Isotopes & Atomic Mass

Why the periodic table's atomic masses are decimals — and never whole numbers.

High schoolIntro Gen Chem
⏱️ About 16 min

Look at carbon on the periodic table: its atomic mass is 12.011, not a tidy 12. That stray decimal is a clue — not every carbon atom is identical. Some carry extra neutrons. Those are isotopes, and they explain every decimal mass on the table.

💡
The big idea: Isotopes are atoms of the same element (same number of protons) that differ in their number of neutrons. They behave identically in chemistry but differ in mass — and the periodic table's atomic mass is the weighted average across all of an element's natural isotopes.
🎯 By the end, you'll be able to
  • Define an isotope and write isotope notation (e.g. carbon-14, ¹⁴C)
  • Explain why isotopes share chemistry but differ in mass
  • Calculate an element's average atomic mass as a weighted average
📎 Helpful to know first
  • Protons, neutrons & electrons

Same element, different mass

Recall that the number of protons defines the element. Neutrons don't change what element you have — so you can add or remove neutrons and still have the same element, just a heavier or lighter version. These variants are isotopes.

Carbon always has 6 protons. But carbon-12 has 6 neutrons, carbon-13 has 7, and carbon-14 has 8. All three are carbon; all three do carbon chemistry identically (chemistry is run by electrons, and neutral carbon always has 6). They simply weigh different amounts.

🔑 Isotope notation
We name an isotope by its mass number: carbon-14 or ¹⁴C. The 14 is protons + neutrons, so neutrons = 14 − 6 = 8. The element symbol already tells you the proton count.
e⁻ e⁻ n=1 (2 e⁻) e⁻ e⁻ e⁻ e⁻ n=2 (4 e⁻) 6 p⁺ 6 n Carbon · C · Z=6 · mass number 12
Carbon-12: 6 protons, 6 neutrons. Generated from Z = 6, A = 12.
e⁻ e⁻ n=1 (2 e⁻) e⁻ e⁻ e⁻ e⁻ n=2 (4 e⁻) 6 p⁺ 8 n Carbon · C · Z=6 · mass number 14
Carbon-14: same 6 protons (still carbon, same electron arrangement) — but 8 neutrons. Only the nucleus changed.

Where the decimal masses come from

Elements occur in nature as a fixed mix of isotopes. Chlorine, for instance, is about 75.8% chlorine-35 and 24.2% chlorine-37. The atomic mass on the periodic table is the weighted average of those isotope masses — weighted by how common each one is.

\[ \text{average mass} = \sum (\text{fraction}_i \times \text{mass}_i) \]
Multiply each isotope's mass by its natural fraction, then add.
📝 Worked example: Chlorine is 75.8% chlorine-35 (mass 35) and 24.2% chlorine-37 (mass 37). Find its average atomic mass.
  1. Convert percentages to fractions: 0.758 and 0.242.
  2. Weight each mass: (0.758 × 35) + (0.242 × 37) = 26.53 + 8.95.
  3. Add: 26.53 + 8.95 = 35.5 (matches the 35.45 on the table).
✓ ≈ 35.5 — closer to 35 because chlorine-35 is the more common isotope.
✏️ Practice: A metal has two isotopes: mass-63 at 69% abundance and mass-65 at 31% abundance. What is its average atomic mass? (Round to 2 decimals.)
u
Solution
  1. Fractions: 0.69 and 0.31.
  2. (0.69 × 63) + (0.31 × 65) = 43.47 + 20.15.
  3. = 63.62 u. (This is copper — check the table!)

Check your understanding

1. Two atoms are isotopes of the same element. What must be the same?
Isotopes share the number of protons (that's what makes them the same element). They differ in neutrons, and therefore in mass number.
2. Why is chlorine's atomic mass 35.45 rather than a whole number?
The table lists the weighted average across natural isotopes (mostly Cl-35 with some Cl-37), which lands between whole numbers.
🔎 Related lessons
Previous: Inside the Atom Electron Configuration
✅ Key takeaways
  • Isotopes = same element (same protons), different numbers of neutrons.
  • Isotope notation uses the mass number: carbon-14 has 14 − 6 = 8 neutrons.
  • Isotopes are chemically identical (electrons unchanged) but differ in mass.
  • Average atomic mass = weighted average of isotope masses by natural abundance.
➡️ Neutrons set an atom's mass; electrons set its chemistry. So next we follow the electrons — exactly how they arrange themselves is the key to the whole periodic table.
Want to test yourself on this? Try the Chemistry practice test →