The relative atomic mass, also known as atomic weight or atomic mass, is a dimensionless quantity that represents the average mass of an atom of an element relative to 1/12th the mass of a carbon-12 atom. It is a weighted average of the masses of the naturally occurring isotopes of an element.
The relative atomic mass is expressed in atomic mass units (amu) or unified atomic mass units (u), where 1 atomic mass unit is defined as 1/12th the mass of a carbon-12 atom. The value of the relative atomic mass for an element can be found on the periodic table.
Since most elements exist as a mixture of isotopes, which are atoms of the same element with different numbers of neutrons, the relative atomic mass takes into account the abundance of each isotope in nature. The mass of each isotope is multiplied by its relative abundance, and the resulting values are summed to obtain the average mass.
For example, carbon has two major isotopes: carbon-12 and carbon-13. Carbon-12 is the most abundant isotope, while carbon-13 is less common. The relative atomic mass of carbon is approximately 12.01 amu, which is closer to the mass of carbon-12 because it is more abundant.
The relative atomic mass is useful in various scientific fields, such as chemistry and physics, as it provides a standard reference for comparing the masses of different atoms and calculating molar masses of compounds.