Commission on Isotopic Abundances and Atomic Weights

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Commission on Isotopic Abundances and Atomic Weights
CIAAW logo.png
AbbreviationCIAAW
Formation1899;123 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. 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. [1] The biennial CIAAW Standard Atomic Weights are accepted as the authoritative source in science and appear worldwide on the periodic table wall charts. [2]

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. [3] The latest CIAAW report was published in February 2016. [4] After 20 May 2019 a new definition for kelvin came into force based on the Boltzmann constant.

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. [5] 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. [6]

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. [7] 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. [8] 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. [9]

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. [10]

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. [11] 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. [12] In the 1957 Paris meeting the Commission put forward a proposal for a carbon-12 scale. [13] 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. [14]

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. [15] 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. [16] [17] 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. [18]

Affiliations and name

The Commission on Isotopic Abundances and Atomic Weights has undergone many name changes:

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" [19] while he was a member of the Commission. [20] 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. [21] Incidentally, the 1925 Atomic Weights report was signed by three Nobel laureates. [22]

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), [23] 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:

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 widely used in physics and chemistry. It is defined as 112 of the mass of an unbound neutral atom of carbon-12 in its nuclear and electronic ground state and at rest. The atomic mass constant, denoted mu, is defined identically, giving mu = m(12C)/12 = 1 Da.

Chemical symbol Abbreviations used in chemistry

Chemical symbols are the abbreviations used in chemistry for chemical elements, functional groups and chemical compounds. Element symbols for chemical elements normally consist of one or two letters from the Latin alphabet and are written with the first letter capitalised. In Chinese, each chemical element has a dedicated character, usually created for the purpose. However, Latin symbols are also used, especially in formulas.

Relative atomic mass or 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; hence the value is said to be relative.

In chemistry, the amount of substance n in a given sample of matter is defined as the quantity or number of discrete atomic-scale particles in it divided by the Avogadro constant NA. The particles or entities may be molecules, atoms, ions, electrons, or other, depending on the context, and should be specified (e.g. amount of sodium chloride nNaCl). The value of the Avogadro constant NA has been defined as 6.02214076×1023 mol−1. The mole (symbol: mol) is a unit of amount of substance in the International System of Units, defined (since 2019) by fixing the Avogadro constant at the given value. Sometimes, the amount of substance is referred to as the chemical amount.

Vienna Standard Mean Ocean Water Standard defining the isotopic composition of fresh water originating from ocean water

Vienna Standard Mean Ocean Water (VSMOW) is an isotopic standard for water. Despite the name, VSMOW is pure water with no salt or other chemicals found in the oceans. The VSMOW standard was promulgated by the International Atomic Energy Agency in 1968, and since 1993 continues to be evaluated and studied by the IAEA along with the European Institute for Reference Materials and Measurements and the American National Institute of Standards and Technology. The standard includes both the established values of stable isotopes found in waters and calibration materials provided for standardization and interlaboratory comparisons of instruments used to measure these values in experimental materials.

Gold (79Au) has one stable isotope, 197Au, and 36 radioisotopes, with 195Au being the most stable with a half-life of 186 days. Gold is currently considered the heaviest monoisotopic element. Bismuth formerly held that distinction until alpha-decay of the 209Bi isotope was observed. All isotopes of gold are either radioactive or, in the case of 197Au, observationally stable, meaning that 197Au is predicted to be radioactive but no actual decay has been observed.

Naturally occurring platinum (78Pt) consists of five stable isotopes (192Pt, 194Pt, 195Pt, 196Pt, 198Pt) and one very long-lived (half-life 6.50×1011 years) radioisotope (190Pt). There are also 34 known synthetic radioisotopes, the longest-lived of which is 193Pt with a half-life of 50 years. All other isotopes have half-lives under a year, most under a day. All isotopes of platinium are either radioactive or observationally stable, meaning that they are predicted to be radioactive but no actual decay has been observed.

Naturally occurring rhenium (75Re) is 37.4% 185Re, which is stable (although it is predicted to decay), and 62.6% 187Re, which is unstable but has a very long half-life (4.12×1010 years). Among elements with a known stable isotope, only indium and tellurium similarly occur with a stable isotope in lower abundance than the long-lived radioactive isotope.

Natural hafnium (72Hf) consists of five stable isotopes (176Hf, 177Hf, 178Hf, 179Hf, and 180Hf) and one very long-lived radioisotope, 174Hf, with a half-life of 2×1015 years. In addition, there are 30 other known radionuclides, the most stable of which is 182Hf with a half-life of 8.9×106 years. This extinct radionuclide is used in hafnium–tungsten dating to study the chronology of planetary differentiation.

Naturally occurring thulium (69Tm) is composed of one stable isotope, 169Tm. Thirty-four radioisotopes have been characterized, with the most stable being 171Tm with a half-life of 1.92 years, 170Tm with a half-life of 128.6 days, 168Tm with a half-life of 93.1 days, and 167Tm with a half-life of 9.25 days. All of the remaining radioactive isotopes have half-lives that are less than 64 hours, and the majority of these have half-lives that are less than 2 minutes. This element also has 26 meta states, with the most stable being 164mTm, 160mTm and 155mTm.

Natural holmium (67Ho) contains one 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.

Antimony (51Sb) occurs in two stable isotopes, 121Sb and 123Sb. There are 35 artificial radioactive isotopes, the longest-lived of which are 125Sb, with a half-life of 2.75856 years; 124Sb, with a half-life of 60.2 days; and 126Sb, with a half-life of 12.35 days. All other isotopes have half-lives less than 4 days, most less than an hour.

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.

Naturally occurring cobalt (27Co) is composed of a single stable isotope, 59Co. Twenty-eight radioisotopes have been characterized with the most stable being 60Co with a half-life of 5.2714 years, 57Co with a half-life of 271.8 days, 56Co with a half-life of 77.27 days, and 58Co with a half-life of 70.86 days. All of the remaining radioactive isotopes have half-lives that are less than 18 hours and the majority of these have half-lives that are less than 1 second. This element also has 11 meta states, all of which have half-lives less than 15 minutes.

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.

Naturally occurring scandium (21Sc) is composed of one stable isotope, 45Sc. Twenty-five radioisotopes have been characterized, with the most stable being 46Sc with a half-life of 83.8 days, 47Sc with a half-life of 3.35 days, and 48Sc with a half-life of 43.7 hours and 44Sc with a half-life of 3.97 hours. All the remaining isotopes have half-lives that are less than four hours, and the majority of these have half-lives that are less than two minutes, the least stable being proton unbound 39Sc with a half-life shorter than 300 nanoseconds. This element also has 13 meta states with the most stable being 44m2Sc.

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. Fourteen radioisotopes are also known, with atomic masses ranging from 10 to 25, along with one nuclear isomer, 11mN. 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.

Standard atomic weight Relative atomic mass as defined by IUPAC (CIAAW)

The standard atomic weight of a chemical element 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 constitutes 69% of the copper on Earth, the rest being 65Cu, so

Mass (mass spectrometry) 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.

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

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