| |||||||||||||||||||||||||||||||
Standard atomic weight Ar°(O) | |||||||||||||||||||||||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
There are three known stable
isotopes of
oxygen (8O):
16
O
,
17
O
, and
18
O
.
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 the
unbound 11
O
with a half-life of 198(12)
yoctoseconds, though half-lives have not been measured for the unbound heavy isotopes 27
O
and 28
O
.
[3]
Nuclide [n 1] |
Z | N |
Isotopic mass (
Da)
[4] [n 2] |
Half-life
[5] [ resonance width] |
Decay mode [5] [n 3] |
Daughter isotope [n 4] |
Spin and parity [5] [n 5] [n 6] |
Natural abundance (mole fraction) | |||||||||||
---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
Excitation energy | Normal proportion [5] | 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 ? [n 11] | 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 12] |
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 [3] |
8 | 19 | ≥ 2.5 zs | n | 26 O |
(3/2+, 7/2−) | |||||||||||||
28 O [3] |
8 | 20 | ≥ 650 ys | 2n | 26 O |
0+ | |||||||||||||
This table header & footer: |
n: | Neutron emission |
p: | Proton emission |
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.
[10] 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
.
[10]
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.
[10] 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.
[10] About 109
kelvin is needed to
fuse oxygen into
sulfur.
[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.
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
.
[13]
Water molecules with a lighter isotope are slightly more likely to
evaporate and less likely to fall as
precipitation,
[14] 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. [15] Researchers need to avoid improper or prolonged storage of the samples for accurate measurements. [15]
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. [16]
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. The four heaviest known isotopes (up to 28
O
) decay by
neutron emission to
24
O
, whose half-life is 77.4(4.5) ms. This isotope, along with
28Ne, have been used in the model of reactions in crust of neutron stars.
[17] 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 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-14 is the second most stable radioisotope. 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, used a 14O beam to study the beta decay transition of this isotope to 14N. [18] [19]
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. [20] [21] 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. [22]
Oxygen-15 and nitrogen-13 are produced in air when gamma rays (for example from lightning) knock neutrons out of 16O and 14N: [23]
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-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. [24]
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