What is "relative atomic mass"

n. (context physics English) ratio of the atomic mass of one atom of an isotope to 1/12 (one twelfth) the mass of a Carbon-12 atom.

n. (chemistry) the ratio of the atomic mass of an element to half the atomic mass of carbon-12 [syn: atomic weight]

Relative atomic mass (symbol: A) is a dimensionless physical quantity, the ratio of the average mass of atoms of an element (from a single given sample or source) to of the mass of an atom of carbon-12 (known as the unified atomic mass unit). The relative atomic mass is a statistical term, referring to an abundance-weighted figure involving measurement of many atoms. As in all related terms, the word "relative" refers to making the figure relative to carbon-12, so that the final figure is dimensionless.

The term relative atomic mass is exactly equivalent to atomic weight, which is the older term. In technical usage, these values are sample-specific (i.e., element source-specific) when a natural element source is composed of more than one isotope. Thus, two samples of a chemical element which is naturally found as being composed of more than one isotope, collected from two substantially different sources, are expected to give slightly different relative atomic masses (atomic weights), because isotopic concentrations typically vary slightly due to the history (origin) of the source. These values differences are real and repeatable, and can be used to identify specific samples. For example, a sample of elemental carbon from volcanic methane will have a different relative atomic mass (atomic weight) than one collected from plant or animal tissues (for more, see isotope geochemistry). In short, the atomic weight (relative atomic mass) of carbon varies slightly from place to place and from source to source, a fact that can be useful. However, a typical (standard) figure also can be useful, as follows.

Both the terms relative atomic mass and atomic weight are sometimes loosely used to refer to a technically different standardized expectation value, called the standard atomic weight. This value is the mean value of atomic weights of a number of "normal samples" of the element in question. For this definition, "[a] normal sample is any reasonably possible source of the element or its compounds in commerce for industry and science and has not been subject to significant modification of isotopic composition within a geologically brief period." These standard atomic weights are published at regular intervals by the Commission on Isotopic Abundances and Atomic Weights of the International Union of Pure and Applied Chemistry (IUPAC) The "standard" values are intended as mean values that compensate for small variances in the isotopic composition of the chemical elements across a range of ordinary samples on Earth, and thus to be applicable to normal laboratory materials. However, they may not accurately reflect values from samples from unusual locations or extraterrestrial objects, which often have more widely variant isotopic compositions.

The standard atomic weights are reprinted in a wide variety of textbooks, commercial catalogues, Periodic Table wall charts etc., and in the table below. They are what chemists loosely call "atomic weights."

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. (see below). 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 "standard relative atomic mass") continues to be preferred.