The terms "relative isotopic mass" and "average atomic mass" both relate to the mass of atoms, particularly in the context of isotopes, but they have slightly different meanings.
Relative Isotopic Mass: The relative isotopic mass refers to the mass of a specific isotope of an element relative to a standard reference value. It is a measure of the mass of a single atom of a particular isotope. The relative isotopic mass is typically expressed in atomic mass units (u) or unified atomic mass units (u), which are defined relative to the mass of a carbon-12 atom.
For example, the relative isotopic mass of carbon-12 is defined as exactly 12 atomic mass units. This means that a carbon-12 atom has a relative isotopic mass of 12 u.
Average Atomic Mass: The average atomic mass, also known as the atomic weight or atomic mass, is the weighted average of the masses of all naturally occurring isotopes of an element. It takes into account the relative abundance of each isotope in nature.
Since many elements have multiple isotopes with different masses, the average atomic mass provides a more representative value for the mass of an atom of that element. The average atomic mass is also expressed in atomic mass units (u) or unified atomic mass units (u).
For example, carbon has two stable isotopes: carbon-12 and carbon-13. Carbon-12 is more abundant in nature, making up about 98.9% of naturally occurring carbon, while carbon-13 accounts for the remaining 1.1%. The average atomic mass of carbon is calculated by multiplying the relative isotopic masses of each isotope by their respective abundances and summing them. In this case, the average atomic mass of carbon is approximately 12.01 u.
In summary, the relative isotopic mass refers to the mass of a specific isotope of an element, while the average atomic mass represents the weighted average of the masses of all naturally occurring isotopes of that element. The average atomic mass takes into account the relative abundance of each isotope in nature.