# Why are bromine oxyanions uncommon?

Sodium hypochlorite is used in bleach, calcium hypochlorite is used in public swimming pools, ammonium perchlorate is used in solid rocket fuel.

Sodium periodate is used in the cleaving of syn-vicinal diols, and is generally useful as an oxidizing agent.

Bromates are water contaminants, and perbromates weren't even successfully synthesized until 1968, but only through radioactive decay. Though alternate methods for their production exist today, what makes bromine oxyanions so unstable when chlorine and iodine oxyanions are widely used in chemistry?

Due to the intervention of the ten 3d electrons, which are poorly shielding, the 4s and 4p electrons of the elements Ga through Kr are more tightly bound than one might expect. These elements show a distinct reluctance to take on their group valency. In $\ce{BrO4-}$, the higher-than-expected energy cost of ionisation to form Br(VII) is not fully compensated for by covalent bond formation (the strengths of which decrease linearly down the group). A selection of other observations that can, to various extents, be attributed to this are:
• the anomalous instability of $\ce{AsCl5}$ (decomposes above $-50~\mathrm{^\circ C}$) even though $\ce{PCl5}$ and $\ce{SbCl5}$ are well-known,
• the fact that nitric acid oxidises sulfur to $\ce{H2SO4}$, but selenium only to $\ce{H2SeO3}$