The atomic weight of an element is also referred to as the relative atomic mass because it is a relative measurement that compares the mass of an atom of a particular element to a reference standard.
The concept of atomic weight emerged from the early understanding of atoms and their constituents. Initially, scientists considered atoms as indivisible particles, but it was later discovered that atoms consist of subatomic particles, namely protons, neutrons, and electrons. Protons and neutrons contribute to the majority of an atom's mass, while electrons have a negligible mass in comparison.
The atomic weight is determined by considering the average mass of all the naturally occurring isotopes of an element, taking into account their relative abundance. Isotopes are variants of an element that have the same number of protons but differ in the number of neutrons. Since different isotopes of an element have different masses, the atomic weight accounts for this variation by calculating a weighted average.
To express the atomic weight in a standardized and comparable manner, scientists use a reference standard called the unified atomic mass unit (u or Da). The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom, which is assigned a mass of exactly 12 atomic mass units.
By comparing the mass of an atom of a particular element to this reference standard, the atomic weight or relative atomic mass is obtained. It provides a relative measure of an atom's mass compared to other atoms. For example, if an atom has an atomic weight of 16, it means that it is approximately 16 times heavier than one-twelfth the mass of a carbon-12 atom.
The term "relative" emphasizes that the atomic weight is a comparative value rather than an absolute measure of an atom's mass. It allows scientists to compare the masses of different atoms and elements in a meaningful and consistent way, facilitating various calculations and chemical studies.