RELATIVE ATOMIC MASS (symbol: Ar), or ATOMIC WEIGHT, 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. Since both values in the ratio are expressed in the same unit (u), the resulting value is dimensionless; hence the value is relative.
For one sample, its relative atomic mass (atomic weight) 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. 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 is an application of this relative atomic mass values from different samples. It is expected range of the relative atomic mass values, with the various sources being terrestrial (taken from Earth ). These standard atomic weights are what chemists loosely and so incorrectly call "atomic weights" (incorrectly, because they are not from a single sample). They are the most published form of the relative atomic mass, because they are not sample-specific, but cover a broad range of expected Earth samples.
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. 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.
* 1 Definition
* 1.1 Current definition * 1.2 Historical amu
* 2 Standard atomic weight * 3 Other definitions of the mass of atoms * 4 Determination of relative atomic mass * 5 See also * 6 References * 7 External links
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. 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 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 these will have overlapping values if the relative atomic mass used is that for an element from Earth under defined conditions. However, relative atomic mass (atomic weight) covers more than standard atomic weights, and is more general 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.
and relative atomic mass (standard atomic weight) — The ratio of the average mass of the atom to the unified atomic mass unit.
Here the "unified atomic mass unit" refers to 1⁄12 of the mass of an atom of 12C in its ground state.
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
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.
STANDARD ATOMIC WEIGHT
Main article: Standard atomic weight
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
Main article: Isotope geochemistry
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 for virtually all non-radioactive nuclides, but
isotopic compositions are both harder to measure to high precision and
more subject to variation between samples. 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
ISOTOPE ATOMIC MASS ABUNDANCE
28Si 7001279769265324600♠27.97692653246(194) 92.2297(7)% 92.21–92.25%
29Si 7001289764947000000♠28.976494700(22) 4.6832(5)% 4.67–4.69%
30Si 7001299737701710000♠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 .
The estimation of the uncertainty is complicated, especially as the
sample distribution is not necessarily symmetrical: 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, in 2010
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Quantities, Units and Symbols in Physical Chemistry
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Selected Elements" (PDF),
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1987–2017, doi :10.1351/pac200274101987
* ^ Meija, Juris; Mester, Zoltán (2008). "Uncertainty propagation
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