| |||||||||||||||||||||||||||||||
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. 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]
The atom is the basic particle of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons but a different number of neutrons are called isotopes of the same element.
Nuclear physics is the field of physics that studies atomic nuclei and their constituents and interactions, in addition to the study of other forms of nuclear matter.
Nucleosynthesis is the process that creates new atomic nuclei from pre-existing nucleons and nuclei. According to current theories, the first nuclei were formed a few minutes after the Big Bang, through nuclear reactions in a process called Big Bang nucleosynthesis. After about 20 minutes, the universe had expanded and cooled to a point at which these high-energy collisions among nucleons ended, so only the fastest and simplest reactions occurred, leaving our universe containing hydrogen and helium. The rest is traces of other elements such as lithium and the hydrogen isotope deuterium. Nucleosynthesis in stars and their explosions later produced the variety of elements and isotopes that we have today, in a process called cosmic chemical evolution. The amounts of total mass in elements heavier than hydrogen and helium remains small, so that the universe still has approximately the same composition.
The abundance of the chemical elements is a measure of the occurrence of the chemical elements relative to all other elements in a given environment. Abundance is measured in one of three ways: by mass fraction, by mole fraction, or by volume fraction. Volume fraction is a common abundance measure in mixed gases such as planetary atmospheres, and is similar in value to molecular mole fraction for gas mixtures at relatively low densities and pressures, and ideal gas mixtures. Most abundance values in this article are given as mass fractions.
A radioactive tracer, radiotracer, or radioactive label is a synthetic derivative of a natural compound in which one or more atoms have been replaced by a radionuclide. By virtue of its radioactive decay, it can be used to explore the mechanism of chemical reactions by tracing the path that the radioisotope follows from reactants to products. Radiolabeling or radiotracing is thus the radioactive form of isotopic labeling. In biological contexts, experiments that use radioisotope tracers are sometimes called radioisotope feeding experiments.
In nuclear physics, a magic number is a number of nucleons such that they are arranged into complete shells within the atomic nucleus. As a result, atomic nuclei with a 'magic' number of protons or neutrons are much more stable than other nuclei. The seven most widely recognized magic numbers as of 2019 are 2, 8, 20, 28, 50, 82, and 126.
Neutron activation is the process in which neutron radiation induces radioactivity in materials, and occurs when atomic nuclei capture free neutrons, becoming heavier and entering excited states. The excited nucleus decays immediately by emitting gamma rays, or particles such as beta particles, alpha particles, fission products, and neutrons. Thus, the process of neutron capture, even after any intermediate decay, often results in the formation of an unstable activation product. Such radioactive nuclei can exhibit half-lives ranging from small fractions of a second to many years.
Hydrogen (1H) has three naturally occurring isotopes, sometimes denoted 1
H
, 2
H
, and 3
H
. 1
H
and 2
H
are stable, while 3
H
has a half-life of 12.32(2) years. Heavier isotopes also exist, all of which are synthetic and have a half-life of less than one zeptosecond (10−21 s). Of these, 5
H
is the least stable, while 7
H
is the most.
Fluorine (9F) has 18 known isotopes ranging from 13
F
to 31
F
and two isomers. Only fluorine-19 is stable and naturally occurring in more than trace quantities; therefore, fluorine is a monoisotopic and mononuclidic element.
Technetium (43Tc) is one of the two elements with Z < 83 that have no stable isotopes; the other such element is promethium. It is primarily artificial, with only trace quantities existing in nature produced by spontaneous fission or neutron capture by molybdenum. The first isotopes to be synthesized were 97Tc and 99Tc in 1936, the first artificial element to be produced. The most stable radioisotopes are 97Tc, 98Tc, and 99Tc.
Naturally occurring zirconium (40Zr) is composed of four stable isotopes (of which one may in the future be found radioactive), and one very long-lived radioisotope (96Zr), a primordial nuclide that decays via double beta decay with an observed half-life of 2.0×1019 years; it can also undergo single beta decay, which is not yet observed, but the theoretically predicted value of t1/2 is 2.4×1020 years. The second most stable radioisotope is 93Zr, which has a half-life of 1.53 million years. Thirty other radioisotopes have been observed. All have half-lives less than a day except for 95Zr (64.02 days), 88Zr (83.4 days), and 89Zr (78.41 hours). The primary decay mode is electron capture for isotopes lighter than 92Zr, and the primary mode for heavier isotopes is beta decay.
Bromine (35Br) has two stable isotopes, 79Br and 81Br, and 32 known radioisotopes, the most stable of which is 77Br, with a half-life of 57.036 hours.
Natural nitrogen (7N) consists of two stable isotopes: the vast majority (99.6%) of naturally occurring nitrogen is nitrogen-14, with the remainder being nitrogen-15. Thirteen radioisotopes are also known, with atomic masses ranging from 9 to 23, along with three nuclear isomers. All of these radioisotopes are short-lived, the longest-lived being nitrogen-13 with a half-life of 9.965(4) min. All of the others have half-lives below 7.15 seconds, with most of these being below 620 milliseconds. 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 143(36) yoctoseconds, though the half-life of nitrogen-9 has not been measured exactly.
Carbon (6C) has 15 known isotopes, from 8
C
to 22
C
, of which 12
C
and 13
C
are stable. The longest-lived radioisotope is 14
C
, with a half-life of 5.70(3)×103 years. This is also the only carbon radioisotope found in nature, as trace quantities are formed cosmogenically by the reaction 14
N
+
n
→ 14
C
+ 1
H
. The most stable artificial radioisotope is 11
C
, which has a half-life of 20.3402(53) min. All other radioisotopes have half-lives under 20 seconds, most less than 200 milliseconds. The least stable isotope is 8
C
, with a half-life of 3.5(1.4)×10−21 s. Light isotopes tend to decay into isotopes of boron and heavy ones tend to decay into isotopes of nitrogen.
Although there are nine known isotopes of helium (2He), only helium-3 and helium-4 are stable. All radioisotopes are short-lived, the longest-lived being 6
He
with a half-life of 806.92(24) milliseconds. The least stable is 10
He
, with a half-life of 260(40) yoctoseconds, although it is possible that 2
He
may have an even shorter half-life.
Nuclear binding energy in experimental physics is the minimum energy that is required to disassemble the nucleus of an atom into its constituent protons and neutrons, known collectively as nucleons. The binding energy for stable nuclei is always a positive number, as the nucleus must gain energy for the nucleons to move apart from each other. Nucleons are attracted to each other by the strong nuclear force. In theoretical nuclear physics, the nuclear binding energy is considered a negative number. In this context it represents the energy of the nucleus relative to the energy of the constituent nucleons when they are infinitely far apart. Both the experimental and theoretical views are equivalent, with slightly different emphasis on what the binding energy means.
A table or chart of nuclides is a two-dimensional graph of isotopes of the elements, in which one axis represents the number of neutrons and the other represents the number of protons in the atomic nucleus. Each point plotted on the graph thus represents a nuclide of a known or hypothetical chemical element. This system of ordering nuclides can offer a greater insight into the characteristics of isotopes than the better-known periodic table, which shows only elements and not their isotopes. The chart of the nuclides is also known as the Segrè chart, after the Italian physicist Emilio Segrè.
Isotopes are distinct nuclear species of the same chemical element. They have the same atomic number and position in the periodic table, but differ in nucleon numbers due to different numbers of neutrons in their nuclei. While all isotopes of a given element have almost the same chemical properties, they have different atomic masses and physical properties.
Nuclear transmutation is the conversion of one chemical element or an isotope into another chemical element. Nuclear transmutation occurs in any process where the number of protons or neutrons in the nucleus of an atom is changed.
In nuclear physics, properties of a nucleus depend on evenness or oddness of its atomic number Z, neutron number N and, consequently, of their sum, the mass number A. Most importantly, oddness of both Z and N tends to lower the nuclear binding energy, making odd nuclei generally less stable. This effect is not only experimentally observed, but is included in the semi-empirical mass formula and explained by some other nuclear models, such as the nuclear shell model. This difference of nuclear binding energy between neighbouring nuclei, especially of odd-A isobars, has important consequences for beta decay.
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