Sodium chlorite is a chemical compound used in the manufacture of
The free acid, chlorous acid, HClO2, is only stable at low
concentrations. Since it cannot be concentrated, it is not a
commercial product. However, the corresponding sodium salt, sodium
chlorite, NaClO2 is stable and inexpensive enough to be
commercially available. The corresponding salts of heavy metals
(Ag+, Hg+, Tl+, Pb2+, and also Cu2+ and NH4+) decompose explosively
with heat or shock.
Sodium chlorite is derived indirectly from sodium chlorate, NaClO3.
First, the explosively unstable gas chlorine dioxide, ClO2 is
produced by reducing sodium chlorate in a strong acid solution with
a suitable reducing agent (for example, sodium chloride, sulfur
dioxide, or hydrochloric acid). The chlorine dioxide is then
absorbed into an alkaline solution and reduced with hydrogen
peroxide, H2O2 yielding sodium chlorite.
The main application of sodium chlorite is the generation of
chlorine dioxide for bleaching and stripping of textiles, pulp, and
paper. It is also used for disinfection in a few municipal water
treatment plants after conversion to chlorine dioxide. An advantage
in this application, as compared to the more commonly used
chlorine, is that trihalomethanes are not produced from organic
contaminants. Sodium chlorite, NaClO2 also finds application as a
component in therapeutic rinses, mouthwashes, toothpastes and gels,
mouth sprays chewing gums and lozenges, and also in contact lens
cleaning solution under the trade name purite.
In organic synthesis, sodium chlorite is frequently used for the
oxidation of aldehydes to carboxylic acids. The reaction is usually
performed in the presence of a chlorine scavenger.
Sodium chlorite, like many oxidizing agents, should be protected
from inadvertent contamination by organic materials to avoid the
formation of an explosive mixture.