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The mass number is a fundamental concept in understanding the composition of an atom. It is defined as the total number of protons and neutrons present in the nucleus of an atom. This value is always a whole number. Protons and neutrons are collectively referred to as "nucleons," and the difference in their count is what gives different isotopes of an element their unique mass numbers.
For example, consider an atom of carbon. The most common carbon isotope is carbon-12, meaning it has a mass number of 12, since it contains 6 protons and 6 neutrons. It's important to note that mass number differs from atomic mass, which accounts for isotopic distribution. This makes atomic mass a weighted average, often not an integer.

Molar mass is a critical concept when dealing with chemical reactions and substance quantification. It represents the mass of one mole of a substance, be it an element or a compound, and is expressed in grams per mole (g/mol).
Unlike mass number, which describes individual atoms, molar mass relates to a bulk quantity of substance that contains Avogadro's number of particles, approximately 6.022 脳 10^{23}. For a single atom or molecule, its molar mass in grams per mole is numerically equal to its atomic or molecular mass expressed in atomic mass units (amu).
Using carbon once again as an example, one mole of carbon-12 atoms has a molar mass of exactly 12 g/mol. This makes molar mass a direct linkage between macroscopic quantities used in laboratory measurements and microscopic entities like atoms and molecules.

Isotopes are variants of a particular chemical element. They have the same number of protons, which means they share the same atomic number, but they have different numbers of neutrons. This variance in neutrons causes isotopes to have different mass numbers and atomic masses.
For instance, hydrogen has three common isotopes: protium (with no neutrons), deuterium (one neutron), and tritium (two neutrons). Though chemically similar, isotopes can have permutations in physical properties, such as different stability and radioactivity levels.
Understanding isotopes is essential, especially in fields such as nuclear chemistry and medicine, where specific isotopes are utilized for diagnostics and treatment plans. Their presence impacts the atomic mass of elements since the atomic mass is the weighted average of all naturally occurring isotopes of an element.