The terms "relative atomic mass" and "average atomic mass" are often used interchangeably, but they do have slightly different meanings.
Relative Atomic Mass: Relative atomic mass refers to the mass of an atom of a specific element relative to the mass of an atom of carbon-12, which is assigned a value of exactly 12 atomic mass units (u). It is a dimensionless quantity and is represented by a numerical value without any unit. For example, the relative atomic mass of carbon is approximately 12.01.
Relative atomic mass takes into account the different isotopes of an element and their relative abundances in nature. Since isotopes have different masses, the relative atomic mass is a weighted average of the masses of all the isotopes, with the abundances of each isotope serving as the weights. It provides a measure of the typical mass of atoms of an element found in nature.
Average Atomic Mass: Average atomic mass, also known as atomic weight, refers to the average mass of the atoms in a naturally occurring sample of the element. It is expressed in atomic mass units (u). The average atomic mass takes into account the different isotopes of an element, their masses, and their relative abundances.
The average atomic mass is determined by considering the isotopic composition of the element, which may vary in different samples. It represents the weighted average of the atomic masses of all the isotopes of an element present in a sample, with the abundances of each isotope serving as the weights. The average atomic mass is the value that is typically listed in the periodic table for each element.
In summary, relative atomic mass refers to the mass of an atom relative to carbon-12, while average atomic mass represents the weighted average of the masses of all isotopes of an element in a naturally occurring sample. Both values are important in understanding the atomic properties and behavior of elements.