Atomic Masses

An atom is extraordinarily small (diameter of the order of 10⁻¹⁰ m) and extraordinarily light (mass of the order of 10⁻²⁷ kg), so its absolute mass cannot be measured directly on any balance. Chemists solve this by using a relative scale: instead of measuring an atom's mass in kilograms, we measure it relative to a chosen standard atom.

The standard. IUPAC (the International Union of Pure and Applied Chemistry) fixed the carbon-12 (¹²C) atom in its ground state as this standard, and defined its mass as exactly 12 amu (also written 12 u). One atomic mass unit (amu, or u) is therefore defined as one-twelfth of the mass of a carbon-12 atom in its ground state:

1 amu (or 1 u) ≈ 1.6605 × 10⁻²⁷ kg

Relative atomic mass. Using this unit, the relative atomic mass (Aᵣ) of an element is defined as the ratio of the average mass of its atom to the unified atomic mass unit:

Aᵣ = (average mass of the atom) ÷ (unified atomic mass unit, 1.6605 × 10⁻²⁷ kg)

For hydrogen, whose average atomic mass is 1.6736 × 10⁻²⁷ kg, this gives Aᵣ(H) = 1.6736×10⁻²⁷ ÷ 1.6605×10⁻²⁷ ≈ 1.0078, rounded to 1.008 u.

Why "average"? Most elements occur naturally as a mixture of isotopes -- atoms of the same element with the same number of protons but different numbers of neutrons, and hence different masses. The relative atomic mass we quote for an element is therefore an average atomic mass: the weighted average of the atomic masses of all its naturally occurring isotopes, weighted by how abundant each isotope is. …