Oxygen fluorides

    Oxygen difluoride, OF₂ (1.1), is highly toxic and may be prepared by reaction 1.1. Selected properties are given in Table 1.1. Although OF₂ is formally the anhydride of hypofluorous acid, HOF, only reaction 1.2 occurs with water and this is very slow at 298 K. With concentrated alkali, decomposition is much faster, and with steam, it is explosive.

(1.1)

             (1.1)

                                (1.2)

   Pure OF₂ can be heated to 470Kwithout decomposition, but it reacts with many elements (to form fluorides and oxides) at, or slightly above, room temperature. When subjected to UV radiation in an argon matrix at 4 K, the OF^- radical is formed (equation 1.3) and on warming, the radicals combine to give dioxygen difluoride, O₂F₂.

                           (1.3)

    Dioxygen difluoride may also be made by the action of a high-voltage discharge on a mixture of O₂ and F₂ at 77–90K and 1–3 kPa pressure. Selected properties of O₂F₂ arelisted in Table 1.1.

Table 1.1 Selected physical properties of oxygen and sulfur fluorides.

   The low-temperature decomposition of O₂F₂ initially yields O₂F^- radicals. Even at low temperatures, O₂F₂ is an extremely powerful fluorinating agent, e.g. it inflames with S at 93 K, and reacts with BF₃ (equation 15.8) and SbF₅ (reaction 1.4).

                   (1.4)

      The molecular shape of O₂F₂ (1.2) resembles that of H₂O₂ (Figure 15.9) although the internal dihedral angle is smaller (878). The very long O^_F bond probably accounts for the ease of dissociation into O₂F^* and F^*. Structures 1.3 show valence bond representations which reflect the long O^_F and short O^_O bonds; compare the O^_O bond distance with those for O₂ and derived ions (Section 15.4) and H₂O₂ (Table 15.3).

(1.2)                                         (1.3)