The halogens

Every halogen atom has seven outer electrons and needs exactly one more to complete a stable shell, which is why the group is so uniformly reactive — and also why it is not uniform in how reactive. Reactivity in this group means how eagerly an atom pulls in that one electron, and that pull weakens predictably as the group is descended: fluorine's outer shell sits close to a strongly felt nucleus, while iodine's sits several shells further out, shielded and less attracted. Fluorine is consequently the most powerful oxidiser known in ordinary chemistry; iodine is comparatively mild.

That gradient has a direct, testable consequence: a more reactive halogen displaces a less reactive one from its compounds. Bubble chlorine through potassium bromide solution and chlorine takes the electron bromide had, forming chloride and releasing free bromine — never the reverse. The displacement series follows the group exactly, top to bottom, which is what makes it a genuine test of relative reactivity rather than an arbitrary list to memorise.

Free halogens are diatomic — Cl₂, Br₂, I₂ — because a lone halogen atom, one electron short of stability, immediately pairs with another to share the missing electron covalently. They are essentially never found as isolated atoms in nature, and the compounds they form with metals, the halides, are generally stable ionic salts, since a metal readily gives up the electron a halogen wants. This combination — extremely reactive as an element, extremely stable once combined — is characteristic of the whole group and is the reason halogens turn up as compounds throughout the periodic table's chemistry.