Relative atomic mass

The relative atomic mass (Ar) is the average mass of one atom of an element compared with one-twelfth of the mass of one carbon-12 atom. It is a ratio, so it carries no unit. Malay: jisim atom relatif · Chinese: 相对原子质量.

Because an atom is far too small to weigh directly, chemists compare its mass against an agreed standard rather than measuring it in grams. That standard is the carbon-12 atom, whose mass is defined as exactly 12 units, so one-twelfth of it is the yardstick against which every other atom is measured. The word average matters: most elements exist as a mixture of isotopes, so the Ar is the mean of the isotopic masses weighted by how common each isotope is.

Example. Chlorine exists as two isotopes, and their weighted average gives a relative atomic mass of 35.5. The value is not a whole number precisely because it is an average of the isotopes present, not the mass of a single atom.

Confusion to avoid. Do not confuse relative atomic mass with relative molecular mass. The Ar describes a single atom of an element, while the Mr adds together the Ar values of all the atoms in a molecule. Also remember that Ar has no unit, unlike molar mass, which is measured in grams per mole even though it shares the same number.

In this chapter the relative atomic mass is the starting point for almost every calculation: you add Ar values to build a relative molecular or formula mass, and you use it as the molar mass in grams when converting between mass and moles. Getting the definition exact, including the comparison with carbon-12, earns a straightforward Paper 2 mark.