Every element is defined by its atomic number Z — the number of protons in the nucleus. The mass number A is the total number of nucleons (protons + neutrons). These are written in the standard ˢX notation:
For example, carbon-12 is written 126C — it has 6 protons and 6 neutrons. The number of neutrons N is always A − Z.
Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A). They have identical chemical properties because their electron configurations are the same, but different nuclear properties — some isotopes are stable while others are radioactive.
Common isotopes
| Isotope | Notation | Protons | Neutrons | Stable? |
|---|---|---|---|---|
| Hydrogen-1 | 11H | 1 | 0 | ✓ Stable |
| Hydrogen-2 (deuterium) | 21H | 1 | 1 | ✓ Stable |
| Hydrogen-3 (tritium) | 31H | 1 | 2 | ✗ Radioactive |
| Carbon-12 | 126C | 6 | 6 | ✓ Stable |
| Carbon-14 | 146C | 6 | 8 | ✗ Radioactive |
| Uranium-235 | 23592U | 92 | 143 | ✗ Radioactive |
| Uranium-238 | 23892U | 92 | 146 | ✗ Radioactive |
Nuclear stability depends on the neutron-to-proton ratio. For light elements (Z < 20), stable nuclei have roughly equal numbers of protons and neutrons (N/Z ≈ 1). For heavier elements, more neutrons are needed to overcome the electrostatic repulsion between protons (N/Z ≈ 1.5 for the heaviest stable nuclei). Isotopes with too many or too few neutrons are unstable and undergo radioactive decay.
The relative atomic mass of an element is a weighted average of the masses of its naturally occurring isotopes, taking into account their relative abundances.