Oxides of sulfur

  The most important oxides of sulfur are SO₂ and SO₃, but there are also a number of unstable oxides. Among these are S₂O (1.1) and S₈O (1.2), made by reactions 1.1 and 1.2; the oxides S_nO      (n = 6–10) can be prepared by reaction 1.3, exemplified for S₈O.

(1.1)                                                           (1.2)

                         (1.1)

                           (1.2)

                                        (1.3)

  Sulfur dioxide is manufactured on a large scale by burning sulfur (the most important process) or H₂S, by roasting sulfide ores (e.g. equation 1.4), or reducing CaSO₄ (equation 1.5). Desulfurization processes to limit SO₂ emissions and reduce acid rain (see Box 15.5) are now in use. In the laboratory, SO₂ may be prepared by, for example, reaction 1.6, and it is commercially available in cylinders.

                                              (1.4)

                                              (1.5)

                                            (1.6)

Selected physical properties of SO₂ are listed in Table 1.1.

Table 1.1 Selected physical properties of SO₂ and SO₃.

   At 298 K, SO2 is a liquid and a good solvent (see Section 8.5). Sulfur dioxide has a molecular structure (1.3).

(1.3)

    Sulfur dioxide reacts with O₂ (see below), F₂ and Cl₂ (equation 1.7). It also reacts with the heavier alkali metal fluorides to give metal fluorosulfites (equation 1.8), and with CsN₃ to give the Cs‏ salt of [SO₂N₃]^- (Figure 1.1).

Fig. 1.1 The structure of the azidosulfite anion, [SO₂N₃]^-, determined by X-ray diffraction at 173K for the Cs‏ salt [K.O. Christe et al. (2002) Inorg. Chem., vol. 41, p. 4275]. Colour code: N, blue; S, yellow; O, red.

                                                                (1.7)

                                                  (1.8)

   In aqueous solution, it is converted to only a small extent to sulfurous acid; aqueous solutions of H₂SO₃ contain significant amounts of dissolved SO₂ (see equations 6.18–6.20). Sulfur dioxide is a weak reducing agent in acidic solution, and a slightly stronger one in basic media (equations 1.9 and 1.10).

                           (1.9)

                  (1.10)

   Thus, aqueous solutions of SO₂ are oxidized to sulfate by many oxidizing agents (e.g. I₂, [MnO₄]^-, [Cr₂O₇]^2- and Fe³⁺‏ in acidic solutions). However, if the concentration of H‏ is very high, [SO₄]^2- can be reduced to SO₂ as in, for example, reaction 1.11; the dependence of E on [H⁺‏] was detailed in Section 7.2.

                                                                                          (1.11)

  In the presence of concentrated HCl, SO₂ will itself act as an oxidizing agent; in reaction 1.12, the Fe(III) produced is then complexed by Cl^-.

                                                                               (1.12)

  The oxidation of SO₂ by atmospheric O₂ (equation 1.13) is very slow, but is catalysed by V₂O₅.   This is the first step in the Contact process for the manufacture of sulfuric acid; operating conditions are crucial since equilibrium 1.13 shifts further towards the left-hand side as the temperature is raised, although the yield can be increased somewhat by use of high pressures of air. In practice, the industrial catalytic process operates at ≈750K and achieves conversion factors >98%.

                                                                     (1.13)

    In the manufacture of sulfuric acid, gaseous SO₃ is removed from the reaction mixture by passage through concentrated H₂SO₄, in which it dissolves to form oleum (see Section 15.9). Absorption into water to yield H₂SO₄ directly is not a viable option; SO₃ reacts vigorously and very exothermically with H₂O, forming a thick mist. On a small scale, SO₃ can be prepared by heating oleum.

                                                   (1.4)                                                                (1.5)

   Table 1.1 lists selected physical properties of SO₃. In the gas phase, it is an equilibrium mixture of monomer (planar molecules, 1.4) and trimer. Resonance structures 1.5 are consistent with three equivalent S^_O bonds, and with the S atom possessing an octet of electrons. Solid SO₃ is polymorphic, with all forms containing SO₄-tetrahedra sharing two oxygen atoms. Condensation of the vapour at low temperatures yields γ-SO₃ which contains trimers (Figure 1.2a); crystals of γ-SO₃ have an ice-like appearance. In the presence of traces of water, white crystals of β-SO₃ form; β -SO₃ consists of polymeric chains (Figure 1.2b), as does α-SO₃ in which the chains are arranged into layers in the solid state lattice. Differences in the thermodynamic properties of the different polymorphs are very small, although they do react with water at different rates. Sulfur trioxide is very reactive and representative reactions are given in scheme 1.14.

                     (1.14)

Fig. 1.2 The structures of solid state polymorphs of sulfur trioxide contains tetrahedral SO₄ units: (a)  γ-SO₃ consists of trimeric units and (b) α- and β-SO₃ contain polymeric chains.  Colour code: S, yellow; O, red.