Definition · Plain-language
Isotope
An isotope is one of two or more forms of an element that share the same number of protons but differ in the number of neutrons.
The step most authors miss
Doing CRediT right? Don’t stop at the statement.
A CRediT statement credits you inside one paper. The recognition CRediT was built for happens when those roles are tied to you, persistently. Sign in with your ORCID — free — and claim your CRediT contributions on casrai.org, the home of the standard. They become a verified, portable part of your identity, not a line that disappears into one PDF.
Free: claim your contributions, then export a journal-ready CRediT statement, schema.org structured data, JATS XML, CSV or BibTeX — and preview your public profile. A membership publishes that profile publicly and verifies the journals you serve.
Same element, different neutrons
Every atom of a given element has the same number of protons — that proton count, the atomic number, is what makes it that element. Isotopes are atoms of the same element that differ only in how many neutrons sit in the nucleus. Since neutrons add mass but no charge, isotopes of an element have different mass numbers but identical chemical behaviour. Carbon, for instance, always has six protons, but its isotopes carbon-12, carbon-13 and carbon-14 have six, seven and eight neutrons respectively.
Why chemistry stays the same
Chemical behaviour is governed by electrons, especially the outer valence electrons, and the number of electrons in a neutral atom equals the number of protons. Because isotopes share the same proton count, they have the same electron arrangement and therefore the same chemistry — they form the same compounds and undergo the same reactions. The extra neutrons change only the mass, which can have subtle physical effects (heavier isotopes react slightly more slowly) but does not change which reactions occur. This is why the relative atomic mass on the periodic table is a weighted average across an element’s natural isotopes.
Stable and radioactive isotopes
Some isotopes have nuclei that are stable indefinitely, while others are unstable and break down over time, emitting radiation — these are radioactive isotopes, or radioisotopes. The rate of decay is described by the half-life, the time for half a sample to decay. Radioisotopes have many uses: carbon-14 dating estimates the age of once-living material, medical isotopes image the body or treat disease, and others power instruments. Whether an isotope is stable or radioactive depends on the balance of protons and neutrons in its nucleus.
Key facts
At a glance
- Definition: atoms of one element with the same protons but different neutrons
- Same: atomic number (protons) and chemistry
- Different: number of neutrons, and therefore mass number
- Example: carbon-12, carbon-13 and carbon-14
- Two kinds: stable isotopes and radioactive isotopes
- Periodic mass: a weighted average across an element’s natural isotopes
Common misconceptions
What people often get wrong
Often heard: Isotopes of an element have different chemical properties.
Actually: Isotopes share the same number of protons and electrons, so their chemistry is essentially the same. Only their mass differs, which has minor physical effects but does not change the reactions they undergo.
Often heard: All isotopes are radioactive.
Actually: Many isotopes are perfectly stable and do not decay. Only some isotopes are radioactive, breaking down over time and emitting radiation; most elements have at least one stable isotope.
Often heard: A different number of protons makes an isotope of the same element.
Actually: Changing the number of protons changes the element itself. Isotopes differ only in neutron number; the proton count, which defines the element, is always the same.
Going deeper
