The relative atomic mass, also known as atomic weight or atomic mass, is a weighted average of the masses of the naturally occurring isotopes of an element. Isotopes are variants of an element that have the same number of protons but different numbers of neutrons in their nuclei.
The relative atomic mass is expressed in atomic mass units (AMU) and is typically listed for each element on the periodic table. The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom.
To calculate the relative atomic mass, you need to consider the abundance of each isotope and its respective mass. The formula for calculating the relative atomic mass is:
Relative Atomic Mass = (Isotope Mass 1 × Abundance 1) + (Isotope Mass 2 × Abundance 2) + ...
The abundance of each isotope is usually given as a percentage or a decimal fraction. By multiplying the mass of each isotope by its abundance, you obtain the weighted contribution of each isotope to the overall atomic mass. Adding up these contributions gives you the relative atomic mass of the element.
It's important to note that the relative atomic mass listed on the periodic table is an average value because it considers the natural abundance of isotopes. However, for some elements, particularly those with radioactive or artificially created isotopes, the atomic mass may be expressed as a range or a single value for the most stable isotope.