Concept:The 'atomic mass' found on the periodic table is not the mass of a single atom, but rather a weighted average of all naturally occurring variants (isotopes) of that element.
Formula:$$ \text{Atomic Mass} = \sum (\text{Mass of Isotope} \times \text{Fractional Abundance}) $$
Solution:- Because elements exist in nature as a mixture of isotopes (atoms with the same number of protons but different neutrons), the overall atomic mass is inherently a weighted average.
- Therefore, the final published atomic mass depends strictly on two factors: the exact mass of each individual isotope, and how common (abundant) each isotope is in nature.
Why other options are incorrect:- Atomic number: Only tells you the number of protons, which doesn't account for the neutrons that contribute heavily to mass.
- Number of electrons: Electrons have negligible mass (\( \sim 1/1836 \) of a proton) and contribute virtually zero to the atomic mass.
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