Atomic Mass
Atomic mass is the weighted average mass of an element's atoms in atomic mass units, reflecting how common each isotope is in nature.
Definition
Atomic mass is how heavy an atom of an element is on average, the decimal number under each symbol on the periodic table. It is the weighted average of the masses of an element's naturally occurring isotopes in atomic mass units (amu), where one amu is $1/12$ the mass of a carbon-12 atom, and more abundant isotopes count more toward the average. Individual isotopic masses are not integers because of the mass defect from nuclear binding energy and the small mass of the electrons. Atomic mass in amu is numerically equal to molar mass in grams per mole, which is what links the periodic table to laboratory balances.
Example
Carbon is listed at about $12.01$: most atoms weigh $12$ units and a few weigh $13$, so the average is slightly over $12$. Chlorine is about $75.8\%$ chlorine-35 and $24.2\%$ chlorine-37, giving $(0.758 \times 35) + (0.242 \times 37) \approx 35.5\text{ amu}$. Copper-63 ($62.930\text{ amu}$, $69.15\%$) and copper-65 ($64.928\text{ amu}$, $30.85\%$) average to $63.55\text{ amu}$, so a mole of copper has a mass of $63.55\text{ g}$ and contains $6.022 \times 10^{23}$ atoms.
Key Insight
No single carbon atom weighs $12.01$ units; the decimal exists because it is an average. An atomic mass near a whole number means one isotope dominates, while a value between whole numbers, like chlorine's, means two isotopes are both common. Isotope abundances vary slightly by geographic source, so IUPAC now publishes some atomic masses as intervals, a reminder that the periodic table number is a convention rather than a constant of nature.