* Isotopes: Isotopes of an element have the same number of protons (and thus the same atomic number) but different numbers of neutrons. This means they have different atomic masses.
* Abundance: Isotopes occur in nature with varying abundances. For example, carbon-12 (6 protons, 6 neutrons) is the most abundant isotope of carbon, while carbon-13 (6 protons, 7 neutrons) and carbon-14 (6 protons, 8 neutrons) are less abundant.
* Average Atomic Mass: The average atomic mass is calculated by taking into account the masses of all the isotopes of an element and their relative abundances. This weighted average represents the typical atomic mass of the element found in nature.
Example:
* Carbon-12 has an atomic mass of 12.0000 amu and is 98.9% abundant.
* Carbon-13 has an atomic mass of 13.0034 amu and is 1.1% abundant.
* Carbon-14 is a trace isotope and has a much lower abundance.
The average atomic mass of carbon is calculated as:
(0.989 * 12.0000 amu) + (0.011 * 13.0034 amu) + (very small contribution from C-14) ≈ 12.011 amu
So, the atomic mass of carbon listed on the periodic table is 12.011 amu, reflecting the average atomic mass of all its naturally occurring isotopes.
