Isotopes & atomic mass
The idea
Not all atoms of an element are identical. Isotopes share the same number of protons — so they are the same element with the same chemistry — but carry different numbers of neutrons, and therefore different masses. Carbon-12 and carbon-14 are both carbon; the suffix is the mass number. Since you already know protons fix identity, isotopes are simply the allowed variations in the neutron count on top of that fixed identity.
This explains an odd fact about the periodic table: the atomic mass printed there is almost never a whole number, because it is a weighted average over every naturally occurring isotope. Think of it like a course grade: each isotope's mass counts in proportion to its abundance. Multiply each isotope mass by its fractional abundance, then add. The classic mistake is averaging the isotope masses straight down the middle — the true average sits closer to the more abundant isotope, sometimes much closer, and no single atom actually has the average mass.
Worked example
Natural copper is 69.15% copper-63 (mass 62.93 amu) and 30.85% copper-65 (mass 64.93 amu). Calculate the average atomic mass of copper.
- Convert the percent abundances to decimal fractions: 0.6915 and 0.3085. Check that they cover everything: 0.6915 + 0.3085 = 1.0000, so these two isotopes account for all natural copper.
- Weight each mass by its abundance. Copper-63 contributes 62.93 × 0.6915 = 43.52 amu of the average.
- Copper-65 contributes 64.93 × 0.3085 = 20.03 amu.
- Add the contributions: 43.52 + 20.03 = 63.55 amu.
- Sanity-check: 63.55 lies between the two isotope masses but noticeably closer to 62.93, exactly as it should, because copper-63 is more than twice as abundant — and 63.55 matches the value printed on the periodic table.
Answer. The average atomic mass of copper is 63.55 amu, pulled toward the more abundant copper-63.
Check your understanding
- Why does the average atomic mass land closer to the more abundant isotope, and what would equal abundances do to it?
- Why do two isotopes of the same element behave almost identically in chemical reactions despite their different masses?
- Given an element's average atomic mass and the masses of its two isotopes, how could you work backwards to find the abundances?
- Why is the mass on the periodic table almost never a whole number, even though every individual atom has a whole-number mass number?
Build the foundations first
Isotopes & atomic mass builds on these concepts. If any feel shaky, start there.