BASIC CONCEPT
1 mole = 6.022 × 10^23
= Avogadro’s numberAs we have seen, Avogadro’s number serves as a conversion factor between one of something and a
mole of something. Since 12 amu is the mass of 1 carbon-12 atom while 12 g is the mass of 1 mole of
carbon-12 atoms, Avogadro’s number also helps to convert between the mass units. Specifically:
which is the conversion factor we gave above. We can now see how this is derived from (or related
to) the concept of the mole.
The atomic weight of an element is also found in the periodic table, as the number appearing below
the symbol for the element. Notice, however, that these numbers are not whole numbers, which is
odd considering that a proton and a neutron each have a mass of 1 amu and an atom can only have
a whole number of these. Furthermore, even carbon, the element with which we have set the
standards, does not have a mass of 12.000 exactly. This is due to the presence of isotopes, as
mentioned above. The masses listed in the periodic table are weighted averages that account for
the relative abundance of various isotopes. The word weighted is important: It is not simply the
average of the masses of individual isotopes, but takes into account how frequently one encounters
that isotope in a common sample of the element. There are, for example, 3 isotopes of hydrogen,
with 0, 1, and 2 neutrons, respectively. Together with the one proton that makes it hydrogen in the
first place, the mass numbers for these isotopes are 1, 2, and 3. The atomic weight of hydrogen,
however, is not simply 2 (the average of 1, 2, and 3) but about 1.008, that is, much closer to 1. This is
because the isotope with no neutrons is so much more abundant that we count it much more
heavily in calculating the average. The following example provides a more concrete illustration of
the idea.
Example: Element Q consists of three different isotopes, A, B, and C. Isotope A has an atomic