The terms "atomic weight" and "relative atomic mass" are sometimes used interchangeably, but they can have slightly different meanings depending on the context. Generally, the concepts are as follows:
Atomic Weight: Atomic weight refers to the average weight of an atom of an element, taking into account the different isotopes and their respective abundances. It is often represented by the symbol "A" and is expressed in atomic mass units (amu) or grams per mole (g/mol). The atomic weight considers the isotopic distribution of an element found in nature.
Relative Atomic Mass: Relative atomic mass is a similar concept to atomic weight, but it is specifically used to describe the mass of an atom of an element relative to the mass of an atom of carbon-12, which is assigned a mass of exactly 12 atomic mass units. Relative atomic masses are dimensionless quantities and are represented by the symbol "Ar" or "Mr".
When calculating masses from isotopic abundance data, it is more appropriate to use relative atomic masses. This is because relative atomic masses are based on the comparison to a specific standard (carbon-12) and allow for consistent and accurate comparisons between elements. By using relative atomic masses, one can calculate the average atomic mass of an element by taking into account the abundance of each isotope and its respective mass.
It's worth noting that the terms "atomic weight" and "relative atomic mass" have historically been used differently in different contexts, leading to some confusion and variations in usage. However, the current recommended practice is to use relative atomic mass when referring to the average mass of an atom relative to carbon-12 and its isotopic abundance data.