Isotopes of oxygen

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]

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] Temperatures on the order of 10⁹ kelvins are 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 nuclear 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]

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:^[23]

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-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]

• Dole effect

• 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.