Sulfuric acid, H₂SO₄

   Sulfuric acid is by far the most important of the oxoacids of sulfur and is manufactured on a huge scale by the Contact process. The first stages of this process (conversion of SO₂ to SO₃ and formation of oleum) were described in Section 15.8; the oleum is finally diluted with water to give H₂SO₄. Pure H₂SO₄ is a colourless liquid with a high viscosity caused by extensive intermolecular hydrogen bonding.  Gas-phase H₂SO₄ molecules have C₂ symmetry  with S^_O bond distances that reflect two different types of S^_O bond. Diagram 1.1 shows a hypervalent structure for H₂SO₄, and 1.2 gives a bonding scheme in which the S atom obeys the octet rule (refer back to the discussion of bonding in Section 15.3). In the sulfate ion, all four S^_O bond distances are equal (149 pm) because of charge delocalization, and in [HSO₄]^- , the S^_OH bond distance is 156 pm and the remaining S^_O bonds are of equal length (147pm).

                                                 (1.1)                     (1.2)                     (1.3)

    In aqueous solution,H₂SO₄ acts as a strong acid (equation 1.1) but the [HSO₄]^- ion is a fairly weak acid (equation 1.2 and Table 15.8). Two series of salts are formed and can be isolated, e.g. KHSO₄ and K₂SO₄.

                                                           (1.1)

                                                             (1.2)

   Dilute aqueous H₂SO₄ (typically 2M) neutralizes bases (e.g. equation 1.3), and reacts with electropositive metals, liberating H₂, and metal carbonates (equation 1.4).

                                     (1.3)

              (1.4)

   Commercial applications of sulfate salts are numerous, e.g. (NH₄(₂SO₄ as a fertilizer, CuSO₄ in fungicides, MgSO₄ as a laxative, and hydrated CaSO4  uses of H₂SO₄ were included in Figure 15.3. Concentrated H₂SO₄ is a good oxidizing agent (e.g. reaction 15.88) and a powerful dehydrating agent (see Box 11.4); its reaction with HNO₃ is important for organic nitrations (equation 1.5).

                                   (1.4)

Commercial applications of sulfate salts are numerous, e.g.)NH₄(₂SO₄  as a fertilizer, CuSO₄ in fungicides, MgSO₄ as a laxative, and hydrated CaSO₄  uses of H₂SO₄ were included in Figure 15.3.  Concentrated H₂SO₄ is a good oxidizing agent (e.g. reaction 15.88) and a powerful dehydrating agent; its reaction with HNO₃ is important for organic nitrations (equation 1.5).

                                                     (1.5)

   Although HF/SbF₅ is a superacid, attempts to use it to protonate pure H₂SO₄ are affected by the fact that pure sulfuric acid undergoes reaction 1.6 to a small extent. The presence of the [H₃O] ⁺‏‏ ions in the HF/SbF₅ system prevents complete conversion of H₂SO₄ to [H₃SO₄]⁺‏.

                                                                                        (1.6)

An ingenious method of preparing a salt of [H₃SO₄]‏ is to use reaction 1.7 which is thermodynamically driven by the high Si^_F bond enthalpy term in Me₃SiF. In the solid state structure of [D₃SO₄]‏[SbF₆]^- (made by using DF in place of HF), the cation has structure 1.3 and there are extensive O^_D……F interactions between cations and anions.

                       (1.7)