Atomic mass is the total mass of particles of matter in an atom, which includes the masses of protons, neutrons, and electrons in an atom added together. The atomic mass of an element is expressed in atomic mass units (amu) or daltons (Da), which are equivalent. The atomic mass unit is defined as 1/12 of the mass of a free carbon-12 atom at rest in its ground state. The observed atomic mass is slightly less than the sum of the mass of the protons, neutrons, and electrons that make up the atom. The difference, called the mass defect, is accounted for during the combination of these particles by conversion into binding energy, according to an equation in which the energy (E) released equals the product of the mass (m) consumed and the square of the velocity of light in vacuum (c); thus, E = mc^2.

It is important to note that the atomic mass of an isotope and the relative isotopic mass refer to a certain specific isotope of an element. Since substances are usually not isotopically pure, it is convenient to use the elemental atomic mass, which is the average (mean) atomic mass of an element, weighted by the abundance of the isotopes. The relative atomic mass is an average of the atomic masses of all the different isotopes in a sample, with each isotopes contribution to the average determined by how big a fraction of the sample it makes up. The relative atomic masses given in periodic table entries are calculated for all the naturally occurring isotopes of each element, weighted by the abundance of those isotopes on earth.