Isotopes and Nuclide Notation

Isotopes — atoms of the same element with the same number of protons but different numbers of neutrons.

Because the proton number is unchanged, isotopes of an element behave identically chemically — they have the same number of electrons arranged the same way. What differs is their mass, and whether they are stable.

Nuclide notation writes both numbers beside the symbol, with the mass number above and the atomic number below:

²³⁸₉₂U uranium-238: 92 protons, 238 − 92 = 146 neutrons

²³⁵₉₂U uranium-235: 92 protons, 143 neutrons

Both are uranium because both have 92 protons. Only the neutron count differs, and that difference is enough to make one usable as a reactor fuel and the other not.

Three sets worth knowing:

• Hydrogen: ¹₁H (protium, no neutrons), ²₁H (deuterium, 1 neutron), ³₁H (tritium, 2 neutrons, radioactive).
• Carbon: ¹²₆C (stable, the common form), ¹⁴₆C (radioactive, used for dating).
• Uranium: ²³⁵₉₂U (fissile) and ²³⁸₉₂U (not readily fissile).

Why some isotopes are radioactive. The nucleus is held together by a short-range attractive force competing with the electrostatic repulsion between protons. An unfavourable balance of protons and neutrons leaves the nucleus unstable, and it decays by emitting radiation to reach a more stable arrangement.

Radioactive decay is random and spontaneous. Random means you cannot predict which nucleus will decay next, or when a particular one will. Spontaneous means nothing triggers it: heating, cooling, chemical reactions and pressure make no difference. You can only say what fraction of a large number will decay in a given time, which is why half-life is a statistical quantity.

Example: carbon-12 and carbon-14 are chemically indistinguishable — a living organism takes both up without discriminating. That is precisely what makes radiocarbon dating work, since the ratio in a living thing matches the atmosphere, and only starts to fall once it dies.