Relative atomic mass

Atoms of the same element can have different masses. They all have the same number of protons, but they can have different numbers of neutrons. These different forms of an element are called isotopes.

Most elements occur as mixtures of isotopes. The relative atomic mass of an element is the weighted average mass of its atoms, compared with one-twelfth of the mass of a carbon-12 atom. It has no unit because it is a comparison of masses.

The average depends on both the mass of each isotope and its abundance: how common it is in the mixture. A more abundant isotope contributes more to the average, so the relative atomic mass is closer to its mass. It is not usually a whole number, and it does not have to equal the mass of any individual atom.

Relative atomic mass as a weighted average

Consider a simplified chlorine sample in which 75% of the atoms are $$\ce{^{35}Cl}$$ and 25% are $$\ce{^{37}Cl}$$. Using approximate relative isotope masses of 35 and 37, multiply each mass by its fraction of the sample and add the results:

The weighted average is 35.5, closer to 35 because $$\ce{^{35}Cl}$$ is more abundant. A simple mean of 35 and 37 would give 36, but that would treat the two isotopes as equally common.