Standard atomic weight

  • example: copper in terrestrial sources. two isotopes are present: copper-63 (62.9) and copper-65 (64.9), in abundances 69% + 31%. the standard atomic weight (ar, standard) for copper is the average, weighted by their natural abundance, and then divided by the atomic mass constant mu.[1]

    the standard atomic weight (ar, standard) of a chemical element is the arithmetic mean of the relative atomic masses (ar) of all isotopes of that element weighted by each isotope's abundance on earth. the standard atomic weight of each chemical element is determined and published by the commission on isotopic abundances and atomic weights (ciaaw) of the international union of pure and applied chemistry (iupac) based on natural, stable, terrestrial sources of the element. it is the most common and practical atomic weight used by scientists.

    for example, isotope 63cu (ar = 62.929) constitutes 69% of the copper on earth, the rest being 65cu (ar = 64.927), then

    the specified definition is to use many representative sources (samples) from the earth, so that the value can widely be used as 'the' atomic weight for real life substances—for example, in pharmaceuticals and scientific research. atomic weights are specific to single sources and samples of an element, such as the atomic weight of carbon in a particular bone from a particular archeological site. standard atomic weight generalizes such values to the range of atomic weights that a chemist might expect to derive from many random samples from earth. this range is the cause of the interval notation in some standard atomic weight values.

    out of the 118 known chemical elements, 80 have stable isotopes and 84 have this earth-environment based value. typically, such a value is, for example helium: ar, standard(he) = 4.002602(2). the "(2)" indicates the uncertainty in the last digit shown, to read 4.002602±0.000002. iupac also publishes abridged values, rounded to five significant figures. for helium, ar, abridged(he) = 4.0026.

    for thirteen elements the samples diverge on this value, because their sample sources have had a different decay history. for example, thallium (tl) in sedimentary rocks has a different isotopic composition than in igneous rocks and volcanic gases. for these elements, the standard atomic weight is noted as an interval: ar, standard(tl) = [204.38, 204.39]. with such an interval, for less demanding situations, iupac also publishes a conventional value. for thallium, ar, conventional(tl) = 204.38.

  • definition
  • determination of relative atomic mass
  • naming controversy
  • published values
  • list of atomic weights
  • see also
  • references
  • external links

Example: copper in terrestrial sources. Two isotopes are present: copper-63 (62.9) and copper-65 (64.9), in abundances 69% + 31%. The standard atomic weight (Ar, standard) for copper is the average, weighted by their natural abundance, and then divided by the atomic mass constant mu.[1]

The standard atomic weight (Ar, standard) of a chemical element is the arithmetic mean of the relative atomic masses (Ar) of all isotopes of that element weighted by each isotope's abundance on Earth. The standard atomic weight of each chemical element is determined and published by the Commission on Isotopic Abundances and Atomic Weights (CIAAW) of the International Union of Pure and Applied Chemistry (IUPAC) based on natural, stable, terrestrial sources of the element. It is the most common and practical atomic weight used by scientists.

For example, isotope 63Cu (Ar = 62.929) constitutes 69% of the copper on Earth, the rest being 65Cu (Ar = 64.927), then

The specified definition is to use many representative sources (samples) from the Earth, so that the value can widely be used as 'the' atomic weight for real life substances—for example, in pharmaceuticals and scientific research. Atomic weights are specific to single sources and samples of an element, such as the atomic weight of carbon in a particular bone from a particular archeological site. Standard atomic weight generalizes such values to the range of atomic weights that a chemist might expect to derive from many random samples from Earth. This range is the cause of the interval notation in some standard atomic weight values.

Out of the 118 known chemical elements, 80 have stable isotopes and 84 have this Earth-environment based value. Typically, such a value is, for example helium: Ar, standard(He) = 4.002602(2). The "(2)" indicates the uncertainty in the last digit shown, to read 4.002602±0.000002. IUPAC also publishes abridged values, rounded to five significant figures. For helium, Ar, abridged(He) = 4.0026.

For thirteen elements the samples diverge on this value, because their sample sources have had a different decay history. For example, thallium (Tl) in sedimentary rocks has a different isotopic composition than in igneous rocks and volcanic gases. For these elements, the standard atomic weight is noted as an interval: Ar, standard(Tl) = [204.38, 204.39]. With such an interval, for less demanding situations, IUPAC also publishes a conventional value. For thallium, Ar, conventional(Tl) = 204.38.