The discussion centers on the atomic weight of carbon, specifically the value of 12.0107 g/mole. This figure reflects the average atomic mass of all naturally occurring carbon isotopes, rather than solely the Carbon-12 isotope, which is defined as having a mass of exactly 12 g/mole. The presence of heavier isotopes, such as Carbon-13 and Carbon-14, contributes to this average. Participants clarify that the atomic mass unit (a.m.u.) is a unit of mass for nucleons, while relative atomic mass (r.a.m.) considers the proportions of isotopes in nature, leading to the confusion in terminology. The discussion highlights the distinction between the definition of the mole in relation to Carbon-12 and the average atomic mass of carbon as found in nature.
