Ozone

    Ozone, O₃, is usually prepared in up to 10% concentration by the action of a silent electrical discharge between two concentric metallized tubes in an apparatus called an ozonizer. Electrical discharges in thunderstorms convert O₂ into ozone. The action of UV radiation on O₂, or heating O₂ above 2750K followed by rapid quenching, also produces O₃. In each of these processes, O atoms are produced and combine with O₂ molecules. Pure ozone can be separated from reaction mixtures by fractional liquefaction; the liquid is blue and boils at 163K to give a perceptibly blue gas with a characteristic ‘electric’ smell. Molecules of O₃ are bent (Figure 1.1). Ozone absorbs strongly in the UV region, and its presence in the upper atmosphere of the Earth is essential in protecting the planet’s surface from over-exposure to UV radiation from the Sun. Ozone is highly endothermic (equation 1.1). The pure liquid is dangerously explosive, and the gas is a very powerful oxidizing agent (equation 1.2).

                  (1.1)   

                     (1.2)

Fig. 1.1 The structures of O₃ and [O₃]^-, and contributing resonance structures in O₃. The O^_O bond order in O₃ is taken to be 1.5.

       The value of E^o in equation 1.2 refers to pH = 0, and at higher pH, E diminishes: ‏1.65V at pH = 7, and 1.24V at pH = 14. The presence of high concentrations of alkali stabilizes O₃ both thermodynamically and kinetically.   Ozone is much more reactive than O₂ (hence the use of O₃ in water purification). Reactions 1.3–1.5 typify this high reactivity, as does its reaction with alkenes to give ozonides.

                                           (1.3)

                        (1.4)

                                        (1.5)

Potassium ozonide, KO₃ (formed in reaction 1.6), is an unstable red salt which contains the paramagnetic [O₃]^- ion (Figure 1.1). Ozonide salts are known for all the alkali metals. The compounds [Me₄N][O₃] and [Et₄N][O₃] have been prepared using reactions of the type shown in equation 1.7. Ozonides are explosive, but [Me₄N][O₃] is relatively stable, decomposing above 348K.

                        (1.6)

        (1.7)

   Phosphite ozonides, (RO)₃PO₃, have been known since the early 1960s, and are made in situ as precursors to singlet oxygen. The ozonides are stable only at low temperatures, and it is only with the use of modern low-temperature crystallographic methods that structural data are now available. Figure 1.1 shows the structure of the phosphite ozonide

Fig. 1.1 The structure (X-ray diffraction at 188 K) of the phosphite ozonide EtC(CH₂O)₃PO₃ [A. Dimitrov et al. (2001) Eur. J. Inorg. Chem., p. 1929]. Colour code: P, brown; O, red; C, grey; H, white.

Fig. 1.2 Schematic representations of the structures of some allotropes of sulfur: (a) S₆, (b) S₇, (c) S₈ and (d) catena-S1 (the chain continues at each end).

prepared by the steps in scheme 1.8. In the PO₃ ring, the P^_O and O^_O bond lengths are 167 and 146 pm, respectively; the ring is close to planar, with a dihedral angle of 78.

      (1.8)