Atomic Molecular Mass :

 It is known that the atomic mass of an element/compound is not the mass of one atom or molecule of the said element/compound because they occur as a mixture of isotopes. As the atomic and molecular masses are expressed on a relative scale based on a mass of ¹²C atom, all atomic and molecular masses are in fact the weighted average of the mass of these isotopes. Thus, atomic mass of an element is defined as the average relative mass of an atom of an element as compared to the mass of an atom of carbon (¹²C ) taken as 12.

Relative atomic mass and atomic weight: The relative atomic mass (A_r) of an element is the ratio of the mass of an atom of the element to one-twelfth the mass of an atom of carbon-12. Because an element in nature is usually a mixture of isotopes, the relative atomic mass is also the weighted mean of the atomic masses of all the atoms in a particular sample of the element, weighted by isotopic abundance. In this sense, relative atomic mass was once known as atomic weight. Mass number: The mass number of an isotope is the total number of nucleons (neutrons plus protons) in the nucleus of each atom of the isotope. Rounding the atomic mass of an isotope usually gives the total nucleon count. The neutron count can then be derived by subtracting the atomic number (number of protons) from the mass number.

Often an element has one predominant isotope. The actual numerical difference between the atomic mass of that main isotope and the relative atomic mass or standard atomic weight of the element can be very small, such that it does not affect most bulk calculations; but such an error can be critical when considering individual atoms. For elements with more than one common isotope, the difference between the atomic mass of the most common isotope and the relative atomic mass of the element can be as much as half a mass unit or more .The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.