Relative atomic mass
Relative atomic mass (symbol: Ar; sometimes abbreviated RAM or r.a.m.), also known by the
For a single given sample, the relative atomic mass of a given element is the
The more common, and more specific quantity known as standard atomic weight (Ar,standard) is an application of the relative atomic mass values obtained from many different samples. It is sometimes interpreted as the expected range of the relative atomic mass values for the atoms of a given element from all terrestrial sources, with the various sources being taken from Earth.[8] "Atomic weight" is often loosely and incorrectly used as a synonym for standard atomic weight (incorrectly because standard atomic weights are not from a single sample). Standard atomic weight is nevertheless the most widely published variant of relative atomic mass.
Additionally, the continued use of the term "atomic weight" (for any element) as opposed to "relative atomic mass" has attracted considerable controversy since at least the 1960s, mainly due to the technical difference between
Definition
Relative atomic mass is determined by the average atomic mass, or the
It is a synonym for atomic weight, though it is not to be confused with
Current definition
The prevailing IUPAC definitions (as taken from the "
- atomic weight — See: relative atomic mass[11]
and
- relative atomic mass (atomic weight) — The ratio of the average mass of the atom to the unified atomic mass unit.[12]
Here the "unified atomic mass unit" refers to 1⁄12 of the mass of an atom of 12C in its ground state.[13]
The IUPAC definition[1] 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.[14][15] Nevertheless, given the cost and difficulty of isotope analysis, it is common practice to instead substitute the tabulated values of standard atomic weights, which are ubiquitous in chemical laboratories and which are revised biennially by the IUPAC's Commission on Isotopic Abundances and Atomic Weights (CIAAW).[16]
Historical usage
Older (pre-1961) historical relative scales based on the atomic mass unit (symbol: a.m.u. or amu) used either the
Standard atomic weight
The
Also, CIAAW has published abridged (rounded) values and simplified values (for when the Earthly sources vary systematically).
Other measures of the mass of atoms
Atomic mass (ma) is the mass of a single atom. It defines the mass 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 below. A convenient unit of mass for atomic mass is the dalton (Da), which is also called the unified atomic mass unit (u).
The relative isotopic mass is the ratio of the mass of a single atom to the
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[17][18] for virtually all non-radioactive nuclides, but isotopic compositions are both harder to measure to high precision and more subject to variation between samples.[19][20] 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[18] | Abundance[19] | |
---|---|---|---|
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 as follows:
- Ar(Si) = (27.97693 × 0.922297) + (28.97649 × 0.046832) + (29.97377 × 0.030872) = 28.0854
The estimation of the
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, the IUPAC made the decision in 2010 to list the standard relative atomic masses of 10 elements as an interval rather than a fixed number.[23]
See also
- International Union of Pure and Applied Chemistry (IUPAC)
- Commission on Isotopic Abundances and Atomic Weights (CIAAW)
Notes
- reciprocal moles by definition, whereas previously it had to be determined experimentally and thus had an uncertainty.[3]: 134
- ^ Immediately following the 2019 redefinition, M(12C) was equal to 12.0000000000(54) g/mol, corresponding to a relative standard uncertainty[6] of 4.5 × 10-10. This uncertainty was “inherited” from the relative standard uncertainty that the product h NA had immediately prior to the redefinition: also 4.5 × 10-10. (Here h is the Planck constant. Following the redefinition, the product h NA has an exact value by definition.)[7]: 143 Conversely, immediately prior to the redefinition, the Avogadro constant NA had a measured value of 6.022140758(62) × 1023 reciprocal moles, corresponding to a relative standard uncertainty of 1.0 × 10-8. Note that immediately prior to the redefinition, the product h NA was known far more precisely than either h or NA individually[7]: 139 ).
References
- ^ .
- ISBN 0-632-03583-8. p. 41. Electronic version.
- ^ ISBN 978-92-822-2272-0, archivedfrom the original on 18 October 2021
- ^ "2018 CODATA Value: molar mass of carbon-12". The NIST Reference on Constants, Units, and Uncertainty. NIST. 20 May 2019. Retrieved 2023-08-30.
- PMC 9890581.
- NIST. Archivedfrom the original on 24 July 2023. Retrieved 30 August 2023.
- ^ .
- ^ Definition of element sample
- .
- ^ IUPAC Gold Book - atomic weight
- ^ IUPAC Gold Book - relative atomic mass (atomic weight), A r
- ^ IUPAC Gold Book - unified atomic mass unit
- ISBN 978-0-08-022057-4.
- S2CID 96800435.
- ^ IUPAC Gold Book - standard atomic weights
- ^ National Institute of Standards and Technology. Atomic Weights and Isotopic Compositions for All Elements.
- ^ a b Wapstra, A.H.; Audi, G.; Thibault, C. (2003), The AME2003 Atomic Mass Evaluation (Online ed.), National Nuclear Data Center. Based on:
- Wapstra, A.H.; Audi, G.; Thibault, C. (2003), "The AME2003 atomic mass evaluation (I)",
- Audi, G.; Wapstra, A.H.; Thibault, C. (2003), "The AME2003 atomic mass evaluation (II)",
- ^
- S2CID 122229901.
- ^ Holden, Norman E. (2004). "Atomic Weights and the International Committee—A Historical Review". Chemistry International. 26 (1): 4–7.
- ^ "Changes to the Periodic Table". Archived from the original on 2019-07-15.
Further reading
- Possolo, Antonio; van der Veen, Adriaan M.H.; Meija, Juris; Brynn Hibbert, D. (2018-01-04). "Interpreting and propagating the uncertainty of the standard atomic weights (IUPAC Technical Report)". Pure and Applied Chemistry. 90 (2): 395–424. S2CID 145931362. Retrieved 2019-02-08.