Acid Equilibria Sulfite ion is a weak base, but does undergo some hydrolysis to produce basic solutions. In acidic solution, the equilibria are shifted to form sulfurous acid, resulting in the evolution of SO2 gas. Sulfur dioxide is a colorless gas with a characteristic choking odor. \[\ce{SO3^{2-}(aq) + H2O(l) <=> HSO3^{-}(aq) + OH^{-}(aq)} \nonumber \] wiht \(K_b = 1.8 \times 10^{-7}\) \[\ce{HSO3^{-}(aq) + H2O(l) <=> H2SO3(aq) + OH^{-}(aq) } \nonumber \] with \(K_b = 1 \times 10^{-12}\) \[\ce{H2SO3(aq) <=> H2O(l) + SO2(g)} \nonumber \] Solubility The sulfites of \(\ce{Na^{+}}\), \(\ce{K^{+}}\), and \(\ce{NH4^{+}}\) are soluble in water. Most other sulfites are insoluble in water. However, due to the basic nature of \(\ce{SO3^{2-}}\), all sulfites dissolve in acidic solution. Oxidation-Reduction Sulfite ion is readily oxidized to sulfate. On prolonged exposure to air, this oxidation occurs with atmospheric oxygen: \[\ce{2SO3^{2-}(aq) + O2(g) -> 2SO4^{2-}(aq)} \nonumber \] Sulfite or sulfur dioxide will decolorize permanganate. This decolorization serves as a convenient test for sulfur dioxide: \[\ce{2MnO4^{-}(aq) + 5SO2(g) + 2H2O(l) -> 5SO4^{2-}(aq) + 2Mn^{2+}(aq) + 4H^{+}(aq)} \nonumber \]