A chemist should use average atomic mass instead of atomic mass or atomic weight because average atomic mass takes into account the relative abundance of each isotope of an element.
Elements can exist in multiple isotopic forms, meaning they have different numbers of neutrons in their nuclei. Each isotope has a different atomic mass due to the varying number of neutrons. For example, chlorine has two isotopes, 17Cl35 and 17Cl37, with atomic masses of 35 and 37 atomic mass units (u), respectively.
In nature, these isotopes occur in different proportions. The average atomic mass of an element is the weighted average of the atomic masses of its isotopes, taking into account their relative abundances. The weight of each isotope is determined by its abundance, which can vary depending on the sample and the source of the element.
By using the average atomic mass, chemists can obtain a more accurate representation of the typical mass of atoms of a specific element. This is especially important in calculations involving chemical reactions, stoichiometry, and determining molecular masses. It allows chemists to consider the contributions of different isotopes and their relative abundance in a given sample, leading to more precise calculations and predictions.
On the other hand, atomic mass or atomic weight typically refers to the standard atomic weight, which is the average atomic mass of an element based on the naturally occurring isotopes and their abundances in the Earth's crust. It is often used as a reference value in periodic tables and general chemical calculations. However, in specific contexts or when dealing with isotopically enriched or depleted samples, the average atomic mass based on the specific isotopic composition becomes more relevant and accurate.
In summary, chemists should use average atomic mass because it considers the relative abundances of isotopes, providing a more accurate representation of the typical mass of atoms of an element in a given sample or system.