Natural abundance

In physics, natural abundance (NA) refers to the abundance of isotopes of a chemical element as naturally found on a planet. The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet, and even from place to place on the Earth, but remains relatively constant in time (on a short-term scale).

As an example, uranium has three naturally occurring isotopes: 238U, 235U and 234U. Their respective natural mole-fraction abundances are 99.2739–99.2752%, 0.7198–0.7202%, and 0.0050–0.0059%.[1] For example, if 100,000 uranium atoms were analyzed, one would expect to find approximately 99,274 238U atoms, approximately 720 235U atoms, and very few (most likely 5 or 6) 234U atoms. This is because 238U is much more stable than 235U or 234U, as the half-life of each isotope reveals: 4.468 × 109 years for 238U compared with 7.038 × 108 years for 235U and 245,500 years for 234U.

Exactly because the different uranium isotopes have different half-lives, when the Earth was younger, the isotopic composition of uranium was different. As an example, 1.7×109 years ago the NA of 235U was 3.1% compared with today's 0.7%, and for that reason a natural nuclear fission reactor was able to form, something that cannot happen today.

However, the natural abundance of a given isotope is also affected by the probability of its creation in nucleosynthesis (as in the case of samarium; radioactive 147Sm and 148Sm are much more abundant than stable 144Sm) and by production of a given isotope as a daughter of natural radioactive isotopes (as in the case of radiogenic isotopes of lead).

Relative abundance of elements

Deviations from natural abundance

It is now known from study of the sun and primitive meteorites that the solar system was initially almost homogeneous in isotopic composition. Deviations from the (evolving) galactic average, locally sampled around the time that the sun's nuclear burning began, can generally be accounted for by mass fractionation (see the article on mass-independent fractionation) plus a limited number of nuclear decay and transmutation processes.[2] There is also evidence for injection of short-lived (now extinct) isotopes from a nearby supernova explosion that may have triggered solar nebula collapse.[3] Hence deviations from natural abundance on earth are often measured in parts per thousand (per mille or ‰‰) because they are less than one percent (%).

The single exception to this lies with the presolar grains found in primitive meteorites. These bypassed the homogenization, and often carry the nuclear signature of specific nucleosynthesis processes in which their elements were made.[4] In these materials, deviations from "natural abundance" are sometimes measured in factors of 100.

Natural abundance of some elements

The next table gives the isotope distributions for some elements. Some elements like phosphorus and fluorine only exist as a single isotope, with a natural abundance of 100%.

Natural abundance of some elements [5]
Isotope % nat. abundance atomic mass
1H 99.985 1.007825
2H 0.015 2.0140
12C 98.89 12 (definition)
13C 1.11 13.00335
14N 99.64 14.00307
15N 0.36 15.00011
16O 99.76 15.99491
17O 0.04 16.99913
18O 0.2 17.99916
28Si 92.23 27.97693
29Si 4.67 28.97649
30Si 3.10 29.97376
32S 95.0 31.97207
33S 0.76 32.97146
34S 4.22 33.96786
37Cl 24.23
35Cl 75.77 34.96885
79Br 50.69 78.9183
81Br 49.31 80.9163

See also

Footnotes and References

  1. ^ Uranium Isotopes, retrieved 14 March 2012
  2. ^ Clayton, Robert N. (1978). "Isotopic anomalies in the early solar system". Annual Review of Nuclear and Particle Science. 28: 501–522. Bibcode:1978ARNPS..28..501C. doi:10.1146/annurev.ns.28.120178.002441.
  3. ^ Zinner, Ernst (2003). "An isotopic view of the early solar system". Science. 300 (5617): 265–267. doi:10.1126/science.1080300. PMID 12690180.
  4. ^ Zinner, Ernst (1998). "Stellar nucleosynthesis and the isotopic composition of presolar grains from primitive meteorites". Annual Review of Earth and Planetary Sciences. 26: 147–188. Bibcode:1998AREPS..26..147Z. doi:10.1146/annurev.earth.26.1.147.
  5. ^ Lide, D. R., ed. (2002). CRC Handbook of Chemistry and Physics (83rd ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0483-0.

External links

Abundance

Abundance may refer to:

In science and technology:

Abundance (economics), the opposite of scarcities

Abundance (ecology), the relative representation of a species in a community

Abundance (programming language), a Forth-like computer programming language

Abundance, a property of abundant numbers

In chemistry:

Abundance (chemistry), when a substance in a reaction is present in high quantities

Abundance of the chemical elements, a measure of how common elements are

Natural abundance, the natural prevalence of different isotopes of an element on Earth

Abundance of elements in Earth's crustIn literature:

Abundance (play), a 1990 stage play written by Beth Henley

Al-Kawthar ("Abundance"), the 108th sura of the Qur'an

Abundance: The Future Is Better Than You Think, a 2012 book by Peter Diamandis and Steven KotlerOther usesAbundance Generation, a renewable energy investment platform

Fountain de la Abundancia, a former fountain in Madrid

Abundance, Royal Abundance and Abundance Declared, bids in the card game Solo whist; sometimes spelled "abondance"

Carbon-13 nuclear magnetic resonance

Carbon-13 (C13) nuclear magnetic resonance (most commonly known as carbon-13 NMR or 13C NMR or sometimes simply referred to as carbon NMR) is the application of nuclear magnetic resonance (NMR) spectroscopy to carbon. It is analogous to proton NMR (1H NMR) and allows the identification of carbon atoms in an organic molecule just as proton NMR identifies hydrogen atoms. As such 13C NMR is an important tool in chemical structure elucidation in organic chemistry. 13C NMR detects only the 13C isotope of carbon, whose natural abundance is only 1.1%, because the main carbon isotope, 12C, is not detectable by NMR since it has zero net spin.

Conversion electron Mössbauer spectroscopy

Conversion electron Mössbauer spectroscopy (CEMS) is a Mössbauer spectroscopy technique based on conversion electron.

The CEM spectrum can be obtained either by collecting essentially all the electrons leaving the surface (integral technique), or by selecting the ones in a given energy range by means of a beta ray spectrometer (differential or depth selective CEMS).

This method allows the use of simple and inexpensive detecting equipment, mainly flow-type proportional detectors in which large counting rates can be obtained. This last characteristic makes possible the study of samples with the natural abundance of the Mössbauer isotope. The information furnished by the integral measurements can be increased by using various angles of incidence or by depositing thin layers of inert material on the sample.

Inert gas

An inert gas is a gas which does not undergo chemical reactions under a set of given conditions. The noble gases often do not react with many

substances, and were historically referred to as the inert gases. Inert gases are used generally to avoid unwanted chemical reactions degrading a sample. These undesirable chemical reactions are often oxidation and hydrolysis reactions with the oxygen and moisture in air. The term inert gas is context-dependent because several of the noble gases can be made to react under certain conditions.

Purified argon and nitrogen gases are most commonly used as inert gases due to their high natural abundance (78.2% N2, 1% Ar in air) and low relative cost.

Unlike noble gases, an inert gas is not necessarily elemental and is often a compound gas. Like the noble gases the tendency for non-reactivity is due to the valence, the outermost electron shell, being complete in all the inert gases. This is a tendency, not a rule, as noble gases and other "inert" gases can react to form compounds.Some name of inert gases are (1)Helium,(2)Radon,(3)Neon

(4)Argon (5)Xenon

Isotopes of aluminium

Aluminium or aluminum (13Al) has 25 known isotopes from 19Al to 43Al and 4 known isomers. Only 27Al (stable isotope) and 26Al (radioactive isotope, t1/2 = 7.2 × 105 y) occur naturally, however 27Al has a natural abundance of >99.9%. 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.Cosmogenic aluminium-26 was first applied in studies of the Moon and meteorites. Meteorite fragments, after departure from their parent bodies, are exposed to intense cosmic-ray bombardment during their travel through space, causing substantial 26Al production. After falling to Earth, atmospheric shielding protects the meteorite fragments from further 26Al production, and its decay can then be used to determine the meteorite's terrestrial age. Meteorite research has also shown that 26Al was relatively abundant at the time of formation of our planetary system. Most meteoriticists believe that the energy released by the decay of 26Al was responsible for the melting and differentiation of some asteroids after their formation 4.55 billion years ago.

Isotopes of chromium

Naturally occurring chromium (24Cr) is composed of four stable isotopes; 50Cr, 52Cr, 53Cr, and 54Cr with 52Cr being the most abundant (83.789% natural abundance). 50Cr is suspected of decaying by β+β+ to 50Ti with a half-life of (more than) 1.8×1017 years. Twenty-two radioisotopes, all of which are entirely synthetic, have been characterized with the most stable being 51Cr with a half-life of 27.7 days. All of the remaining radioactive isotopes have half-lives that are less than 24 hours and the majority of these have half-lives that are less than 1 minute, the least stable being 66Cr with a half-life of 10 milliseconds. This element also has 2 meta states, 45mCr, the more stable one, and 59mCr, the least stable isotope or isomer.

53Cr is the radiogenic decay product of 53Mn. Chromium isotopic contents are typically combined with manganese isotopic contents and have found application in isotope geology. Mn-Cr isotope ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites indicate an initial 53Mn/55Mn ratio that suggests Mn-Cr isotope systematics must result from in-situ decay of 53Mn in differentiated planetary bodies. Hence 53Cr provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system. The same isotope is preferentially involved in certain leaching reactions, thereby allowing its abundance in seawater sediments to be used as a proxy for atmospheric oxygen concentrations.The isotopes of chromium range from 42Cr to 67Cr. The primary decay mode before the most abundant stable isotope, 52Cr, is electron capture and the primary mode after is beta decay.

Isotopes of erbium

Naturally occurring erbium (68Er) is composed of 6 stable isotopes, with 166Er being the most abundant (33.503% natural abundance). Thirty radioisotopes have been characterized with between 74 and 108 neutrons, or 142 to 177 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 13 meta states, with the most stable being 167mEr (t1/2 2.269 seconds).

The isotopes of erbium range in atomic weight from 141.9723 u (142Er) to 176.9541 u (177Er). The primary decay mode before the most abundant stable isotope, 166Er, is electron capture, and the primary mode after is beta decay. The primary decay products before 166Er are holmium isotopes, and the primary products after are thulium isotopes.

Isotopes of germanium

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).

Stable 74Ge is the most common isotope, having a natural abundance of approximately 36%. 76Ge is the least common with a natural abundance of approximately 7%. When bombarded with alpha particles, the isotopes 72Ge and 76Ge will generate stable 75As and 77Se, releasing high energy electrons in the process.At least 27 radioisotopes have also been synthesized ranging in atomic mass from 58 to 89. The most stable of these is 68Ge, decaying by electron capture with a half-life of 270.95 d. It decays to the medically useful positron-emitting isotope 68Ga. (See gallium-68 generator for notes on the source of this isotope, and its medical use). The least stable known germanium isotope is 60Ge with a half-life of 30 ms.

While most of germanium's radioisotopes decay by beta decay, 61Ge and 64Ge decay by β+ delayed proton emission. 84Ge through 87Ge also have minor β− delayed neutron emission decay paths.

Isotopes of lanthanum

Naturally occurring lanthanum (57La) is composed of one stable (139La) and one radioactive (138La) isotope, with the stable isotope, 139La, being the most abundant (99.91% natural abundance). There are 38 radioisotopes that have been characterized, with the most stable being 138La, with a half-life of 1.02×1011 years; 137La, with a half-life of 60,000 years and 140La, with a half-life of 1.6781 days. The remaining radioactive isotopes have half-lives that are less than a day and the majority of these have half-lives that are less than 1 minute. This element also has 12 nuclear isomers, the longest-lived of which is 132mLa, with a half-life of 24.3 minutes.

The isotopes of lanthanum range in atomic weight from 116.95 u (117La) to 154.96 u (155La).

Isotopes of lutetium

Naturally occurring lutetium (71Lu) is composed of 1 stable isotope 175Lu (97.41% natural abundance) and one long-lived radioisotope, 176Lu with a half-life of 3.78 × 1010 years (2.59% natural abundance). Thirty-four radioisotopes have been characterized, with the most stable, besides 176Lu, being 174Lu with a half-life of 3.31 years, and 173Lu with a half-life of 1.37 years. All of the remaining radioactive isotopes have half-lives that are less than 9 days, and the majority of these have half-lives that are less than half an hour. This element also has 18 meta states, with the most stable being 177mLu (t1/2 160.4 days), 174mLu (t1/2 142 days) and 178mLu (t1/2 23.1 minutes).

The isotopes of lutetium range in atomic weight from 149.973 (150Lu) to 183.961 (184Lu). The primary decay mode before the most abundant stable isotope, 175Lu, is electron capture (with some alpha and positron emission), and the primary mode after is beta emission. The primary decay products before 175Lu are isotopes of ytterbium and the primary products after are isotopes of hafnium.

Isotopes of neodymium

Naturally occurring neodymium (60Nd) is composed of 5 stable isotopes, 142Nd, 143Nd, 145Nd, 146Nd and 148Nd, with 142Nd being the most abundant (27.2% natural abundance), and 2 long-lived radioisotopes, 144Nd and 150Nd. In all, 33 radioisotopes of neodymium have been characterized up to now, with the most stable being naturally occurring isotopes 144Nd (alpha decay, a half-life (t1/2) of 2.29×1015 years) and 150Nd (double beta decay, t1/2 of 7×1018 years). All of the remaining radioactive isotopes have half-lives that are less than 12 days, and the majority of these have half-lives that are less than 70 seconds; the most stable artificial isotope is 147Nd with a half-life of 10.98 days. This element also has 13 known meta states with the most stable being 139mNd (t1/2 5.5 hours), 135mNd (t1/2 5.5 minutes) and 133m1Nd (t1/2 ~70 seconds).

The primary decay modes before the most abundant stable isotope, 142Nd, are electron capture and positron decay, and the primary mode after is beta decay. The primary decay products before 142Nd are element Pr (praseodymium) isotopes and the primary products after are element Pm (promethium) isotopes.

Isotopes of nickel

Naturally occurring nickel (28Ni) is composed of five stable isotopes; 58Ni, 60Ni, 61Ni, 62Ni and 64Ni with 58Ni being the most abundant (68.077% natural abundance). 26 radioisotopes have been characterised with the most stable being 59Ni with a half-life of 76,000 years, 63Ni with a half-life of 100.1 years, and 56Ni with a half-life of 6.077 days. All of the remaining radioactive isotopes have half-lives that are less than 60 hours and the majority of these have half-lives that are less than 30 seconds. This element also has 1 meta state.

Isotopes of titanium

Naturally occurring titanium (22Ti) is composed of five stable isotopes; 46Ti, 47Ti, 48Ti, 49Ti and 50Ti with 48Ti being the most abundant (73.8% natural abundance). Twenty-one radioisotopes have been characterized, with the most stable being 44Ti with a half-life of 60 years, 45Ti with a half-life of 184.8 minutes, 51Ti with a half-life of 5.76 minutes, and 52Ti with a half-life of 1.7 minutes. All of the remaining radioactive isotopes have half-lives that are less than 33 seconds, and the majority of these have half-lives that are less than half a second.The isotopes of titanium range in atomic mass from 38.01 u (38Ti) to 62.99 u (63Ti). The primary decay mode before the most abundant stable isotope, 48Ti, is β+ and the primary mode after is β−. The primary decay products before 48Ti are scandium isotopes and the primary products after are vanadium isotopes.

Isotopes of zinc

Naturally occurring zinc (30Zn) is composed of the 5 stable isotopes 64Zn, 66Zn, 67Zn, 68Zn, and 70Zn with 64Zn being the most abundant (48.6% natural abundance). Twenty-five radioisotopes have been characterised with the most abundant and stable being 65Zn with a half-life of 244.26 days, and 72Zn with a half-life of 46.5 hours. All of the remaining radioactive isotopes have half-lives that are less than 14 hours and the majority of these have half-lives that are less than 1 second. This element also has 10 meta states.

Zinc has been proposed as a "salting" material for nuclear weapons. A jacket of isotopically enriched 64Zn, irradiated by the intense high-energy neutron flux from an exploding thermonuclear weapon, would transmute into the radioactive isotope 65Zn with a half-life of 244 days and produce approximately 1.115 MeV of gamma radiation, significantly increasing the radioactivity of the weapon's fallout for several days. Such a weapon is not known to have ever been built, tested, or used.

Ivyton, Kentucky

Ivyton is an unincorporated community in Magoffin County, Kentucky, United States. It lies along Route 114 southeast of the city of Salyersville, the county seat of Magoffin County. Its elevation is 984 feet (300 m).A post office was established in the community in 1883. The town was named for the natural abundance of ivy.

Kinetic isotope effect

In physical organic chemistry, a kinetic isotope effect (KIE) is the change in the reaction rate of a chemical reaction when one of the atoms in the reactants is replaced by one of its isotopes. Formally, it is the ratio of rate constants for the reactions involving the light (kL) and the heavy (kH) isotopically substituted reactants (isotopologues):

This change in reaction rate is a quantum mechanical effect that primarily results from heavier isotopologues having lower vibrational frequencies compared to their lighter counterparts. In most cases, this implies a greater energetic input needed for heavier isotopologues to reach the transition state (or, in rare cases, the dissociation limit), and consequently, a slower reaction rate. The study of kinetic isotope effects can help the elucidation of the reaction mechanism of certain chemical reactions and is occasionally exploited in drug development to improve unfavorable pharmacokinetics by protecting metabolically-vulnerable C-H bonds.

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.

Podophyllotoxin

Podophyllotoxin (PPT), also known as podofilox, is a medical cream that is used to treat genital warts and molluscum contagiosum. It is not recommended in HPV infections without external warts. It can be applied either by a healthcare provider or the person themselves.It is a non-alkaloid toxin lignan extracted from the roots and rhizomes of Podophyllum species. A less refined form known as podophyllum resin is also available, but has greater side effects.Podophyllotoxin was isolated in 1880. In the United Kingdom the price for the NHS of 3.5 ml of medication is about 14.49 pounds. In the United States the price of a course of treatment is more than $200.

Two-dimensional nuclear magnetic resonance spectroscopy

Two-dimensional nuclear magnetic resonance spectroscopy (2D NMR) is a set of nuclear magnetic resonance spectroscopy (NMR) methods which give data plotted in a space defined by two frequency axes rather than one. Types of 2D NMR include correlation spectroscopy (COSY), J-spectroscopy, exchange spectroscopy (EXSY), and nuclear Overhauser effect spectroscopy (NOESY). Two-dimensional NMR spectra provide more information about a molecule than one-dimensional NMR spectra and are especially useful in determining the structure of a molecule, particularly for molecules that are too complicated to work with using one-dimensional NMR.

The first two-dimensional experiment, COSY, was proposed by Jean Jeener, a professor at the Université Libre de Bruxelles, in 1971. This experiment was later implemented by Walter P. Aue, Enrico Bartholdi and Richard R. Ernst, who published their work in 1976.

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