Defining Atomic Weight
Atomic weight, also known as relative atomic mass, is the weighted average mass of all naturally occurring isotopes of an element. It is expressed in atomic mass units (amu) and reflects the average mass of an atom of that element, taking into account the abundance of each isotope. Unlike the mass number, which is a whole number for a specific isotope, atomic weight is usually not a whole number due to this averaging.
The Role of Isotopes
Elements naturally exist as a mixture of isotopes, which are atoms of the same element with different numbers of neutrons and thus different masses. The atomic weight calculation accounts for each isotope's individual mass and its relative abundance in nature. Isotopes that are more abundant contribute more significantly to the element's overall atomic weight, making it a representative average.
Calculating Atomic Weight: A Simple Example
Consider a hypothetical element 'X' with two isotopes: Isotope X-10 (mass = 10.00 amu, abundance = 20%) and Isotope X-11 (mass = 11.00 amu, abundance = 80%). The atomic weight would be calculated as (10.00 amu * 0.20) + (11.00 amu * 0.80) = 2.00 amu + 8.80 amu = 10.80 amu. This demonstrates how a heavier, more abundant isotope pulls the average closer to its own mass.
Importance in Chemistry
Atomic weight is fundamental in chemistry for practical applications such as stoichiometry, where it's used to convert between mass and moles of a substance. It allows chemists to accurately predict the mass of reactants and products in chemical reactions, determine empirical formulas, and understand the macroscopic behavior of substances based on their atomic composition.