Isotopes and atomic mass
Isotopes are atoms of the same element with different numbers of neutrons. They have the same atomic number but different mass numbers.
Carbon-12 has 6 protons and 6 neutrons. Carbon-14 has 6 protons and 8 neutrons. Both are carbon and behave almost identically in reactions.
The atomic mass on the periodic table is a weighted average of an element’s natural isotopes. Common isotopes count for more.
average atomic mass = sum of (isotope mass × fraction abundance)
Worked example: chlorine
| Isotope | Mass (amu) | Abundance |
|---|---|---|
| Cl-35 | 34.97 | 75.8% |
| Cl-37 | 36.97 | 24.2% |
(34.97 × 0.758) + (36.97 × 0.242) = 26.51 + 8.95 = 35.45 amu
No single chlorine atom has a mass of 35.45 amu. It is the average over a large sample.
Common mistake
Averaging the isotope masses without weighting. (34.97 + 36.97) ÷ 2 = 35.97, which is wrong because chlorine-35 is about three times as common.