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Relative atomic mass (symbol: A r; sometimes abbreviated RAM or r.a.m.), also known by the deprecated synonym atomic weight, is a dimensionless physical quantity defined as the ratio of the average mass of atoms of a chemical element in a given sample to the atomic mass constant.
As such, relative atomic mass and standard atomic weight often differ numerically from the relative isotopic mass. The atomic mass (relative isotopic mass) is defined as the mass of a single atom, which can only be one isotope (nuclide) at a time, and is not an abundance-weighted average, as in the case of relative atomic mass/atomic weight ...
Formerly called atomic/molecular weight. Example: A r (Cl) = 35.453. Both quantities depend on the nuclidic composition. relative molecular mass: M r: Ratio of the average mass per molecule or specified entity of a substance to 1/12 of the mass of an atom of the nuclide 12 C number of molecules or other elementary entities: N
Atomic weight and relative atomic mass are synonyms. The standard atomic weight is a special value of the relative atomic mass. It is defined as the "recommended values" of relative atomic masses of sources in the local environment of the Earth's crust and atmosphere as determined by the IUPAC Commission on Atomic Weights and Isotopic ...
The relative atomic mass (a weighted average, weighted by mole-fraction abundance figures) of these isotopes is the atomic weight listed for the element in the periodic table. The abundance of an isotope varies from planet to planet, and even from place to place on the Earth, but remains relatively constant in time (on a short-term scale).
When atomic mass is shown, it is usually the weighted average of naturally occurring isotopes; but if no isotopes occur naturally in significant quantities, the mass of the most stable isotope usually appears, often in parentheses. [8] In the standard periodic table, the elements are listed in order of increasing atomic number.
The mass number gives an estimate of the isotopic mass measured in atomic mass units (u). For 12 C, the isotopic mass is exactly 12, since the atomic mass unit is defined as 1/12 of the mass of 12 C. For other isotopes, the isotopic mass is usually within 0.1 u of the mass number.
The molar mass constant, usually denoted by M u, is a physical constant defined as one twelfth of the molar mass of carbon-12: M u = M(12 C)/12. [1] The molar mass of an element or compound is its relative atomic mass (atomic weight) or relative molecular mass (molecular weight or formula weight) multiplied by the molar mass constant.