* Isotopes: Most elements exist in nature as a mixture of different isotopes. Isotopes are atoms of the same element that have the same number of protons (defining the element) but different numbers of neutrons. This difference in neutrons results in different masses for each isotope.
* Weighted Average: The atomic mass listed on the periodic table is a weighted average of the masses of all naturally occurring isotopes of that element. The weighting is based on the relative abundance of each isotope.
Example: Carbon
* Carbon-12 (6 protons, 6 neutrons) has a mass of 12 atomic mass units (amu).
* Carbon-13 (6 protons, 7 neutrons) has a mass of 13 amu.
* Carbon-14 (6 protons, 8 neutrons) has a mass of 14 amu (very rare).
Carbon-12 is the most abundant isotope, making up about 98.9% of natural carbon. Carbon-13 is about 1.1% abundant. The average atomic mass of carbon is calculated as:
(0.989 x 12 amu) + (0.011 x 13 amu) ≈ 12.01 amu
Key Point: The atomic mass of an element is a weighted average, not the mass of a single atom. Since the isotopes have different masses and abundances, the resulting average often appears as a fraction.
