Relative atomic mass

Example for a copper source (think: various ores). Two isotopes are present: copper-63 (~62.93 u) and copper-65 (~64.93 u), in different abundances (% in the sample). The relative atomic mass (Ar) for this element in these samples is their average, taken their abundance into account, and then divided by a standardised 12C12 unit.
Incidentally, the sources are from our Earth environment, and these values have been used by CIAAW to define the standard atomic weight of copper.[1] (Note: numbers in this example are simplified)

Relative atomic mass (symbol: Ar) is a dimensionless (number only) physical quantity. In its modern definition, it is the ratio of the average mass of atoms of an element in a given sample to one unified atomic mass unit. The unified atomic mass unit, symbol u, is defined being  1⁄12 of the mass of a carbon-12 atom.[2][3] The mass of atoms can vary (between atoms of the same element), due to the presence of various isotopes of that element. Since both values in the ratio are expressed in the same unit (u), the resulting value is dimensionless; hence the value is relative.

Within one source (sample), it is a straight average over the individual atom weights (isotopes) present. Between sources, the atomic weight can vary when the source's origin (radioactive history) resulted in different isotopic concentrations. These differences are real and measurable, and can be used to identify a sample to its origin. For example, a sample of elemental carbon from volcanic methane will have a different relative atomic mass than one collected from plant or animal tissues.

The well-known standard atomic weight, or atomic weight, is a specific usage of relative atomic mass: it is the relative atomic mass, with the sources being terrestrial (taken from Earth).[4] The standard atomic weights are reprinted in a wide variety of textbooks, commercial catalogues, and periodic table wall charts. They are what chemists loosely call "atomic weights". It is the most published form of the relative atomic mass.

The continued use of the term "atomic weight" (of any element), as opposed to "relative atomic mass" has attracted considerable controversy, since at least the 1960s, mainly due to the technical difference between weight and mass in physics.[5] Both terms are officially sanctioned by IUPAC. The term "relative atomic mass" now seems to be gaining as the preferred term over "atomic weight", although in the case of "standard atomic weight", this shorter term (as opposed to the more correct "standard relative atomic mass") continues to be used.

Definition

Relative atomic mass is determined by the average atomic mass, or the weighted mean of the atomic masses of all the atoms of a particular chemical element found in a particular sample, which is then compared to the atomic mass of carbon-12.[6] This comparison is the quotient of the two weights, which makes the value dimensionless (no unit appended). This quotient also explains the word relative: the sample mass value is made relative to carbon-12.

It is used as a synonym for atomic weight (and is not to be confused with relative isotopic mass). Relative atomic mass is frequently used as a synonym for the standard atomic weight and it is correct to do so if the relative atomic mass used is that for an element from Earth under defined conditions. However, relative atomic mass covers more than standard atomic weights, and is a less specific term that may more broadly refer to non-terrestrial environments and highly specific terrestrial environments that deviate from Earth-average or have different certainties (number of significant figures) than do the standard atomic weights. In some circumstances may even be synonymous with standard atomic weight (depending on the sample, see below).

Current definition

Prevailing IUPAC definitions taken from the "Gold Book" are

atomic weight — See: relative atomic mass[7]

and

relative atomic mass (standard atomic weight) — The ratio of the average mass of the atom to the unified atomic mass unit.[8]

Here the "unified atomic mass unit" refers to  1⁄12 of the mass of an atom of 12C in its ground state.[9]

The IUPAC definition[2] of relative atomic mass is:

An atomic weight (relative atomic mass) of an element from a specified source is the ratio of the average mass per atom of the element to 1/12 of the mass of an atom of 12C.

The definition deliberately specifies "An atomic weight…", as an element will have different relative atomic masses depending on the source. For example, boron from Turkey has a lower relative atomic mass than boron from California, because of its different isotopic composition.[10][11] Nevertheless, given the cost and difficulty of isotope analysis, it is usual to use the tabulated values of standard atomic weights which are ubiquitous in chemical laboratories.


Historical amu

Older (pre-1961) historical relative scales (based on the atomic mass unit, or a.m.u., or amu) used either the oxygen-16 relative isotopic mass for reference, or else the oxygen relative atomic mass (i.e., atomic weight) for reference. See the article on the history of the modern unified atomic mass unit for the resolution of these problems in 1961.

The Standard atomic weight

The IUPAC commission CIAAW maintains a more strict definition of the relative atomic weight, named standard atomic weight or atomic weight. Requires is that the sources are terrestrial, natural, and stable with regard to radioactivity. Also there are requirements for the research process. For 84 stable elements CIAAW has determined this standard atomic weight. These values is widely published and used as 'the' atomic weight of elements for real life substances like pharmaceuticals and commercial trade.

Also, CIAAW has published abridged (rounded) values, and simplified values (for when the Earthly sources vary systematically).

Other definitions of the mass of atoms

Atomic mass (ma) is the mass of a single atom, with unit Da or u (the unified atomic mass unit). It defines the weight of a specific isotope, which is an input value for the determination of the relative atomic mass. An example for three silicon isotopes is given in determination of relative atomic mass

From this mass, the relative isotopic mass is specifically the ratio of the mass of a single atom to the mass of a unified atomic mass unit. This value too is relative, and so dimensionless.

Determination of relative atomic mass

Modern relative atomic masses (a term specific to a given element sample) are calculated from measured values of atomic mass (for each nuclide) and isotopic composition of a sample. Highly accurate atomic masses are available[12][13] for virtually all non-radioactive nuclides, but isotopic compositions are both harder to measure to high precision and more subject to variation between samples.[14][15] For this reason, the relative atomic masses of the 22 mononuclidic elements (which are the same as the isotopic masses for each of the single naturally occurring nuclides of these elements) are known to especially high accuracy. For example, there is an uncertainty of only one part in 38 million for the relative atomic mass of fluorine, a precision which is greater than the current best value for the Avogadro constant (one part in 20 million).

Isotope Atomic mass[13] Abundance[14]
Standard Range
28Si 27.97692653246(194) 92.2297(7)% 92.21–92.25%
29Si 28.976494700(22) 4.6832(5)% 4.67–4.69%
30Si 29.973770171(32) 3.0872(5)% 3.08–3.10%

The calculation is exemplified for silicon, whose relative atomic mass is especially important in metrology. Silicon exists in nature as a mixture of three isotopes: 28Si, 29Si and 30Si. The atomic masses of these nuclides are known to a precision of one part in 14 billion for 28Si and about one part in one billion for the others. However the range of natural abundance for the isotopes is such that the standard abundance can only be given to about ±0.001% (see table). The calculation is

Ar(Si) = (27.97693 × 0.922297) + (28.97649 × 0.046832) + (29.97377 × 0.030872) = 28.0854

The estimation of the uncertainty is complicated,[16] especially as the sample distribution is not necessarily symmetrical: the IUPAC standard relative atomic masses are quoted with estimated symmetrical uncertainties,[17] and the value for silicon is 28.0855(3). The relative standard uncertainty in this value is 1×10–5 or 10 ppm.

Apart from this uncertainty by measurement, some elements have variation over sources. That is, different sources (ocean water, rocks) have a different radioactive history, and so different isotopic composition. To reflect this natural variability, in 2010 IUPAC made the decision to list the standard relative atomic masses of 12 elements as an interval rather than a fixed number.[18]

See also

References

  1. ↑ Meija, J.; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure Appl. Chem. 88 (3): 265–91. doi:10.1515/pac-2015-0305.
  2. 1 2 International Union of Pure and Applied Chemistry (1980). "Atomic Weights of the Elements 1979". Pure Appl. Chem. 52 (10): 2349–84. doi:10.1351/pac198052102349.
  3. ↑ International Union of Pure and Applied Chemistry (1993). Quantities, Units and Symbols in Physical Chemistry, 2nd edition, Oxford: Blackwell Science. ISBN 0-632-03583-8. p. 41. Electronic version.
  4. ↑ Definition of element sample
  5. ↑ de Bièvre, P.; Peiser, H. S. (1992). "'Atomic Weight'—The Name, Its History, Definition, and Units". Pure Appl. Chem. 64 (10): 1535–43. doi:10.1351/pac199264101535.
  6. ↑ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "relative atomic mass".
  7. ↑ IUPAC Gold Book - atomic weight
  8. ↑ IUPAC Gold Book - relative atomic mass (atomic weight), A r
  9. ↑ IUPAC Gold Book - unified atomic mass unit
  10. ↑ Greenwood, Norman N.; Earnshaw, Alan (1984). Chemistry of the Elements. Oxford: Pergamon Press. pp. 21, 160. ISBN 0-08-022057-6.
  11. ↑ International Union of Pure and Applied Chemistry (2003). "Atomic Weights of the Elements: Review 2000". Pure Appl. Chem. 75 (6): 683–800. doi:10.1351/pac200375060683.
  12. ↑ National Institute of Standards and Technology. Atomic Weights and Isotopic Compositions for All Elements.
  13. 1 2 Wapstra, A.H.; Audi, G.; Thibault, C. (2003), The AME2003 Atomic Mass Evaluation (Online ed.), National Nuclear Data Center. Based on:
  14. 1 2 Rosman, K. J. R.; Taylor, P. D. P. (1998), "Isotopic Compositions of the Elements 1997" (PDF), Pure and Applied Chemistry, 70 (1): 217–35, doi:10.1351/pac199870010217
  15. ↑ Coplen, T. B.; et al. (2002), "Isotopic Abundance Variations of Selected Elements" (PDF), Pure and Applied Chemistry, 74 (10): 1987–2017, doi:10.1351/pac200274101987
  16. ↑ Meija, Juris; Mester, Zoltán (2008). "Uncertainty propagation of atomic weight measurement results". Metrologia. 45: 53–62. Bibcode:2008Metro..45...53M. doi:10.1088/0026-1394/45/1/008.
  17. ↑ Holden, Norman E. (2004). "Atomic Weights and the International Committee—A Historical Review". Chemistry International. 26 (1): 4–7.
  18. ↑ IUPAC - International Union of Pure and Applied Chemistry: Atomic Weights of Ten Chemical Elements About to Change
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