Relative Atomic Mass
A comparison scale that lets us work with atoms too small to weigh one by one
1. Big idea and learning goals
Individual atoms are extraordinarily light. Writing the mass of one atom in kilograms would give a long string of zeros that is awkward for ordinary chemistry. Chemists therefore compare atomic masses with one agreed reference. The result is the relative atomic mass, written $A_r$. It tells us how heavy an atom of an element is on a common scale.
By the end, you can
- explain why a relative scale is useful and name the carbon-12 reference;
- read an $A_r$ value from a periodic table and state that it has no unit;
- compare atoms using $A_r$ and calculate simple totals from atom counts;
- distinguish atomic number $Z$ from relative atomic mass $A_r$;
- interpret a decimal table value without imagining a fractional atom.
2. Building the carbon-12 scale
A scale needs a fixed reference. Chemists chose one atom of the isotope carbon-12 and defined its mass as exactly 12 parts. One part is therefore one twelfth of the mass of a carbon-12 atom. The relative atomic mass of an element compares the average mass of its atoms with that one-part reference:
If $A_r(\mathrm{O})=16.00$, an oxygen atom is about 16 reference parts in mass. If $A_r(\mathrm{Fe})=55.85$, an iron atom is about 55.85 reference parts. This is a ratio, so $A_r$ is dimensionless: do not attach g, kg, or u to an $A_r$ value.