oxoacids and oxoanions of the halogens

Bleach, swimming-pool chlorine, and the perchlorate in some fireworks all share a common backbone: a halogen atom bonded to oxygen and to an acidic O-H. These are the oxoacids of the halogens, and chlorine shows the full set in a tidy ladder. As you add oxygen atoms one at a time you climb from hypochlorous acid HOCl, to chlorous acid HClO2, to chloric acid HClO3, to perchloric acid HClO4, with chlorine's oxidation state rising +1, +3, +5, +7. Each loses a proton to give its oxoanion: hypochlorite, chlorite, chlorate, perchlorate.

Two opposite trends run along this ladder, and both come from the oxygen count. Acid strength increases sharply with more oxygens: HOCl is weak, HClO4 is one of the strongest acids known. The reason follows Pauling's rule — each extra terminal oxygen pulls electron density away and spreads the negative charge of the anion over more atoms by resonance, stabilising the conjugate base and making the proton easier to lose. The oxidizing power, however, tends to run the other way in the practical sense: the lower oxoacids and oxoanions (hypochlorite, chlorate) are the aggressive, fast oxidizers used in bleach and weedkillers, while perchlorate, though it holds chlorine in its highest +7 state, is kinetically sluggish and surprisingly unreactive at room temperature.

These compounds run through everyday and industrial chemistry: sodium hypochlorite is household bleach and water disinfectant, chlorates and perchlorates are oxidisers in matches, fireworks, and rocket propellant. They also teach the difference between thermodynamics and kinetics, since perchlorate is a powerful oxidiser on paper yet safe to handle dry. An honest caveat: bromine and especially iodine give a less complete and less stable set of oxoacids than chlorine (periodic acid, for instance, prefers a different, more hydrated structure), so the neat chlorine ladder is the cleanest case, not a universal template for the whole group.