Relative mass is a cornerstone of modern chemistry, allowing chemists to express the mass of an atom or molecule relative to a universal standard: one‑twelfth of a carbon‑12 atom. Because protons and neutrons are roughly 1 × 10⁻²⁷ kg and electrons are 1 × 10⁻³⁰ kg, using a simple, unitless scale makes calculations both intuitive and precise.
In this system, every proton or neutron counts as 1 unit. Electrons, being far lighter, are omitted from the calculation. Thus the relative atomic mass of an element is simply the sum of its protons and neutrons:
Relative atomic mass = number of protons + number of neutrons
For example, a hydrogen atom (1 proton, 0 neutrons) has a relative mass of 1, while a carbon‑12 atom (6 protons, 6 neutrons) has a relative mass of 12.
Periodic tables list a single value for each element, often a non‑integer, because most elements exist as a mixture of isotopes. The table value is a weighted average based on natural abundance:
Weighted average = Σ (isotope mass × abundance) / 100
For chlorine, the calculation is:
((35 × 75) + (37 × 25)) / 100 = 35.5
Thus the relative atomic mass shown for chlorine is 35.5.
When a particular isotope is referenced—such as uranium‑238—the number following the element name is its exact relative mass (238). Knowing the isotope lets you compute the mass directly by adding protons and neutrons.
Once the relative atomic masses are known, the relative molecular mass of a compound is the sum of each element’s contribution:
Relative molecular mass = Σ (number of atoms of element × relative atomic mass of that element)
Examples:
Apply the same procedure to any chemical formula.
With these principles, you can confidently determine the relative mass of any atom or molecule, whether in academic research or everyday chemistry tasks.
