What Is an Isotope?

Isotopes are atoms of the same element that contain the same number of protons but different numbers of neutrons. Therefore, they have the same atomic number but different mass numbers.

This definition is required at both SL and HL in Structure 1.2, The nuclear atom.

The identity of an element is determined by its number of protons. For example, every carbon atom contains six protons. Carbon-12 and carbon-14 are both carbon because they each have six protons, but they contain different numbers of neutrons.

The atomic number, $Z$, equals the number of protons. The mass number, $A$, equals the total number of protons and neutrons, which are collectively called nucleons:

$$A = Z + N$$

where $N$ is the number of neutrons. Therefore:

$$N = A - Z$$

In nuclide notation, $^{A}_{Z}\mathrm{X}$, the mass number is written at the upper left and the atomic number at the lower left of the element symbol.

Because isotopes of an element have the same number of electrons when neutral, they have the same electron configuration and generally very similar chemical properties. Their masses and some physical properties differ because their neutron numbers differ.

For example, if a sample contains 75.0% carbon-12 and 25.0% carbon-13, its relative atomic mass is the weighted mean: $(0.750 \times 12) + (0.250 \times 13) = 12.25$. This value is not a mass number; it is an abundance-weighted average for the naturally occurring isotope mixture. No single carbon atom in this sample has mass number 12.25.

A common misconception is that isotopes are ions. This is incorrect: isotopes differ in neutron number, whereas ions form when atoms gain or lose electrons.

Exam technique: For a question asking you to “state” what an isotope is, include both marking points: same number of protons and different numbers of neutrons. If given isotope abundances, use a weighted mean to calculate the element’s relative atomic mass.