# Isotopes of oxygen

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

Radioactive isotopes ranging from 11O to 26O have also been characterized, all short-lived. The longest-lived radioisotope is 15O with a half-life of 122.24 seconds, while the shortest-lived isotope is 12O with a half-life of 580(30)×10−24 seconds (the half-life of the unbound 11O was not measured).

Main isotopes of oxygen (8O)
Iso­tope Decay
abun­dance half-life (t1/2) mode pro­duct
16O 99.76% stable
17O 0.04% stable
18O 0.20% stable
Standard atomic weight Ar, standard(O)
• [15.9990315.99977][1]
• Conventional: 15.999

## Stable isotopes

Late in a massive star's life, 16O concentrates in the O-shell, 17O in the H-shell and 18O in the He-shell

Naturally occurring oxygen is composed of three stable isotopes, 16O, 17O, and 18O, with 16O 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).

The relative and absolute abundance of 16O is high 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 made exclusively of hydrogen.[2] Most 16O is synthesized at the end of the helium fusion process in stars; the triple-alpha reaction creates 12C, which captures an additional 4He to make 16O. The neon burning process creates additional 16O.[2]

Both 17O and 18O are secondary isotopes, meaning that their nucleosynthesis requires seed nuclei. 17O is primarily made by the burning of hydrogen into helium during the CNO cycle, making it a common isotope in the hydrogen burning zones of stars.[2] Most 18O is produced when 14N (made abundant from CNO burning) captures a 4He nucleus, making 18O common in the helium-rich zones of stars.[2] Approximately a billion degrees Celsius is required for two oxygen nuclei to undergo nuclear fusion to form the heavier nucleus of sulfur.[3]

Measurements of the ratio of oxygen-18 to oxygen-16 are often used to interpret changes in paleoclimate. The isotopic composition of oxygen atoms in the Earth's atmosphere is 99.759% 16O, 0.037% 17O and 0.204% 18O.[4] Because water molecules containing the lighter isotope are slightly more likely to evaporate and fall as precipitation,[5] fresh water and polar ice on earth contains slightly less (0.1981%) of the heavy isotope 18O than air (0.204%) or seawater (0.1995%). This disparity allows analysis of temperature patterns via historic ice cores.

An atomic mass of 16 was assigned to oxygen prior to the definition of the unified atomic mass unit based upon 12C.[6] Since physicists referred to 16O only, while chemists meant the naturally-abundant mixture of isotopes, this led to slightly different mass scales between the two disciplines.

Thirteen radioisotopes have been characterized, with the most stable being 15O with a half-life of 122.24 s and 14O with a half-life of 70.606 s.[7] All of the remaining radioactive isotopes have half-lives that are less than 27 s and the majority of these have half-lives that are less than 83 milliseconds (ms).[7] For example, 24O has a half-life of 61 ms.[8] The most common decay mode for isotopes lighter than the stable isotopes is β+ decay (to nitrogen)[9][10][11] and the most common mode after is β decay (to fluorine).

### Oxygen-13

Oxygen-13 is an unstable isotope of oxygen. It consists of 8 protons and electrons, and 5 neutrons. It has a spin of 3/2-, and a half-life of 8.58 ms. Its atomic mass is 13.0248 Da. It decays to nitrogen-13 by electron capture, and has a decay energy of 17.765 MeV.[12] Its parent nuclide is fluorine-14.[13]

### Oxygen-15

Oxygen-15 is an isotope of oxygen, frequently used in positron emission tomography, or PET imaging. It can be used, amongst other applications, in water for PET myocardial perfusion imaging and for brain imaging.[14][15] It has 8 protons, 7 neutrons, and 8 electrons. The total atomic mass is 15.0030654 amu. It has a half-life of 122.24 seconds.[16] Oxygen-15 is synthesized through deuteron bombardment of nitrogen-14 using a cyclotron.[17]

Oxygen-15 and nitrogen-13 are produced in the atmosphere when gamma rays (for example from lightning) knock neutrons out of oxygen-16 and nitrogen-14:[18]

16O + γ → 15O + n
14N + γ → 13N + n

The oxygen-15 isotope decays with a half-life of about two minutes to nitrogen-15, 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 a half-life of two 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 15O and 13N floats by, carried by the wind.[19]

## List of isotopes

nuclide
symbol
Z(p) N(n)
isotopic mass (u)[20]

half-life decay mode(s)[21] daughter
isotope(s)[n 1]
nuclear
spin and
parity
representative
isotopic
composition
(mole fraction)
range of natural
variation
(mole fraction)
11O[22] 8 3 [~3.4 MeV] 2p 9C 3/2-, 5/2+
12O 8 4 12.034262(26) > 6.3(30)×10−21 s
[0.40(25) MeV]
2p (60.0%) 10C 0+
p (40.0%) 11N
13O 8 5 13.024815(10) 8.58(5) ms β+ (89.1%) 13N (3/2−)
β+, p (10.9%) 12C
14O 8 6 14.008596706(27) 70.620(13) s β+ 14N 0+
15O 8 7 15.0030656(5) 122.24(16) s β+ 15N 1/2−
16O[n 2] 8 8 15.99491461960(17) Stable 0+ 0.99757(16) 0.99738–0.99776
17O[n 3] 8 9 16.9991317566(7) Stable 5/2+ 3.8(1)×10−4 3.7×10−4–4.0×10−4
18O[n 2][n 4] 8 10 17.9991596128(8) Stable 0+ 2.05(14)×10−3 1.88×10−3–2.22×10−3
19O 8 11 19.0035780(28) 26.470(6) s β 19F 5/2+
20O 8 12 20.0040754(9) 13.51(5) s β 20F 0+
21O 8 13 21.008655(13) 3.42(10) s β 21F (5/2+)
22O 8 14 22.00997(6) 2.25(9) s β (78%) 22F 0+
β, n (22%) 21F
23O 8 15 23.01570(13) 97(8) ms β (93%) 23F 1/2+
β, n (7%) 22F
24O 8 16 24.01986(18) 77.4(45) ms β (57%) 24F 0+
β, n (43%) 23F
25O 8 17 25.02934(18) 5.18(0.35)×10−21 s n 24O 3/2+#
26O 8 18 26.03721(18) 4.2(3.3) ps 2n 24O
1. ^ Bold for stable isotopes
2. ^ a b The ratio between 16O and 18O is used to deduce ancient temperatures
3. ^ Can be used in NMR studies of metabolic pathways
4. ^ Can be used in studying certain metabolic pathways

### Notes

• The precision of the isotope abundances and atomic mass is limited through variations. The given ranges should be applicable to any normal terrestrial material.
• Values marked # are not purely derived from experimental data, but at least partly from systematic trends. Spins with weak assignment arguments are enclosed in parentheses.
• Uncertainties are given in concise form in parentheses after the corresponding last digits. Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, which use expanded uncertainties.
• Nuclide masses are given by IUPAP Commission on Symbols, Units, Nomenclature, Atomic Masses and Fundamental Constants (SUNAMCO).
• Isotope abundances are given by IUPAC Commission on Isotopic Abundances and Atomic Weights (CIAAW).

## Notes and references

1. ^ Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
2. ^ a b c d B. S. Meyer (September 19–21, 2005). "Nucleosynthesis and galactic chemical evolution of the isotopes of oxygen" (PDF). Proceedings of the NASA Cosmochemistry Program and the Lunar and Planetary Institute. Workgroup on Oxygen in the Earliest Solar System. Gatlinburg, Tennessee. 9022.
3. ^ Emsley 2001, p. 297.
4. ^ Cook 1968, p. 500.
5. ^ Dansgaard, W (1964). "Stable isotopes in precipitation" (PDF). Tellus. 16 (4): 436–468. Bibcode:1964TellA..16..436D. doi:10.1111/j.2153-3490.1964.tb00181.x.
6. ^ Parks & Mellor 1939, Chapter VI, Section 7.
7. ^ a b K. L. Barbalace. "Periodic Table of Elements: O - Oxygen". EnvironmentalChemistry.com. Retrieved 2007-12-17.
8. ^ Ekström, L. P.; Firestone, R. B. (28 February 1999). "Oxygen-24". WWW Table of Radioactive Isotopes. LUNDS Universitet, LBNL Isotopes Project. Archived from the original on 13 August 2009. Retrieved 2009-06-08.
9. ^ "NUDAT". Retrieved 2009-07-06.
10. ^ "NUDAT". Retrieved 2009-07-06.
11. ^ "NUDAT". Retrieved 2009-07-06.
12. ^ "Periodic Table of Elements: O - Oxygen". EnvironmentalChemistry.com. 1995-10-22. Retrieved 2014-12-02.
13. ^ "Periodic Table of Elements: F - Fluorine". EnvironmentalChemistry.com. 1995-10-22. Retrieved 2014-12-02.
14. ^ Rischpler, Christoph; Higuchi, Takahiro; Nekolla, Stephan G. (22 November 2014). "Current and Future Status of PET Myocardial Perfusion Tracers". Current Cardiovascular Imaging Reports. 8 (1): 333–343. doi:10.1007/s12410-014-9303-z. PMC 4333146. PMID 25234078.
15. ^ Kim, E. Edmund; Lee, Myung-Chul; Inoue, Tomio; Wong, Wai-Hoi (2012). Clinical PET and PET/CT: Principles and Applications. Springer. p. 182. ISBN 9781441908025.
16. ^ "oxygen 15 - definition of oxygen 15 by Medical dictionary". Medical-dictionary.thefreedictionary.com. Retrieved 2014-12-02.
17. ^ "Production of PET Radionuclides". Austin Hospital, Austin Health. Archived from the original on 15 January 2013. Retrieved 6 December 2012.
18. ^ Timmer, John (25 November 2017). "Lightning strikes leave behind a radioactive cloud". Ars Technica.
19. ^ Teruaki Enoto; et al. (Nov 23, 2017). "Photonuclear reactions triggered by lightning discharge". Nature. 551 (7681): 481–484. arXiv:1711.08044. Bibcode:2017Natur.551..481E. doi:10.1038/nature24630. PMID 29168803.
20. ^ Wang, Meng; Audi, Georges; Kondev, Filip G.; Huang, Wen Jian; Naimi, Sarah; Xu, Xing (2017), "The AME2016 atomic mass evaluation (II). Tables, graphs, and references" (PDF), Chinese Physics C, 41 (3): 030003–1—030003–442, doi:10.1088/1674-1137/41/3/030003
21. ^ Audi, Georges; Kondev, Filip G.; Wang, Meng; Huang, Wen Jia; Naimi, Sarah (2017), "The NUBASE2016 evaluation of nuclear properties" (PDF), Chinese Physics C, 41 (3): 030001–1—030001–138, Bibcode:2017ChPhC..41c0001A, doi:10.1088/1674-1137/41/3/030001
22. ^ Webb, T.B.; et al. (2019). "First Observation of Unbound 11O, the Mirror of the Halo Nucleus 11Li". Physical Review Letters. 122 (12): 122501–1—122501–7. arXiv:1812.08880. doi:10.1103/PhysRevLett.122.122501.

## References

For the table
For the prose
• 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.

Diademodon is an extinct genus of cynodonts. It was about 2 metres (6.6 ft) long. Although Diademodon is the most well accepted name for the genera to date, it was originally named Cynochampsa laniarius by Owen in 1860. The proposed name change occurred in 1982, where Grine defended the name proposed by Harry Seeley: Diademodon tetragonus and to be place in the group Therapsida, which was a group Owen had tiptoed around in his works on paleontology. Though Harry Govier Seeley had named Diademodon in 1894, which was after Owen had dubbed the genus Cynochampsa, Seeley had not realized the two were one and the same as the fossil that Owen named was claimed to have been found in a claystone nodule in the Renosterberg Mountains. A later paleontologist explored the same area where the fossil was claimed to have been found and declared no evidence of Cynognathus fossils.

Dole effect

The Dole effect, named after Malcolm Dole, describes an inequality in the ratio of the heavy isotope 18O (a "standard" oxygen atom with two additional neutrons) to the lighter 16O, measured in the atmosphere and seawater. This ratio is usually denoted δ18O.

It was noticed in 1935 that air contained more 18O than seawater; this was quantified in 1975 to 23.5‰, but later refined as 23.88‰ in 2005. The imbalance arises mainly as a result of respiration in plants and in animals. Due to thermodynamics of isotope reactions, respiration removes the lighter — hence more reactive — 16O in preference to 18O, increasing the relative amount of 18O in the atmosphere.

The inequality is balanced by photosynthesis. Photosynthesis emits oxygen with the same isotopic composition (i.e. the ratio between 18O and 16O) as the water (H2O) used in the reaction, which is independent of the atmospheric ratio. Thus when atmospheric 18O levels are high enough, photosynthesis will act as a reducing factor. However, as a complicating factor, the degree of fractionation (i.e. change in isotope ratio) occurring due to photosynthesis is not entirely dependent on the water drawn up by the plant, as fractionation can occur as a result of preferential evaporation of H216O - water bearing lighter oxygen isotopes, and other small but significant processes.

Equilibrium fractionation

Equilibrium isotope fractionation is the partial separation of isotopes between two or more substances in chemical equilibrium. Equilibrium fractionation is strongest at low temperatures, and (along with kinetic isotope effects) forms the basis of the most widely used isotopic paleothermometers (or climate proxies): D/H and 18O/16O records from ice cores, and 18O/16O records from calcium carbonate. It is thus important for the construction of geologic temperature records. Isotopic fractionations attributed to equilibrium processes have been observed in many elements, from hydrogen (D/H) to uranium (238U/235U). In general, the light elements (especially hydrogen, boron, carbon, nitrogen, oxygen and sulfur) are most susceptible to fractionation, and their isotopes tend to be separated to a greater degree than heavier elements.

Foraminifera

Foraminifera (; Latin for "hole bearers"; informally called "forams") are members of a phylum or class of amoeboid protists characterized by streaming granular ectoplasm for catching food and other uses; and commonly an external shell (called a "test") of diverse forms and materials. Tests of chitin (found in some simple genera, and Textularia in particular) are believed to be the most primitive type. Most foraminifera are marine, the majority of which live on or within the seafloor sediment (i.e., are benthic), while a smaller variety float in the water column at various depths (i.e., are planktonic). Fewer are known from freshwater or brackish conditions, and some very few (nonaquatic) soil species have been identified through molecular analysis of small subunit ribosomal DNA.Foraminifera typically produce a test, or shell, which can have either one or multiple chambers, some becoming quite elaborate in structure. These shells are commonly made of calcium carbonate (CaCO3) or agglutinated sediment particles. Over 50,000 species are recognized, both living (10,000) and fossil (40,000). They are usually less than 1 mm in size, but some are much larger, the largest species reaching up to 20 cm.In modern Scientific English, the term foraminifera is both singular and plural (irrespective of the word's Latin derivation), and is used to describe one or more specimens or taxa: its usage as singular or plural must be determined from context. Foraminifera is frequently used informally to describe the group, and in these cases is generally lowercase.

Harold Urey

Harold Clayton Urey (April 29, 1893 – January 5, 1981) was an American physical chemist whose pioneering work on isotopes earned him the Nobel Prize in Chemistry in 1934 for the discovery of deuterium. He played a significant role in the development of the atom bomb, as well as contributing to theories on the development of organic life from non-living matter.Born in Walkerton, Indiana, Urey studied thermodynamics under Gilbert N. Lewis at the University of California. After he received his PhD in 1923, he was awarded a fellowship by the American-Scandinavian Foundation to study at the Niels Bohr Institute in Copenhagen. He was a research associate at Johns Hopkins University before becoming an associate professor of Chemistry at Columbia University. In 1931, he began work with the separation of isotopes that resulted in the discovery of deuterium.

During World War II, Urey turned his knowledge of isotope separation to the problem of uranium enrichment. He headed the group located at Columbia University that developed isotope separation using gaseous diffusion. The method was successfully developed, becoming the sole method used in the early post-war period. After the war, Urey became professor of chemistry at the Institute for Nuclear Studies, and later Ryerson professor of chemistry at the University of Chicago.

Urey speculated that the early terrestrial atmosphere was composed of ammonia, methane, and hydrogen. One of his Chicago graduate students was Stanley L. Miller, who showed in the Miller–Urey experiment that, if such a mixture were exposed to electric sparks and water, it can interact to produce amino acids, commonly considered the building blocks of life. Work with isotopes of oxygen led to pioneering the new field of paleoclimatic research. In 1958, he accepted a post as a professor at large at the new University of California, San Diego (UCSD), where he helped create the science faculty. He was one of the founding members of UCSD's school of chemistry, which was created in 1960. He became increasingly interested in space science, and when Apollo 11 returned moon rock samples from the moon, Urey examined them at the Lunar Receiving Laboratory. Lunar astronaut Harrison Schmitt said that Urey approached him as a volunteer for a one-way mission to the Moon, stating "I will go, and I don't care if I don't come back."

Isotopes of nitrogen

Natural nitrogen (7N) consists of two stable isotopes, nitrogen-14, which makes up the vast majority of naturally occurring nitrogen, and nitrogen-15, which is less common. Fourteen radioactive isotopes (radioisotopes) have also been found so far, with atomic masses ranging from 10 to 25, and one nuclear isomer, 11mN. All of these radioisotopes are short-lived, with the longest-lived one being nitrogen-13 with a half-life of 9.965 minutes. All of the others have half-lives below 7.15 seconds, with most of these being below five-eighths of a second. Most of the isotopes with atomic mass numbers below 14 decay to isotopes of carbon, while most of the isotopes with masses above 15 decay to isotopes of oxygen. The shortest-lived known isotope is nitrogen-10, with a half-life of about 200 yoctoseconds.

Jacob Bigeleisen

Jacob Bigeleisen (pronounced BEEG-a-lie-zen; May 2, 1919 – August 7, 2010) was an American chemist who worked on the Manhattan Project on techniques to extract uranium-235 from uranium ore, an isotope that can sustain nuclear fission and would be used in developing an atomic bomb but that is less than 1% of naturally occurring uranium. While the method of using photochemistry that Bigeleisen used as an approach was not successful in isolating useful quantities of uranium-235 for the war effort, it did lead to the development of isotope chemistry, which takes advantage of the ways that different isotopes of an element interact to form chemical bonds.

Mass-independent fractionation

Mass-independent isotope fractionation or Non-mass-dependent fractionation (NMD), refers to any chemical or physical process that acts to separate isotopes, where the amount of separation does not scale in proportion with the difference in the masses of the isotopes. Most isotopic fractionations (including typical kinetic fractionations and equilibrium fractionations) are caused by the effects of the mass of an isotope on atomic or molecular velocities, diffusivities or bond strengths. Mass-independent fractionation processes are less common, occurring mainly in photochemical and spin-forbidden reactions. Observation of mass-independently fractionated materials can therefore be used to trace these types of reactions in nature and in laboratory experiments.

Oxygen

Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a highly reactive nonmetal, and an oxidizing agent that readily forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after hydrogen and helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up almost half of the Earth's crust.

Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids, carbohydrates, and fats, as do the major constituent inorganic compounds of animal shells, teeth, and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form (allotrope) of oxygen, ozone (O3), strongly absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of smog and thus a pollutant.

Oxygen was isolated by Michael Sendivogius before 1604, but it is commonly believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, and Joseph Priestley in Wiltshire, in 1774. Priority is often given for Priestley because his work was published first. Priestley, however, called oxygen "dephlogisticated air", and did not recognize it as a chemical element. The name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and correctly characterized the role it plays in combustion.

Common uses of oxygen include production of steel, plastics and textiles, brazing, welding and cutting of steels and other metals, rocket propellant, oxygen therapy, and life support systems in aircraft, submarines, spaceflight and diving.

Oxygen-16

Oxygen-16 (16O) is a stable isotope of oxygen, having 8 neutrons and 8 protons in its nucleus. It has a mass of 15.99491461956 u. Oxygen-16 is the most abundant isotope of oxygen and accounts for 99.762% of oxygen's natural abundance. The relative and absolute abundance of 16O are high because it is a principal product of stellar evolution and because it is a primordial isotope, meaning it can be made by stars that were initially made exclusively of hydrogen. Most 16O is synthesized at the end of the helium fusion process in stars; the triple-alpha process creates 12C, which captures an additional 4He to make 16O. The neon-burning process creates additional 16O.

Oxygen-17

Oxygen-17 is a low-abundant, natural, stable isotope of oxygen (0.0373% in seawater; approximately twice as abundant as deuterium).

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

Water used as nuclear reactor coolant is subjected to intense neutron flux. Natural water starts out with 373 ppm of O-17; heavy water starts out incidentally enriched to about 550 ppm of oxygen-17. The neutron flux slowly converts O-16 in the cooling water to O-17 by neutron capture, increasing its concentration. The neutron flux slowly converts O-17 in the cooling water to carbon-14, an undesirable product that escapes to the environment. 17O (n,alpha) → 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 C-14 production.

Oxygen-18

Oxygen-18 (18O) is a natural, stable isotope of oxygen and one of the environmental isotopes.

18O is an important precursor for the production of fluorodeoxyglucose (FDG) used in positron emission tomography (PET). Generally, in the radiopharmaceutical industry, enriched water (H218O) is bombarded with hydrogen ions in either a cyclotron or linear accelerator, creating fluorine-18. This is then synthesized into FDG and injected into a patient. It can also be used to make an extremely heavy version of water when combined with tritium (hydrogen-3): 3H218O or T218O. This compound has a density almost 30% greater than that of natural water

Oxygen evolution

Oxygen evolution is the process of generating molecular oxygen (O2) by a chemical reaction, usually from water. Oxygen evolution from water is effected by oxygenic photosynthesis, electrolysis of water, and thermal decomposition of various oxides. The biological process supports aerobic life. When relatively pure oxygen is required industrially, it is isolated by distillation of liquified air.

Oxygen isotope ratio cycle

Oxygen isotope ratio cycles are cyclical variations in the ratio of the abundance of oxygen with an atomic mass of 18 to the abundance of oxygen with an atomic mass of 16 present in some substances, such as polar ice or calcite in ocean core samples, measured with the isotope fractionation. The ratio is linked to water temperature of ancient oceans, which in turn reflects ancient climates. Cycles in the ratio mirror climate changes in geologic history.

Pleistocene

The Pleistocene ( , often colloquially referred to as the Ice Age) is the geological epoch which lasted from about 2,588,000 to 11,700 years ago, spanning the world's most recent period of repeated glaciations. The end of the Pleistocene corresponds with the end of the last glacial period and also with the end of the Paleolithic age used in archaeology.

The Pleistocene is the first epoch of the Quaternary Period or sixth epoch of the Cenozoic Era. In the ICS timescale, the Pleistocene is divided into four stages or ages, the Gelasian, Calabrian, Middle Pleistocene (unofficially the 'Chibanian') and Upper Pleistocene (unofficially the 'Tarantian'). In addition to this international subdivision, various regional subdivisions are often used.

Before a change finally confirmed in 2009 by the International Union of Geological Sciences, the time boundary between the Pleistocene and the preceding Pliocene was regarded as being at 1.806 million years Before Present (BP), as opposed to the currently accepted 2.588 million years BP: publications from the preceding years may use either definition of the period.

Speleothem

Speleothems ( ; Ancient Greek: "cave deposit"), commonly known as cave formations, are secondary mineral deposits formed in a cave. Speleothems typically form in limestone or dolostone solutional caves. The term "speleothem" as first introduced by Moore (1952), is derived from the Greek words spēlaion "cave" + théma "deposit". The definition of "speleothem" in most publications, specifically excludes secondary mineral deposits in mines, tunnels and on man-made structures. Hill and Forti more concisely defined "secondary minerals" which create speleothems in caves as;

A "secondary" mineral is one which is derived by a physicochemical reaction from a primary mineral in bedrock or detritus, and/or deposited because of a unique set of conditions in a cave; i.e., the cave environment has influenced the mineral's deposition.

Toshiko Mayeda

Toshiko K. Mayeda (née Kuki) (1923–2004) was a Japanese American chemist who worked at the Enrico Fermi Institute in the University of Chicago. She worked on climate science and meteorites from 1958 to 2004.

Δ18O

In geochemistry, paleoclimatology and paleoceanography δ18O or delta-O-18 is a measure of the ratio of stable isotopes oxygen-18 (18O) and oxygen-16 (16O). It is commonly used as a measure of the temperature of precipitation, as a measure of groundwater/mineral interactions, and as an indicator of processes that show isotopic fractionation, like methanogenesis. In paleosciences, 18O:16O data from corals, foraminifera and ice cores are used as a proxy for temperature.

The definition is, in "per mil" (‰, parts per thousand):

${\displaystyle \delta {\ce {^{18}O}}=\left({\frac {\left({\frac {{\ce {^{18}O}}}{{\ce {^{16}O}}}}\right)_{\mathrm {sample} }}{\left({\frac {{\ce {^{18}O}}}{{\ce {^{16}O}}}}\right)_{\mathrm {standard} }}}-1\right)\times 1000}$

where the standard has a known isotopic composition, such as Vienna Standard Mean Ocean Water (VSMOW). The fractionation can arise from kinetic, equilibrium, or mass-independent fractionation.

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