| ||||||||||||||||||||||||||
| Standard atomic weight Ar°(Na) | ||||||||||||||||||||||||||
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
There are 20 isotopes of sodium (11Na), ranging from 17
Na to 39
Na (except for the still-unknown 36Na and 38Na), [4] and two isomers (22m
Na and 24m
Na). 23
Na is the only stable (and the only primordial) isotope. It is considered a monoisotopic element and it has a standard atomic weight of 22.98976928(2). Sodium has two radioactive cosmogenic isotopes (22
Na, with a half-life of 2.6019(6) years; [nb 1] and 24
Na , with a half-life of 14.9560(15) h). With the exception of those two isotopes, all other isotopes have half-lives under a minute, most under a second. The shortest-lived is the unbound 18
Na, with a half-life of 1.3(4)×10−21 seconds (although the half-life of the similarly unbound 17Na is not measured).
Acute neutron radiation exposure (e.g., from a nuclear criticality accident) converts some of the stable 23
Na (in the form of Na+ ion) in human blood plasma to 24
Na. By measuring the concentration of this isotope, the neutron radiation dosage to the victim can be computed.
22
Na is a positron-emitting isotope with a remarkably long half-life. It is used to create test-objects and point-sources for positron emission tomography.
| Nuclide [n 1] | Z | N | Isotopic mass (Da) [5] [n 2] [n 3] | Half-life [1] [n 4] | Decay mode [1] [n 5] | Daughter isotope [n 6] | Spin and parity [1] [n 7] [n 4] | Natural abundance (mole fraction) | |||||||||||
|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|---|
| Excitation energy | Normal proportion [1] | Range of variation | |||||||||||||||||
| 17 Na | 11 | 6 | 17.037270(60) | p | 16 Ne | (1/2+) | |||||||||||||
| 18 Na | 11 | 7 | 18.02688(10) | 1.3(4) zs | p=? [n 8] | 17 Ne | 1−# | ||||||||||||
| 19 Na | 11 | 8 | 19.013880(11) | > 1 as | p | 18 Ne | (5/2+) | ||||||||||||
| 20 Na | 11 | 9 | 20.0073543(12) | 447.9(2.3) ms | β+ (75.0(4)%) | 20 Ne | 2+ | ||||||||||||
| β+α (25.0(4)%) | 16 O | ||||||||||||||||||
| 21 Na | 11 | 10 | 20.99765446(5) | 22.4550(54) s | β+ | 21 Ne | 3/2+ | ||||||||||||
| 22 Na | 11 | 11 | 21.99443742(18) | 2.6019(6) y [nb 1] | β+ (90.57(8)%) | 22 Ne | 3+ | Trace [n 9] | |||||||||||
| ε (9.43(6)%) | 22 Ne | ||||||||||||||||||
| 22m1 Na | 583.05(10) keV | 243(2) ns | IT | 22 Na | 1+ | ||||||||||||||
| 22m2 Na | 657.00(14) keV | 19.6(7) ps | IT | 22 Na | 0+ | ||||||||||||||
| 23 Na | 11 | 12 | 22.9897692820(19) | Stable | 3/2+ | 1 | |||||||||||||
| 24 Na | 11 | 13 | 23.990963012(18) | 14.9560(15) h | β− | 24 Mg | 4+ | Trace [n 9] | |||||||||||
| 24m Na | 472.2074(8) keV | 20.18(10) ms | IT (99.95%) | 24 Na | 1+ | ||||||||||||||
| β− (0.05%) | 24 Mg | ||||||||||||||||||
| 25 Na | 11 | 14 | 24.9899540(13) | 59.1(6) s | β− | 25 Mg | 5/2+ | ||||||||||||
| 26 Na | 11 | 15 | 25.992635(4) | 1.07128(25) s | β− | 26 Mg | 3+ | ||||||||||||
| 26m Na | 82.4(4) keV | 4.35(16) μs | IT | 26 Na | 1+ | ||||||||||||||
| 27 Na | 11 | 16 | 26.994076(4) | 301(6) ms | β− (99.902(24)%) | 27 Mg | 5/2+ | ||||||||||||
| β−n (0.098(24)%) | 26 Mg | ||||||||||||||||||
| 28 Na | 11 | 17 | 27.998939(11) | 33.1(1.3) ms | β− (99.42(12)%) | 28 Mg | 1+ | ||||||||||||
| β−n (0.58(12)%) | 27 Mg | ||||||||||||||||||
| 29 Na | 11 | 18 | 29.002877(8) | 43.2(4) ms | β− (78%) | 29 Mg | 3/2+ | ||||||||||||
| β−n (22(3)%) | 28 Mg | ||||||||||||||||||
| β−2n ? [n 10] | 27 Mg ? | ||||||||||||||||||
| 30 Na | 11 | 19 | 30.009098(5) | 45.9(7) ms | β− (70.2(2.2)%) | 30 Mg | 2+ | ||||||||||||
| β−n (28.6(2.2)%) | 29 Mg | ||||||||||||||||||
| β−2n (1.24(19)%) | 28 Mg | ||||||||||||||||||
| β−α (5.5(2)%×10−5) | 26 Ne | ||||||||||||||||||
| 31 Na | 11 | 20 | 31.013147(15) | 16.8(3) ms | β− (> 63.2(3.5)%) | 31 Mg | 3/2+ | ||||||||||||
| β−n (36.0(3.5)%) | 30 Mg | ||||||||||||||||||
| β−2n (0.73(9)%) | 29 Mg | ||||||||||||||||||
| β−3n (< 0.05%) | 28 Mg | ||||||||||||||||||
| 32 Na | 11 | 21 | 32.020010(40) | 12.9(3) ms | β− (66.4(6.2)%) | 32 Mg | (3−) | ||||||||||||
| β−n (26(6)%) | 31 Mg | ||||||||||||||||||
| β−2n (7.6(1.5)%) | 30 Mg | ||||||||||||||||||
| 33 Na | 11 | 22 | 33.02553(48) | 8.2(4) ms | β−n (47(6)%) | 32 Mg | (3/2+) | ||||||||||||
| β− (40.0(6.7)%) | 33 Mg | ||||||||||||||||||
| β−2n (13(3)%) | 31 Mg | ||||||||||||||||||
| 34 Na | 11 | 23 | 34.03401(64) | 5.5(1.0) ms | β−2n (~50%) | 32 Mg | 1+ | ||||||||||||
| β− (~35%) | 34 Mg | ||||||||||||||||||
| β−n (~15%) | 33 Mg | ||||||||||||||||||
| 35 Na | 11 | 24 | 35.04061(72)# | 1.5(5) ms | β− | 35 Mg | 3/2+# | ||||||||||||
| β−n ? [n 10] | 34 Mg ? | ||||||||||||||||||
| β−2n ? [n 10] | 33 Mg ? | ||||||||||||||||||
| 37 Na | 11 | 26 | 37.05704(74)# | 1# ms [> 1.5 μs] | β− ? [n 10] | 37 Mg ? | 3/2+# | ||||||||||||
| β−n ? [n 10] | 36 Mg ? | ||||||||||||||||||
| β−2n ? [n 10] | 35 Mg ? | ||||||||||||||||||
| 39 Na [4] | 11 | 28 | 39.07512(80)# | 1# ms [> 400 ns] | β− ? [n 10] | 39 Mg ? | 3/2+# | ||||||||||||
| β−n ? [n 10] | 38 Mg ? | ||||||||||||||||||
| β−2n ? [n 10] | 37 Mg ? | ||||||||||||||||||
| This table header & footer: | |||||||||||||||||||
| IT: | Isomeric transition |
| n: | Neutron emission |
| p: | Proton emission |
Sodium-22 is a radioactive isotope of sodium, undergoing positron emission to 22
Ne with a half-life of 2.6019(6) years. 22
Na is being investigated as an efficient generator of "cold positrons" (antimatter) to produce muons for catalyzing fusion of deuterium.[ citation needed ] It is also commonly used as a positron source in positron annihilation spectroscopy. [6]
Sodium-23 is an isotope of sodium with an atomic mass of 22.98976928. It is the only stable isotope of sodium, and because of its abundance, it has been used for nuclear magnetic resonance in various research fields, including materials science and battery research. [7] Sodium-23 relaxation has applications in studying cation-biomolecule interactions, intracellular and extracellular sodium, ion transport in batteries, and quantum information processing. [8]
Sodium-24 is radioactive and can be created from common sodium-23 by neutron activation. With a half-life of 14.9560(15) h, 24
Na decays to 24
Mg by emission of an electron and two gamma rays. [9] [10]
Exposure of the human body to intense neutron radiation creates 24
Na in the blood plasma. Measurements of its quantity can be done to determine the absorbed radiation dose of a patient. [10] This can be used to determine the type of medical treatment required.
When sodium is used as coolant in fast breeder reactors, 24
Na is created, which makes the coolant radioactive. When the 24
Na decays, it causes a buildup of magnesium in the coolant. Since the half-life is short, the 24
Na portion of the coolant ceases to be radioactive within a few days after removal from the reactor. Leakage of the hot sodium from the primary loop may cause radioactive fires, [11] as it can ignite in contact with air (and explodes in contact with water). For this reason the primary cooling loop is within a containment vessel.
Sodium has been proposed as a casing for a salted bomb, as it would convert to 24
Na and produce intense gamma-ray emissions for a few days. [12] [13]
A radionuclide (radioactive nuclide, radioisotope or radioactive isotope) is a nuclide that has excess numbers of either neutrons or protons, giving it excess nuclear energy, and making it unstable. This excess energy can be used in one of three ways: emitted from the nucleus as gamma radiation; transferred to one of its electrons to release it as a conversion electron; or used to create and emit a new particle (alpha particle or beta particle) from the nucleus. During those processes, the radionuclide is said to undergo radioactive decay. These emissions are considered ionizing radiation because they are energetic enough to liberate an electron from another atom. The radioactive decay can produce a stable nuclide or will sometimes produce a new unstable radionuclide which may undergo further decay. Radioactive decay is a random process at the level of single atoms: it is impossible to predict when one particular atom will decay. However, for a collection of atoms of a single nuclide the decay rate, and thus the half-life (t1/2) for that collection, can be calculated from their measured decay constants. The range of the half-lives of radioactive atoms has no known limits and spans a time range of over 55 orders of magnitude.
A nuclide is a class of atoms characterized by their number of protons, Z, their number of neutrons, N, and their nuclear energy state.
Radioactive decay is the process by which an unstable atomic nucleus loses energy by radiation. A material containing unstable nuclei is considered radioactive. Three of the most common types of decay are alpha, beta, and gamma decay. The weak force is the mechanism that is responsible for beta decay, while the other two are governed by the electromagnetism and nuclear force.
Positron emission, beta plus decay, or β+ decay is a subtype of radioactive decay called beta decay, in which a proton inside a radionuclide nucleus is converted into a neutron while releasing a positron and an electron neutrino. Positron emission is mediated by the weak force. The positron is a type of beta particle (β+), the other beta particle being the electron (β−) emitted from the β− decay of a nucleus.
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.
Naturally occurring erbium (68Er) is composed of 6 stable isotopes, with 166Er being the most abundant. 39 radioisotopes have been characterized with between 74 and 112 neutrons, or 142 to 180 nucleons, with the most stable being 169Er with a half-life of 9.4 days, 172Er with a half-life of 49.3 hours, 160Er with a half-life of 28.58 hours, 165Er with a half-life of 10.36 hours, and 171Er with a half-life of 7.516 hours. All of the remaining radioactive isotopes have half-lives that are less than 3.5 hours, and the majority of these have half-lives that are less than 4 minutes. This element also has numerous meta states, with the most stable being 167mEr.
There are 37 known isotopes of iodine (53I) from 108I to 144I; all undergo radioactive decay except 127I, which is stable. Iodine is thus a monoisotopic element.
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.
Germanium (32Ge) has five naturally occurring isotopes, 70Ge, 72Ge, 73Ge, 74Ge, and 76Ge. Of these, 76Ge is very slightly radioactive, decaying by double beta decay with a half-life of 1.78 × 1021 years (130 billion times the age of the universe).
Calcium (20Ca) has 26 known isotopes, ranging from 35Ca to 60Ca. There are five stable isotopes, plus one isotope (48Ca) with such a long half-life that for all practical purposes it can be considered stable. The most abundant isotope, 40Ca, as well as the rare 46Ca, are theoretically unstable on energetic grounds, but their decay has not been observed. Calcium also has a cosmogenic isotope, radioactive 41Ca, which has a half-life of 99,400 years. Unlike cosmogenic isotopes that are produced in the atmosphere, 41Ca is produced by neutron activation of 40Ca. Most of its production is in the upper metre of the soil column, where the cosmogenic neutron flux is still sufficiently strong. 41Ca has received much attention in stellar studies because it decays to 41K, a critical indicator of solar system anomalies. The most stable artificial radioisotopes are 45Ca with a half-life of 163 days and 47Ca with a half-life of 4.5 days. All other calcium isotopes have half-lives measured in minutes or less.
Potassium has 26 known isotopes from 31
K to 57
K, with the exception of still-unknown 32
K, as well as an unconfirmed report of 59
K. Three of those isotopes occur naturally: the two stable forms 39
K (93.3%) and 41
K (6.7%), and a very long-lived radioisotope 40
K (0.012%)
Argon (18Ar) has 26 known isotopes, from 29Ar to 54Ar and 1 isomer (32mAr), of which three are stable. On the Earth, 40Ar makes up 99.6% of natural argon. The longest-lived radioactive isotopes are 39Ar with a half-life of 268 years, 42Ar with a half-life of 32.9 years, and 37Ar with a half-life of 35.04 days. All other isotopes have half-lives of less than two hours, and most less than one minute. The least stable is 29Ar with a half-life of approximately 4×10−20 seconds.
Chlorine (17Cl) has 25 isotopes, ranging from 28Cl to 52Cl, and two isomers, 34mCl and 38mCl. There are two stable isotopes, 35Cl (75.8%) and 37Cl (24.2%), giving chlorine a standard atomic weight of 35.45. The longest-lived radioactive isotope is 36Cl, which has a half-life of 301,000 years. All other isotopes have half-lives under 1 hour, many less than one second. The shortest-lived are proton-unbound 29Cl and 30Cl, with half-lives less than 10 picoseconds and 30 nanoseconds, respectively; the half-life of 28Cl is unknown.
Sulfur (16S) has 23 known isotopes with mass numbers ranging from 27 to 49, four of which are stable: 32S (95.02%), 33S (0.75%), 34S (4.21%), and 36S (0.02%). The preponderance of sulfur-32 is explained by its production from carbon-12 plus successive fusion capture of five helium-4 nuclei, in the so-called alpha process of exploding type II supernovas.
Aluminium or aluminum (13Al) has 22 known isotopes from 22Al to 43Al and 4 known isomers. Only 27Al (stable isotope) and 26Al (radioactive isotope, t1/2 = 7.2×105 y) occur naturally, however 27Al comprises nearly all natural aluminium. Other than 26Al, all radioisotopes have half-lives under 7 minutes, most under a second. The standard atomic weight is 26.9815385(7). 26Al is produced from argon in the atmosphere by spallation caused by cosmic-ray protons. Aluminium isotopes have found practical application in dating marine sediments, manganese nodules, glacial ice, quartz in rock exposures, and meteorites. The ratio of 26Al to 10Be has been used to study the role of sediment transport, deposition, and storage, as well as burial times, and erosion, on 105 to 106 year time scales. 26Al has also played a significant role in the study of meteorites.
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 14 known isotopes, from 8
C
to 20
C
as well as 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.
Californium (98Cf) is an artificial element, and thus a standard atomic weight cannot be given. Like all artificial elements, it has no stable isotopes. The first isotope to be synthesized was 245Cf in 1950. There are 20 known radioisotopes ranging from 237Cf to 256Cf and one nuclear isomer, 249mCf. The longest-lived isotope is 251Cf with a half-life of 898 years.
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