An isotope is the same chemical element, but with a different mass because it has a different number of neutrons in its nucleus.

Core idea in one line

  • All atoms of an element have the same number of protons , but isotopes of that element have different numbers of neutrons , so their mass changes while their chemical behavior stays almost the same.

“Standard form” vs isotope

When people say the “standard form” of an element, they usually mean:

  • The most common (most abundant) isotope in nature, e.g. carbon‑12 for carbon or oxygen‑16 for oxygen.
  • The average atomic mass on the periodic table, which is a weighted average of all naturally occurring isotopes (so it is not a single atom, but an average over many isotopes).

An isotope is:

  • One specific version of that element, with a fixed number of protons and neutrons, like carbon‑12 (6 protons, 6 neutrons) or carbon‑14 (6 protons, 8 neutrons).
  • Written with a mass number: for example, ,^{12}\text{C}, (carbon‑12) or ,^{14}\text{C},.

What stays the same, what changes

Same for all isotopes of an element:

  • Number of protons (defines the element identity).
  • Number of electrons in a neutral atom , so their chemical reactions are almost identical.

Different for isotopes:

  • Number of neutrons.
  • Mass number (protons + neutrons), so each isotope has a different atomic mass.
  • Some isotopes are stable , others are radioactive (they decay over time, like carbon‑14 or tritium).

Quick example story: carbon

  • Carbon‑12: 6 protons, 6 neutrons, stable, and it makes up almost all of the carbon around you.
  • Carbon‑14: 6 protons, 8 neutrons, radioactive, used in carbon dating but still behaves chemically like other carbon when forming compounds like carbon dioxide.

Both are carbon (same element), but they are different isotopes of carbon with different masses and nuclear stability. Information gathered from public forums or data available on the internet and portrayed here.