Isotopic and Relative Atomic Masses:

At this stage it is significant to distinguish among mass number (A), isotopic mass and atomic mass and associative atomic mass (atomic weight) that even by look same but are different. Here Mass number (A) is always an integer quantity represented through the sum of its atomic number (Z) and number of neutrons (N), A = Z + N. The isotopic or nuclear mass (m_a) refers to the mass of an atom expressed as atomic mass unit (u) also called Dalton.

1 u = 1.66 × 10^-24g    = 1.66× 10^-27 kg. = 1 Dalton = 931.5 MeV.

Isotopic mass of a nuclide may be calculated by multiplying the mass of proton (m_Z = 1.673 × 10^-24g = 1.007823 amu) and the mass of neutron (m_N = 1.675× 10^-24g = 1.008665 amu) with their respective numbers in a nucleus. Thus, mass of ²³Na having 11 protons and 12 neutrons may be calculated as;

m_Na = 11× 1.007823 + 12× 1.008665 = 23.190033 amu

The relative atomic mass (also called atomic weight of an element) refers to the weighted average of the masses of all the naturally occurring isotopes in that element. For example, carbon has two naturally occurring isotopes ¹²C and ¹³C with abundances of 98.89 and 1.11% respectively. Thus, its atomic mass will be:

Relative atomic mass (A) of C = (12.000000× 98.89 + 13.003355×1.10 /100) = 12.011

Another term dealing with the mass of nuclides is the mass excess, frequently called the mass defect (?m).  It represents the difference among the mass of nuclide and the mass number or m_a- A. Thus, for the case of ¹³C, ?m = 0.003355 amu.