Atomic Mass Calculator

Average atomic mass is the weighted mean of the isotope masses found in a sample of an element. Atoms of one element always have the same number of protons, but they may have different numbers of neutrons. Those variants are isotopes. Each isotope has its own mass and natural abundance, so the value printed on a periodic table reflects both quantities rather than simply naming the most common isotope. To calculate the average, multiply every isotope mass by its abundance, add those products, and divide by the sum of all abundances. This normalized weighted-average formula works whether abundances are entered as percentages such as 98.93 and 1.07 or fractions such as 0.9893 and 0.0107. The scales must not be mixed within one list, and each mass must have a corresponding abundance. The calculator reports the abundance total so you can quickly spot omitted isotopes or inconsistent data. Carbon provides a straightforward example. Using isotope masses 12.000 and 13.003 with abundances 98.93 and 1.07 gives an average near 12.0107 atomic mass units. Chlorine has two abundant isotopes near masses 35 and 37, producing an average around 35.45 rather than a whole number. A weighted average must fall between the smallest and largest isotope mass in the list. If it does not, recheck the pairing and numeric format. Atomic mass and mass number are related but distinct. Atomic mass is a measured isotope mass or a weighted mean and can contain decimals. Mass number A is the integer count of protons plus neutrons for a specific nuclide. Atomic number Z is the proton count and identifies the element. When optional Z and A values are supplied, this calculator reports Z protons, A minus Z neutrons, and Z electrons for a neutral atom. Ions require a separate adjustment to electron count. Average atomic masses are used throughout stoichiometry, molar-mass calculations, isotope-ratio analysis, geochemistry, environmental tracing, and mass spectrometry. Natural abundances can vary slightly by source, and radioactive or synthetic elements may not have a representative natural average. For high-precision scientific work, use current evaluated isotope masses and abundances for the actual material. For classroom and routine laboratory calculations, the weighted method shown here provides a transparent result that can be checked term by term.