Commission on Isotopic Abundances and Atomic Weights

Last updated
Commission on Isotopic Abundances and Atomic Weights
AbbreviationCIAAW
Formation1899;125 years ago (1899)
TypeInternational scientific organization
PurposeTo provide internationally recommended values of isotopic composition and atomic weights of elements
Region served
Worldwide
Official language
English
Chair
Johanna Irrgeher
Secretary
Jochen Vogl
Parent organization
 IUPAC (since 1920)
Website www.ciaaw.org

The Commission on Isotopic Abundances and Atomic Weights (CIAAW) is an international scientific committee of the International Union of Pure and Applied Chemistry (IUPAC) under its Division of Inorganic Chemistry. [1] Since 1899, it is entrusted with periodic critical evaluation of atomic weights of chemical elements and other cognate data, such as the isotopic composition of elements. [2] The biennial CIAAW Standard Atomic Weights are accepted as the authoritative source in science and appear worldwide on the periodic table wall charts. [3]

Contents

The use of CIAAW Standard Atomic Weights is also required legally, for example, in calculation of calorific value of natural gas (ISO 6976:1995), or in gravimetric preparation of primary reference standards in gas analysis (ISO 6142:2006). In addition, until 2019 the definition of Kelvin, the SI unit for thermodynamic temperature, made direct reference to the isotopic composition of oxygen and hydrogen as recommended by CIAAW. [4] The latest CIAAW report was published in May 2022. [5]

Establishment

U.S. Government FW Clarke Public Domain.jpg
Thomas Edward Thorpe.jpg
The inaugural members of the International Committee on Atomic Weights were:

Although the atomic weight had taken on the concept of a constant of nature like the speed of light, the lack of agreement on accepted values created difficulties in trade. Quantities measured by chemical analysis were not being translated into weights in the same way by all parties and standardization became an urgent matter. [6] With so many different values being reported, the American Chemical Society (ACS), in 1892, appointed a permanent committee to report on a standard table of atomic weights for acceptance by the Society. Clarke, who was then the chief chemist for the U.S. Geological Survey, was appointed a committee of one to provide the report. He presented the first report at the 1893 annual meeting and published it in January 1894. [7]

In 1897, the German Society of Chemistry, following a proposal by Hermann Emil Fischer, appointed a three-person working committee to report on atomic weights. The committee consisted of Chairman Prof. Hans H. Landolt (Berlin University), Prof. Wilhelm Ostwald (University of Leipzig), and Prof. Karl Seubert (University of Hanover). This committee published its first report in 1898, in which the committee suggested the desirability of an international committee on atomic weights. On 30 March 1899 Landolt, Ostwald and Seubert issued an invitation to other national scientific organizations to appoint delegates to the International Committee on Atomic Weights. Fifty-eight members were appointed to the Great International Committee on Atomic Weights, including Frank W. Clarke. [8] The large committee conducted its business by correspondence to Landolt which created difficulties and delays associated with correspondence among fifty-eight members. As a result, on 15 December 1899, the German committee asked the International members to select a small committee of three to four members. [9] In 1902, Prof. Frank W. Clarke (USA), Prof. Karl Seubert (Germany), and Prof. Thomas Edward Thorpe (UK) were elected, and the International Committee on Atomic Weights published its inaugural report in 1903 under the chairmanship of Prof. Clarke. [10]

Function

Since 1899, the Commission periodically and critically evaluates the published scientific literature and produces the Table of Standard Atomic Weights. In recent times, the Table of Standard Atomic Weights has been published biennially. Each recommended standard atomic-weight value reflects the best knowledge of evaluated, published data. In the recommendation of standard atomic weights, CIAAW generally does not attempt to estimate the average or composite isotopic composition of the Earth or of any subset of terrestrial materials. Instead, the Commission seeks to find a single value and symmetrical uncertainty that would include almost all substances likely to be encountered. [11]

Notable decisions

Many notable decisions have been made by the Commission over its history. Some of these are highlighted below.

The inaugural 1903 report of the International Atomic Weights Commission Inaugural Atomic Weights report 1903.png
The inaugural 1903 report of the International Atomic Weights Commission

International atomic weight unit: H=1 or O=16

Though Dalton proposed setting the atomic weight of hydrogen as unity in 1803, many other proposals were popular throughout the 19th century. By the end of the 19th century, two scales gained popular support: H=1 and O=16. This situation was undesired in science and in October 1899, the inaugural task of the International Commission on Atomic Weights was to decide on the international scale and the oxygen scale became the international standard. [12] The endorsement of the oxygen scale created significant backlash in the chemistry community, and the inaugural Atomic Weights Report was thus published using both scales. This practice soon ceded and the oxygen scale remained the international standard for decades to come. Nevertheless, when the Commission joined the IUPAC in 1920, it was asked to revert to the H=1 scale, which it rejected.

Modern unit: 12C=12

With the discovery of oxygen isotopes in 1929, a situation arose where chemists based their calculations on the average atomic mass (atomic weight) of oxygen whereas physicists used the mass of the predominant isotope of oxygen, oxygen-16. This discrepancy became undesired and a unification between the chemistry and physics was necessary. [13] In the 1957 Paris meeting the Commission put forward a proposal for a carbon-12 scale. [14] The carbon-12 scale for atomic weights and nuclide masses was approved by IUPAP (1960) and IUPAC (1961) and it is still in use worldwide. [15]

Uncertainty of the atomic weights

In the early 20th century, measurements of the atomic weight of lead showed significant variations depending on the origin of the sample. These differences were considered to be an exception attributed to lead isotopes being products of the natural radioactive decay chains of uranium. In 1930s, however, Malcolm Dole reported that the atomic weight of oxygen in air was slightly different from that in water. [16] Soon thereafter, Alfred Nier reported natural variation in the isotopic composition of carbon. It was becoming clear that atomic weights are not constants of nature. At the Commission’s meeting in 1951, it was recognized that the isotopic-abundance variation of sulfur had a significant effect on the internationally accepted value of an atomic weight. In order to indicate the span of atomic-weight values that may apply to sulfur from different natural sources, the value ± 0.003 was attached to the atomic weight of sulfur. By 1969, the Commission had assigned uncertainties to all atomic-weight values.

Excerpt of the IUPAC Periodic Table of the Elements 2011 showing the interval notation of the standard atomic weights of boron, carbon, and nitrogen IUPAC Periodic Table of the Elements 2011.jpg
Excerpt of the IUPAC Periodic Table of the Elements 2011 showing the interval notation of the standard atomic weights of boron, carbon, and nitrogen

Interval notation

At its meeting in 2009 in Vienna, the Commission decided to express the standard atomic weight of hydrogen, carbon, oxygen, and other elements in a manner that clearly indicates that the values are not constants of nature. [17] [18] For example, writing the standard atomic weight of hydrogen as [1.007 84, 1.008 11] shows that the atomic weight in any normal material will be greater than or equal to 1.007 84 and will be less than or equal to 1.008 11. [19]

Affiliations and name

The IUPAC Commission on Isotopic Abundances and Atomic Weights has undergone several name changes between its founding in 1899 and 2002, when it received its present name [22] :

Notable members

The Harvard chemist Theodore W. Richards, a member of the International Atomic Weights Commission, was awarded the 1914 Nobel Prize in Chemistry for his work on atomic weight determination Richards Theodore William lab.jpg
The Harvard chemist Theodore W. Richards, a member of the International Atomic Weights Commission, was awarded the 1914 Nobel Prize in Chemistry for his work on atomic weight determination

Since its establishment, many notable chemists have been members of the Commission. Notably, eight Nobel laureates have served in the Commission: Henri Moissan (1903-1907), Wilhelm Ostwald (1906-1916), Francis William Aston, Frederick Soddy, Theodore William Richards, Niels Bohr, Otto Hahn and Marie Curie.

Richards was awarded the 1914 Nobel Prize in Chemistry "in recognition of his accurate determinations of the atomic weight of a large number of chemical elements" [23] while he was a member of the Commission. [24] Likewise, Francis Aston was a member of the Commission when he was awarded the 1922 Nobel Prize in Chemistry for his work on isotope measurements. [25] Incidentally, the 1925 Atomic Weights report was signed by three Nobel laureates. [26]

Among other notable scientists who have served on the Commission were Georges Urbain (discoverer of lutetium, though priority was disputed with Carl Auer von Welsbach), André-Louis Debierne (discoverer of actinium, though priority has been disputed with Friedrich Oskar Giesel), Marguerite Perey (discoverer of francium), Georgy Flyorov (namesake of the element flerovium), [27] Robert Whytlaw-Gray (first isolated radon), and Arne Ölander (Secretary and Member of the Nobel Committee for Chemistry).

Chairs of the Commission

Since its establishment, the chairs of the Commission have been: [28]

In 1950, the Spanish chemist Enrique Moles became the first Secretary of the Commission when this position was created.

See also

Related Research Articles

The dalton or unified atomic mass unit is a unit of mass defined as ⁠1/12⁠ of the mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state and at rest. It is a non-SI unit accepted for use with SI. The atomic mass constant, denoted mu, is defined identically, giving mu = ⁠1/12⁠m(12C) = 1 Da.

Relative atomic mass, also known by the deprecated synonym atomic weight, is a dimensionless physical quantity defined as the ratio of the average mass of atoms of a chemical element in a given sample to the atomic mass constant. The atomic mass constant is defined as being ⁠1/12⁠ of the mass of a carbon-12 atom. Since both quantities in the ratio are masses, the resulting value is dimensionless. These definitions remain valid even after the 2019 revision of the SI.

<span class="mw-page-title-main">Mass number</span> Number of heavy particles in the atomic nucleus

The mass number (symbol A, from the German word: Atomgewicht, "atomic weight"), also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. It is approximately equal to the atomic (also known as isotopic) mass of the atom expressed in atomic mass units. Since protons and neutrons are both baryons, the mass number A is identical with the baryon number B of the nucleus (and also of the whole atom or ion). The mass number is different for each isotope of a given chemical element, and the difference between the mass number and the atomic number Z gives the number of neutrons (N) in the nucleus: N = A − Z.

<span class="mw-page-title-main">Amount of substance</span> Ratio of the number of particles in a sample to Avogadros constant

In chemistry, the amount of substance (symbol n) in a given sample of matter is defined as a ratio (n = N/NA) between the number of elementary entities (N) and the Avogadro constant (NA). The entities are usually molecules, atoms, ions, or ion pairs of a specified kind. The particular substance sampled may be specified using a subscript, e.g., the amount of sodium chloride (NaCl) would be denoted as nNaCl. The unit of amount of substance in the International System of Units is the mole (symbol: mol), a base unit. Since 2019, the value of the Avogadro constant NA is defined to be exactly 6.02214076×1023 mol−1. Sometimes, the amount of substance is referred to as the chemical amount or, informally, as the "number of moles" in a given sample of matter.

Vienna Standard Mean Ocean Water (VSMOW) is an isotopic standard for water, that is, a particular sample of water whose proportions of different isotopes of hydrogen and oxygen are accurately known. VSMOW is distilled from ocean water and does not contain salt or other impurities. Published and distributed by the Vienna-based International Atomic Energy Agency in 1968, the standard and its essentially identical successor, VSMOW2, continue to be used as a reference material.

Naturally occurring lutetium (71Lu) is composed of one stable isotope 175Lu and one long-lived radioisotope, 176Lu with a half-life of 37 billion years. Forty 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, 174mLu and 178mLu.

Naturally occurring ytterbium (70Yb) is composed of seven stable isotopes: 168Yb, 170Yb–174Yb, and 176Yb, with 174Yb being the most abundant. 30 radioisotopes have been characterized, with the most stable being 169Yb with a half-life of 32.014 days, 175Yb with a half-life of 4.185 days, and 166Yb with a half-life of 56.7 hours. All of the remaining radioactive isotopes have half-lives that are less than 2 hours, and the majority of these have half-lives that are less than 20 minutes. This element also has 18 meta states, with the most stable being 169mYb.

Naturally occurring erbium (68Er) is composed of six stable isotopes, with 166Er being the most abundant. Thirty-nine 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.

<span class="mw-page-title-main">Isotopes of holmium</span>

Natural holmium (67Ho) contains one observationally stable isotope, 165Ho. The below table lists 36 isotopes spanning 140Ho through 175Ho as well as 33 nuclear isomers. Among the known synthetic radioactive isotopes; the most stable one is 163Ho, with a half-life of 4,570 years. All other radioisotopes have half-lives not greater than 1.117 days in their ground states, and most have half-lives under 3 hours.

Naturally occurring cerium (58Ce) is composed of 4 stable isotopes: 136Ce, 138Ce, 140Ce, and 142Ce, with 140Ce being the most abundant and the only one theoretically stable; 136Ce, 138Ce, and 142Ce are predicted to undergo double beta decay but this process has never been observed. There are 35 radioisotopes that have been characterized, with the most stable being 144Ce, with a half-life of 284.893 days; 139Ce, with a half-life of 137.640 days and 141Ce, with a half-life of 32.501 days. All of the remaining radioactive isotopes have half-lives that are less than 4 days and the majority of these have half-lives that are less than 10 minutes. This element also has 10 meta states.

<span class="mw-page-title-main">Isotopes of lanthanum</span>

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 39 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. Lighter isotopes mostly decay to isotopes of barium and heavy ones mostly decay to isotopes of cerium. 138La can decay to both.

Indium (49In) consists of two primordial nuclides, with the most common (~ 95.7%) nuclide (115In) being measurably though weakly radioactive. Its spin-forbidden decay has a half-life of 4.41×1014 years, much longer than the currently accepted age of the Universe.

Naturally occurring niobium (41Nb) is composed of one stable isotope (93Nb). The most stable radioisotope is 92Nb with a half-life of 34.7 million years. The next longest-lived niobium isotopes are 94Nb and 91Nb with a half-life of 680 years. There is also a meta state of 93Nb at 31 keV whose half-life is 16.13 years. Twenty-seven other radioisotopes have been characterized. Most of these have half-lives that are less than two hours, except 95Nb, 96Nb and 90Nb. The primary decay mode before stable 93Nb is electron capture and the primary mode after is beta emission with some neutron emission occurring in 104–110Nb.

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, 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. This element also has two meta states, 45mCr, the more stable one, and 59mCr, the least stable isotope or isomer.

Naturally occurring titanium (22Ti) is composed of five stable isotopes; 46Ti, 47Ti, 48Ti, 49Ti and 50Ti with 48Ti being the most abundant. 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.

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

<span class="mw-page-title-main">Standard atomic weight</span> Relative atomic mass as defined by IUPAC (CIAAW)

The standard atomic weight of a chemical element (symbol Ar°(E) for element "E") is the weighted arithmetic mean of the relative isotopic masses of all isotopes of that element weighted by each isotope's abundance on Earth. For example, isotope 63Cu (Ar = 62.929) constitutes 69% of the copper on Earth, the rest being 65Cu (Ar = 64.927), so

<span class="mw-page-title-main">Mass (mass spectrometry)</span> Physical quantities being measured

The mass recorded by a mass spectrometer can refer to different physical quantities depending on the characteristics of the instrument and the manner in which the mass spectrum is displayed.

<span class="mw-page-title-main">Atomic mass</span> Rest mass of an atom in its ground state

The atomic mass (ma or m) is the mass of an atom. Although the SI unit of mass is the kilogram (symbol: kg), atomic mass is often expressed in the non-SI unit dalton (symbol: Da) – equivalently, unified atomic mass unit (u). 1 Da is defined as 1⁄12 of the mass of a free carbon-12 atom at rest in its ground state. The protons and neutrons of the nucleus account for nearly all of the total mass of atoms, with the electrons and nuclear binding energy making minor contributions. Thus, the numeric value of the atomic mass when expressed in daltons has nearly the same value as the mass number. Conversion between mass in kilograms and mass in daltons can be done using the atomic mass constant .

The Inorganic Chemistry Division of the International Union of Pure and Applied Chemistry (IUPAC), also known as Division II, deals with all aspects of inorganic chemistry, including materials and bioinorganic chemistry, and also with isotopes, atomic weights and the periodic table. It furthermore advises the Chemical Nomenclature and Structure Representation Division on issues dealing with inorganic compounds and materials. For the general public, the most visible result of the division's work is that it evaluates and advises the IUPAC on names and symbols proposed for new elements that have been approved for addition to the periodic table. For the scientific end educational community the work on isotopic abundances and atomic weights is of fundamental importance as these numbers are continuously checked and updated.

References

  1. ↑ "IUPAC Commission on Isotopic Abundances and Atomic Weights". ciaaw.org. Retrieved 2024-08-21.
  2. ↑ "CIAAW | Commission on Isotopic Abundances and Atomic Weights". ciaaw.org. Retrieved 2024-08-21.
  3. ↑ "IUPAC, Oxford Reference". Archived from the original on 2013-09-22.{{cite journal}}: Cite journal requires |journal= (help)
  4. ↑ "Clarification of the definition of the kelvin, unit of thermodynamic temperature" (PDF). BIPM. 2005. Archived (PDF) from the original on 2013-06-26.
  5. ↑ Thomas Prohaska; Johanna Irrgeher; Jacqueline Benefield; John Karl Boehlke; Lesley Chesson; Tyler B Coplen; Tiping Ding; Philip J H Dunn; Manfred Gröning; Norman E Holden; Harro A J Meijer; Heiko Moossen; Antonio Possolo; Yoshio Takahashi; Jochen Vogl; Thomas Walczyk; Jun Wang; Michael Wieser; Shigekazu Yoneda; Xiangkun Zhu; Juris Meija (2022). "Standard Atomic weights of the elements 2021 (IUPAC Technical Report)". Pure Appl. Chem. 94 (5): 573–600. doi: 10.1515/pac-2019-0603 .
  6. ↑ E. Crawford (1992). Nationalism and Internationalism in Science, 1880-1939 (p.40). Cambridge University Press. ISBN   9780521524742. Archived from the original on 2016-05-09.
  7. ↑ "Atomic Weights and the International Committee—A Historical Review". Chemistry International. 2004. Archived from the original on 2017-07-09.
  8. ↑ L.M. Dennis, Frank Wigglesworth Clarke (National Academy of Sciences 1932) at p.143 Archived 2012-10-15 at the Wayback Machine
  9. ↑ H. Landolt; W. Ostwald; K. Seubert (1900). "Zweiter Bericht der Commission für die Festsetzung der Atomgewichte". Ber. 22 (2): 1847–1883. doi:10.1002/cber.19000330270.
  10. ↑ F.W. Clarke (1903). "Report of the International Committee on Atomic Weights". J. Am. Chem. Soc. 25 (1): 1–5. doi:10.1021/ja02003a001.
  11. ↑ Michael E. Wieser; Michael Berglund (2009). "Atomic weights of the elements 2007 (IUPAC Technical Report)" (PDF). Pure Appl. Chem. 81 (11): 2131–2156. CiteSeerX   10.1.1.540.9258 . doi:10.1351/PAC-REP-09-08-03. S2CID   98084907. Archived (PDF) from the original on 2016-03-04.
  12. ↑ Richards, Theodore William (1900). "International Atomic Weights". Proceedings of the American Academy of Arts and Sciences. 36 (10): 171–176. doi:10.2307/20020992. JSTOR   20020992.
  13. ↑ F.W. Aston (1931). "The Unit of Atomic Weight". Nature . 128 (3234): 731. Bibcode:1931Natur.128..731.. doi: 10.1038/128731a0 . S2CID   4134425.
  14. ↑ Edward Wichers (1958). "Report on Atomic Weights for 1956-1957". J. Am. Chem. Soc. 80 (16): 4121–4124. doi:10.1021/ja01549a001.
  15. ↑ Encyclopædia Britannica Archived 2013-03-08 at the Wayback Machine
  16. ↑ Malcolm Dole (1935). "The Relative Atomic Weight of Oxygen in Water and in Air". J. Am. Chem. Soc. 57 (12): 2731. doi:10.1021/ja01315a511.
  17. ↑ "Mass Migration: Chemists Revise Atomic Weights of 10 Elements". Scientific American. 16 December 2010. Archived from the original on 30 December 2011.
  18. ↑ "Atomic weights change to reflect natural variations". Chemistry World. 2010. Archived from the original on 2013-09-23.
  19. ↑ Tyler B. Coplen; Norman E. Holden (2011). "Atomic Weights: No Longer Constants of Nature". Chemistry International. Archived from the original on 2013-06-14.
  20. ↑ "Body Details". IUPAC | International Union of Pure and Applied Chemistry. Retrieved 2024-03-16.
  21. ↑ "Our History". IUPAC | International Union of Pure and Applied Chemistry. Retrieved 2024-08-21.
  22. ↑ "CIAAW | Commission on Isotopic Abundances and Atomic Weights". ciaaw.org. Retrieved 2024-11-02.
  23. ↑ "The Nobel Prize in Chemistry 1914". Archived from the original on 2017-06-29.
  24. ↑ "Nobel Prize for Richards; Chemistry Award for 1914 Goes to the Harvard Investigator". The New York Times . 13 Nov 1915. Archived from the original on 2013-09-22.
  25. ↑ F.W. Aston; et al. (1923). "Report of the International Committee on Chemical Elements: 1923". J. Am. Chem. Soc. 45 (4): 867–874. doi:10.1021/ja01657a001.
  26. ↑ F.W. Aston; et al. (1925). "International Atomic Weights 1925". J. Am. Chem. Soc. 47 (3): 597–601. doi:10.1021/ja01680a001.
  27. ↑ "Past and Current Membership Summary, CIAAW". Archived from the original on 2014-10-15.
  28. ↑ "Chairs | Commission on Isotopic Abundances and Atomic Weights". ciaaw.org. Retrieved 2024-08-21.
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