Chemical element
Whereas the mass number simply counts the total number of neutrons and protons and is thus a natural (or whole) number, the atomic mass of a single atom is a real number giving the mass of a particular isotope (or "nuclide") of the element, expressed in atomic mass units (symbol: u). In general, the mass number of a given nuclide differs in value slightly from its atomic mass, since the mass of each proton and neutron is not exactly 1 u; since the electrons contribute a lesser share to the atomic mass as neutron number exceeds proton number; and (finally) because of the nuclear binding energy. For example, the atomic mass of chlorine-35 to five significant digits is 34.969 u and that of chlorine-37 is 36.966 u. However, the atomic mass in u of each isotope is quite close to its simple mass number (always within 1%). The only isotope whose atomic mass is exactly a natural number is C, which by definition has a mass of exactly 12 because u is defined as 1/12 of the mass of a free neutral carbon-12 atom in the ground state.
Molecular mass
In the example above the atomic mass of carbon is given as 12.011, not 12. This is because naturally occurring carbon is a mixture of the isotopes C, C and C which have relative atomic masses of 12, 13 and 14 respectively. Moreover, the proportion of the isotopes varies between samples, so 12.011 is an average value. By contrast, there is less variation in naturally occurring hydrogen so the average atomic mass is known more precisely. The precision of the molecular mass is determined by precision of the least precise atomic mass value, in this case that of carbon. In high-resolution mass spectrometry the isotopomers CH and CH are observed as distinct molecules, with molecular masses of 16 and 17, respectively. The intensity of the mass-spectrometry peaks is proportional to the isotopic abundances in the molecular species. C H H can also be observed with molecular mass of 17.