isotopes and elemental abundance
/ EYE-so-tope /
An element's identity is fixed by its number of protons — six protons is always carbon. But the number of neutrons in the nucleus can vary, and atoms of the same element with different neutron counts are called isotopes. Carbon, for instance, is mostly carbon-12 (6 protons, 6 neutrons) but includes a little carbon-13 and a trace of radioactive carbon-14. They are chemically nearly identical twins that differ only in mass.
Because they have the same number of electrons and the same nuclear charge, isotopes behave almost identically in chemistry — they form the same bonds and compounds. They differ in mass and in nuclear stability: some isotopes are stable forever, while others are radioactive and decay over time. The atomic mass printed on the periodic table is not a whole number because it is the weighted average of an element's naturally occurring isotopes — chlorine's 35.45 reflects a roughly 3-to-1 mix of chlorine-35 and chlorine-37. Elemental abundance is the related idea of how much of each element exists: hydrogen and helium overwhelmingly dominate the universe (forged in the Big Bang), while the heavier elements were and still are cooked inside stars and scattered by supernovae, a process called nucleosynthesis. On Earth, oxygen and silicon dominate the crust, while iron makes up much of the core.
Isotopes and abundance reach into every corner of inorganic chemistry and beyond. Radioactive isotopes power nuclear reactors and medical imaging; the slight mass difference between isotopes lets chemists trace reaction pathways and date rocks and fossils (carbon-14 dating, uranium-lead dating); abundance explains why some elements are cheap and ubiquitous (iron, aluminium) and others vanishingly rare and precious (platinum, the heaviest synthetic elements). The honest nuance: while isotopes are chemically nearly identical, the tiny mass difference does cause measurable kinetic isotope effects, most pronounced for the lightest elements where hydrogen and deuterium can react at noticeably different rates.
Chlorine's atomic mass of 35.45 is not an averaged single atom — it is the weighted mean of about 76% chlorine-35 and 24% chlorine-37, two isotopes that react identically.
The non-whole atomic mass is an abundance-weighted average of isotopes.
Isotopes share chemistry but not nuclear stability — changing neutron count never changes the element. Do not confuse isotopes (same protons, different neutrons) with ions (same protons, different electrons) or with allotropes (same element, different bonding).