Isotopes

Atoms of the same element can be slightly heavier or lighter. These versions are called isotopes.

Key idea

Definition: Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

Same protons means the same element. Different neutrons change the mass.

Why chemistry stays the same

Isotopes of an element have the same number of electrons, so they have the same electronic configuration. This means they have the same chemical properties (they react in the same way).

What can be different

• Mass and density
• Melting/boiling points (slight changes)
• Rate of diffusion (heavier isotopes diffuse more slowly)
• Some isotopes are radioactive (unstable); others are stable

Notation you must read

Mass number $A$ = protons + neutrons. Atomic number $Z$ = protons.

We write isotopes as $_{Z}^{A}\mathrm{X}$. Examples: $_{6}^{12}\mathrm{C}$ (carbon-12), $_{6}^{14}\mathrm{C}$ (carbon-14), $_{17}^{35}\mathrm{Cl}$ and $_{17}^{37}\mathrm{Cl}$.

Analogy: think of the same model of car carrying different numbers of passengers. It is still the same car (same protons), just heavier or lighter (different neutrons).

Relative atomic mass from isotopes

The relative atomic mass $A_r$ on the Periodic Table is a weighted average of an element’s isotopes.

$$A_r=\frac{\sum (\text{isotopic mass} \times \text{abundance})}{\text{total abundance}}$$

Real-world uses

• Carbon-14 is used to date ancient objects
• Iodine-131 is used as a medical tracer

Common misconceptions

• “Different neutrons = different element” — false. Element is set by protons.
• “Isotopes react differently” — false. Same electrons, so same chemistry.
• “Periodic table masses should be whole numbers” — they are averages, so often not whole.