][n 6] 		Natural abundance (mole fraction)
Excitation energy Normal proportion[1] Range of variation
11
O[6] 		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
β+p,α (<0.1%) 24
He[7]
14
O 		8 6 14.008596706(27) 70.621(11) s β+ 14
N 		0+
15
O[n 7] 		8 7 15.0030656(5) 122.266(43) s β+ 15
N 		1/2− Trace[8]
16
O[n 8] 		8 8 15.994914619257(319) Stable 0+ [0.99738, 0.99776][9]
17
O[n 9] 		8 9 16.999131755953(692) Stable 5/2+ [0.000367, 0.000400][9]
18
O[n 8][n 10] 		8 10 17.999159612136(690) Stable 0+ [0.00187, 0.00222][9]
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[n 11] 		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+
27
O[10] 		8 19 ≥ 2.5 zs n 26
O 		(3/2+, 7/2−)
28
O[10] 		8 20 ≥ 650 ys 2n 26
O 		0+
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. ^ Intermediate product of CNO-I in stellar nucleosynthesis as part of the process producing helium from hydrogen
  8. ^ a b The ratio between 16
    O     and 18
    O     is used to deduce ancient temperatures.
  9. ^ Can be used in NMR studies of metabolic pathways.
  10. ^ Can be used in studying certain metabolic pathways.
  11. ^ Heaviest particle-bound isotope of oxygen, see Nuclear drip line

Oxygen-14

[edit]

Oxygen-14 (half-life 70.62 seconds) is the second most stable radioisotope of oxygen, and decays by positron emission to nitrogen-14.

Oxygen-14 ion beams are of interest to researchers of proton-rich nuclei; for example, one early experiment at the Facility for Rare Isotope Beams in East Lansing, Michigan, produced a 14O beam by proton bombardment of 14N,[11][12] using it to determine the absolute strength of the electron capture transition.

Oxygen-15

[edit]

Oxygen-15 (half-life 122.27 seconds) is the most stable radioisotope of oxygen, decaying by positron emission to nitrogen-15.

It is thus the isotope of oxygen 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 is produced for this application through deuteron bombardment of nitrogen-14 using a cyclotron.[15]

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[16] out of 16O and 14N:[17]

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 of 2 minutes, 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
N floats by, carried by the wind.[8]

Oxygen-16

[edit]

Oxygen-16 (symbol: 16O or 16
8O) is a stable isotope of oxygen, with 8 neutrons and 8 protons in its nucleus, making it a doubly magic nuclide. It is the most abundant isotope of oxygen, accounting for about 99.76% of all oxygen.

The relative and absolute abundances of oxygen-16 are high because it is a principal product of stellar evolution. It can be made by stars that were initially made exclusively of hydrogen.[18] Most oxygen-16 is synthesized at the end of the helium fusion process in stars. The triple-alpha process creates carbon-12, which captures an additional helium-4 to make oxygen-16. It is also created by the neon-burning process.

Prior to the definition of the dalton based on 12
C, one atomic mass unit was defined as one sixteenth of the mass of an oxygen-16 atom.[19] Since physicists referred to 16
O only, while chemists meant the natural mix of isotopes, this led to slightly different mass scales.

Oxygen-17

[edit]

Oxygen-17 (17O) is a stable isotope of oxygen with a low isotopic abundance of about 0.038%. 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.[18]

As the only stable isotope of oxygen possessing a nuclear spin (+5⁄2) and a favorable characteristic of field-independent relaxation in liquid water, 17O enables NMR studies of oxidative metabolic pathways through compounds containing 17O (i.e. metabolically produced H217O water by oxidative phosphorylation in mitochondria[20]) at high magnetic fields.

Water used as nuclear reactor coolant is subjected to intense neutron flux. Natural water starts out with 0.038% of 17O; heavy water starts out incidentally enriched to about 0.055% in that isotopes. Further, the neutron flux slowly converts 16O in the cooling water to 17O by neutron capture, increasing its concentration. The neutron flux slowly converts 17O (with much greater cross section) in the cooling water to carbon-14, an undesirable product that can escape to the environment:

17O (n,α) → 14C

Some tritium removal facilities make a point of replacing the oxygen of the water with natural oxygen (mostly 16O) to give the added benefit of reducing 14C production.[21][22]

The isotope was first hypothesized and subsequently imaged by Patrick Blackett in Rutherford's lab in 1925:[23] It was a product out of the first man-made transmutation of 14N and 4He2+ conducted by Frederick Soddy and Ernest Rutherford in 1917–1919.[24] Its presence in Earth's atmosphere was later detected in 1929 by Giauque and Johnson in absorption spectra,[25] demonstrating its natural existece.

Oxygen-18

[edit]

Oxygen-18 (18
O, Ω[26]) is one of the stable isotopes of oxygen, with roughly 0.20% abundance, and considered one of the environmental isotopes. 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.[18] Temperatures on the order of 109 kelvins are needed to fuse oxygen into sulfur.[27]

Fluorine-18 is usually produced by irradiation of 18O-enriched water with high-energy (about 18 MeV) protons prepared in a cyclotron or a linear accelerator, yielding an aqueous solution containing 18F as fluoride ion. This solution is then used for rapid synthesis of a labeled molecule, often with the fluorine atom replacing a hydroxy group. The labeled molecules or radiopharmaceuticals have to be synthesized after the radiofluorine is prepared, as the high energy proton radiation would destroy the molecules. Large amounts of oxygen-18 enriched water are used in positron emission tomography centers, for on-site production of 18F-labeled fluorodeoxyglucose (FDG).[28]

Measurements of the 18O/16O ratio (known as δ18
O) are often used in paleoclimatology. 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%).[29] This disparity allows the study of historical temperature patterns via the analysis of ice cores. Assuming that atmospheric circulation and elevation has not changed significantly over the poles, the temperature of ice formation can be calculated as equilibrium fractionation between phases of water that is known for different temperatures. Water molecules are also subject to Rayleigh fractionation as atmospheric water moves from the equator poleward which results in progressive depletion of 18
O, or lower δ18
O values.[30]

The δ18
O ratio can also be used in paleothermometry for certain types of fossils. The fossil material used is generally calcite or aragonite, however oxygen isotope paleothermometry has also been done of phosphatic fossils using SHRIMP.[31] For determination of ocean temperatures over geologic time, multiple fossils of the same species in different stratigraphic layers would be measured, and the difference between them would indicate long term changes.[32]

18O has also been used to trace ocean composition and temperature which seafood is from.[33]

In the study of plants' photorespiration, the labeling of atmosphere by oxygen-18 allows for the measurement of oxygen uptake by the photorespiration pathway. Labeling by 18
O
2 gives the unidirectional flux of O
2 uptake, while there is a net photosynthetic 16
O
2 evolution. It was demonstrated that, under preindustrial atmosphere, most plants reabsorb, by photorespiration, half of the oxygen produced by photosynthesis. Then, the yield of photosynthesis was halved by the presence of oxygen in atmosphere.[34][35]

Oxygen-20

[edit]

Oxygen-20 has a half-life of 13.51±0.05 s and decays by β− decay to 20F. It is one of the known cluster decay ejected particles, being emitted in the decay of 228Th with a branching ratio of about (1.13±0.22)×10−13.[36]

See also

[edit]

Daughter products other than oxygen

References

[edit]
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