Isotopes of oxygen

Isotopes of oxygen
There are three known stable isotopes of oxygen (8O): and .
Radioactive isotopes ranging from to have also been characterized, all short-lived. The longest-lived radioisotope is with a half-life of , while the shortest-lived isotope is the unbound with a half-life of , though half-lives have not been measured for the unbound heavy isotopes and .
List of isotopes
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 3
|
| []
| 2p
|
| (3/2−)
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 4
|
|
| 2p
|
| 0+
|
|
|-
| rowspan=3|
| rowspan=3 style="text-align:right" | 8
| rowspan=3 style="text-align:right" | 5
| rowspan=3|
| rowspan=3|
| β+
|
| rowspan=3|(3/2−)
| rowspan=3|
| rowspan=3|
|-
| β+p
|
|-
| β+p,α (<)
| 2
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 6
|
|
| β+
|
| 0+
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 7
|
|
| β+
|
| 1/2−
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 8
|
| colspan="3" style="text-align:center;"|Stable
| 0+
| colspan="2" style="text-align:center;"|[, ]
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 9
|
| colspan="3" style="text-align:center;"|Stable
| 5/2+
| colspan="2" style="text-align:center;"|[, ]
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 10
|
| colspan="3" style="text-align:center;"|Stable
| 0+
| colspan="2" style="text-align:center;"|[, ]
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 11
|
|
| β−
|
| 5/2+
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 12
|
|
| β−
|
| 0+
|
|
|-
| rowspan=2|
| rowspan=2 style="text-align:right" | 8
| rowspan=2 style="text-align:right" | 13
| rowspan=2|
| rowspan=2|
| β−
|
| rowspan=2|(5/2+)
| rowspan=2|
| rowspan=2|
|-
| β−n ?
| ?
|-
| rowspan=2|
| rowspan=2 style="text-align:right" | 8
| rowspan=2 style="text-align:right" | 14
| rowspan=2|
| rowspan=2|
| β− (> )
|
| rowspan=2|0+
| rowspan=2|
| rowspan=2|
|-
| β−n (< )
|
|-
| rowspan=2|
| rowspan=2 style="text-align:right" | 8
| rowspan=2 style="text-align:right" | 15
| rowspan=2|
| rowspan=2|
| β−
|
| rowspan=2|1/2+
| rowspan=2|
| rowspan=2|
|-
| β−n
|
|-
| rowspan=2|
| rowspan=2 style="text-align:right" | 8
| rowspan=2 style="text-align:right" | 16
| rowspan=2|
| rowspan=2|
| β−
|
| rowspan=2|0+
| rowspan=2|
| rowspan=2|
|-
| β−n
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 17
|
|
| n
|
| 3/2+
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 18
|
|
| 2n
|
| 0+
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 19
|
|
| 3n
|
|
|
|
|-
|
| style="text-align:right" | 8
| style="text-align:right" | 20
|
|
| 4n
|
| 0+
|
|
Stable isotopes
Natural oxygen is made of three stable isotopes , and , with being the most abundant (99.762% natural abundance). Depending on the terrestrial source, the standard atomic weight varies within the range of [, ] (the conventional value is 15.999).
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 is synthesized at the end of the helium fusion process in stars; the triple-alpha process creates , which captures an additional nucleus to produce . The neon burning process creates additional .
Both and are secondary isotopes, meaning their synthesis requires seed nuclei. 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 is produced when (made abundant from CNO burning) captures a nucleus, becoming . This quickly (half-life around 110 minutes) beta decays to making that isotope common in the helium-rich zones of stars. About 109 kelvin 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 . Since physicists referred to 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 , and . 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 than air or seawater. 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 , 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 , rather than in the also consumed , by isotope tracing experiments. The oxygen contained in in turn is used to make up the sugars formed by photosynthesis.
In heavy water reactors the neutron moderator should preferably be low in and due to their higher neutron absorption cross section compared to . 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 and can be measurably higher. Furthermore the (n,α) 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 with half-life and with half-life . All remaining radioisotopes have half-lives less than and most have half-lives less than 0.1 s. Four heaviest known isotopes (up to ) decay by neutron emission to
, whose half-life is . This isotope, along with 28Ne, have been used in the model of reactions in crust of neutron stars. 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 . Its atomic mass is . It decays to nitrogen-13 by electron capture, with a decay energy of . 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 , and a half-life of . It is produced through deuteron bombardment of nitrogen-14 using a cyclotron.
+ → + n
Oxygen-15 and nitrogen-13 are produced in air when gamma rays (for example from lightning) knock neutrons out of 16O and 14N:
+ γ → + n
+ γ → + n
decays to , 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 and floats by, carried by the wind.
See also
Dole effect
References
Oxygen
Oxygen