We use relative atomic mass instead of actual weight when calculating percent composition by weight for a compound because relative atomic mass provides a standardized value that allows for easier comparison and calculation across different elements.
Relative atomic mass, also known as atomic weight, is the average mass of an element's isotopes, taking into account their natural abundance. It is expressed relative to the mass of a carbon-12 atom, which is assigned a value of exactly 12 atomic mass units (amu). For example, the relative atomic mass of oxygen is approximately 16 amu.
When calculating the percent composition by weight of a compound, we determine the proportion of each element's mass relative to the total mass of the compound. Using relative atomic mass allows us to calculate this proportion accurately and consistently.
Actual weight, on the other hand, can vary depending on the isotopic composition of an element. Isotopes of an element have different masses due to the presence of different numbers of neutrons. However, the relative abundances of isotopes in naturally occurring samples are relatively constant. By using the average mass of isotopes, as represented by the relative atomic mass, we can ensure more accurate and meaningful calculations of percent composition by weight for compounds.
Additionally, using relative atomic mass simplifies the calculation process. The relative atomic masses of elements are readily available on the periodic table, so we can easily access the necessary values for calculations without the need for extensive experimentation or precise measurement of actual weights for individual isotopes.
In summary, relative atomic mass provides a standardized and convenient way to express the masses of elements, enabling consistent and accurate calculations of percent composition by weight for compounds.