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Sulfur Oxyacids and Salts

Sulfur Oxyacids and Salts

The preparation of sulfuric acid, H2SO4 (shown in the figure below), begins with the oxidation of sulfur to sulfur trioxide and then converting the trioxide to sulfuric acid. Pure sulfuric acid is a colorless, oily liquid that freezes at 10.5 °C. It fumes when heated because the acid decomposes to water and sulfur trioxide.

The heating process causes the loss of more sulfur trioxide than water, until reaching a concentration of 98.33% acid. Acid of this concentration boils at 338 °C without further change in concentration (a constant boiling solution) and is commercially concentrated H2SO4. The amount of sulfuric acid used in industry exceeds that of any other manufactured compound.

Sulfuric acid has a tetrahedral molecular structure.

The strong affinity of concentrated sulfuric acid for water makes it a good dehydrating agent. It is possible to dry gases and immiscible liquids that do not react with the acid by passing them through the acid.

Sulfuric acid is a strong diprotic acid that ionizes in two stages. In aqueous solution, the first stage is essentially complete. The secondary ionization is not nearly so complete, and $${\text{HSO}}_{4}{}^{-}$$ is a moderately strong acid (about 25% ionized in solution of a $${\text{HSO}}_{4}{}^{\text{−}}$$ salt: Ka = 1.2 $$×$$ 10−2).

Being a diprotic acid, sulfuric acid forms both sulfates, such as Na2SO4, and hydrogen sulfates, such as NaHSO4. Most sulfates are soluble in water; however, the sulfates of barium, strontium, calcium, and lead are only slightly soluble in water.

Among the important sulfates are Na2SO4⋅10H2O and Epsom salts, MgSO4⋅7H2O. Because the $${\text{HSO}}_{4}{}^{-}$$ ion is an acid, hydrogen sulfates, such as NaHSO4, exhibit acidic behavior, and this compound is the primary ingredient in some household cleansers.

Hot, concentrated sulfuric acid is an oxidizing agent. Depending on its concentration, the temperature, and the strength of the reducing agent, sulfuric acid oxidizes many compounds and, in the process, undergoes reduction to SO2, $${\text{HSO}}_{3}{}^{\text{−}},$$$${\text{SO}}_{3}{}^{2-},$$ S, H2S, or S2−.

Sulfur dioxide dissolves in water to form a solution of sulfurous acid, as expected for the oxide of a nonmetal. Sulfurous acid is unstable, and it is not possible to isolate anhydrous H2SO3. Heating a solution of sulfurous acid expels the sulfur dioxide. Like other diprotic acids, sulfurous acid ionizes in two steps: The hydrogen sulfite ion, $${\text{HSO}}_{3}{}^{\text{−}},$$ and the sulfite ion, $${\text{SO}}_{3}{}^{2-},$$ form. Sulfurous acid is a moderately strong acid. Ionization is about 25% in the first stage, but it is much less in the second (Ka1 = 1.2 $$×$$ 10−2 and Ka2 = 6.2 $$×$$ 10−8).

In order to prepare solid sulfite and hydrogen sulfite salts, it is necessary to add a stoichiometric amount of a base to a sulfurous acid solution and then evaporate the water. These salts also form from the reaction of SO2 with oxides and hydroxides. Heating solid sodium hydrogen sulfite forms sodium sulfite, sulfur dioxide, and water:

$$2{\text{NaHSO}}_{3}(s)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;{\text{Na}}_{2}{\text{SO}}_{3}(s)+{\text{SO}}_{2}(g)+{\text{H}}_{2}\text{O}(l)$$

Strong oxidizing agents can oxidize sulfurous acid. Oxygen in the air oxidizes it slowly to the more stable sulfuric acid:

$$2{\text{H}}_{2}{\text{SO}}_{3}(aq)+{\text{O}}_{2}(g)+2{\text{H}}_{2}\text{O}(l)\;\stackrel{\phantom{\rule{0.4em}{0ex}}\text{Δ}\phantom{\rule{0.4em}{0ex}}}{\to }\;2{\text{H}}_{3}{\text{O}}^{\text{+}}(aq)+2{\text{HSO}}_{4}{}^{\text{−}}(aq)$$

Solutions of sulfites are also very susceptible to air oxidation to produce sulfates. Thus, solutions of sulfites always contain sulfates after exposure to air.