Chlorite

The free acid, chlorous acid HClO2, is the least stable oxoacid of chlorine and has only been observed as an aqueous solution at low concentrations. Since it cannot be concentrated, it is not a commercial product. The alkali metal and alkaline earth metal compounds are all colorless or pale yellow, with sodium chlorite (NaClO2) being the only commercially important chlorite. Heavy metal chlorites (Ag+, Hg+, Tl+, Pb2+, and also Cu2+ and ) are unstable and decompose explosively with heat or shock. ==Structure and properties==

The chlorite ion adopts a bent molecular geometry, due to the effects of the lone pairs on the chlorine atom, with an O–Cl–O bond angle of 111° and Cl–O bond lengths of 156 pm. ==Uses==

The most important chlorite is sodium chlorite (NaClO2), used in the bleaching of textiles, pulp, and paper. However, despite its strongly oxidizing nature, it is often not used directly, being instead used to generate the neutral species chlorine dioxide (ClO2), normally via a reaction with HCl: :5 NaClO2 + 4 HCl → 5 NaCl + 4 ClO2 + 2 H2O == Health risks ==

In 2009, the California Office of Environmental Health Hazard Assessment, or OEHHA, released a public health goal of maintaining amounts lower than 50 parts per billion for chlorite in drinking water. ==Other oxyanions==

Several oxyanions of chlorine exist, in which it can assume oxidation states of −1, +1, +3, +5, or +7 within the corresponding anions Cl−, ClO−, , , or , known commonly and respectively as chloride, hypochlorite, chlorite, chlorate, and perchlorate. These are part of a greater family of other chlorine oxides. ==See also==