'''Atominė masė''', '''atominis svoris''' – santykinis dydis, rodantis, kiek kartų [[cheminis elementas|cheminio elemento]] [[atomas|atomo]] masė didesnė už standartinį atominės masės vienetą (amv). Atominis masės vienetas paskaičiuojamas taip, kad anglies-12 izotopo atominė masė būtų lygi 12.
Atomic mass
Anglies atomo masė m=1,66*10^(-27)kg
[[Atomas|Atomo]] masė būna labai artima atomo [[Branduolys (atomo)|branduolyje]] esančių [[Nuklonas|nuklonų]] ([[Protonas|protonų]] ir [[Neutronas|neutronų]]) skaičiui.
The mass number of an element, A, is the number of nucleons (protons and neutrons) in the atomic nucleus. Different isotopes of a given element are distinguished by their mass numbers, which are conventionally written as a super-index on the left hand side of the atomic symbol (e.g., 238U).
The relative atomic mass of an element is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance, relative to the atomic mass unit (u). This number may be a fraction which is not close to a whole number, due to the averaging process. On the other hand, the atomic mass of a pure isotope is quite close to its mass number. Whereas the mass number is a natural (or whole) number, the atomic mass of a single isotope is a real number which is close to a natural number. In general, it differs slightly from the mass number as the mass of the protons and neutrons is not exactly 1 u, the electrons also contribute slightly to the atomic mass, and because of the nuclear binding energy. For example, the mass of 19F is 18.9984032 u. The only exception to the atomic mass of an isotope not being a natural number is 12C, which has a mass of exactly 12, due to the definition of u (it is fixed as 1/12th of the mass of a free carbon-12 atom, exactly).
