The relative atomic mass, also known as atomic weight, is a weighted average of the masses of all the naturally occurring isotopes of an element. It is expressed in atomic mass units (u) and is used to compare the masses of different atoms.
The relative atomic mass takes into account the abundance of each isotope of an element in nature and the mass of each isotope. Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei, resulting in slightly different masses.
The formula to calculate the relative atomic mass is:
Relative Atomic Mass = (Mass of Isotope 1 × Abundance of Isotope 1) + (Mass of Isotope 2 × Abundance of Isotope 2) + ...
The relative atomic mass is typically reported for elements on the periodic table. It is listed below the element's symbol and name, usually as a decimal number. For example, the relative atomic mass of carbon is approximately 12.01 u. This value takes into account the different isotopes of carbon, including carbon-12, carbon-13, and carbon-14, and their respective abundances in nature.
It's worth noting that the relative atomic mass is different from the atomic number, which represents the number of protons in an atom's nucleus and determines the element's identity.