The Oxidation of Oxygen and Related Chemistry

The following sections are included:

• Acknowledgments

• Contents

• Introduction

The discovery of was accidental. was easily made, but its correct characterization probably involved the most difficult work of my entire career. Both of the ions of this salt were then unknown in chemical compounds, and the oxidation of oxygen, demanded of PtF₆ that it be a one-electron oxidizer of unprecedented power…

Quite naturally, the formulation of O₂PtF₆ as , and the consequent requirement that the electron affinity of PtF₆ exceed 7 eV, was met with some skepticism. But there could, to my mind, only be the one possibility for the salt produced when the red gas PtF₆ precipitated the orange-red solid on mixing with oxygen gas. To fully demonstrate the remarkable one-electron oxidizing power of PtF₆, I pondered on other candidates for oxidation, which would be at least as impressive as the O₂ to example. By this time (end of 1961) our fully documented case for was already in the press (see Chap. 1, Ref. 10), and I realized that at the Argonne National Laboratory (where PtF₆ and other hexafluorides such as RuF₆ and RhF₆ had been first prepared) there was likely to be a quick follow up on our discovery. But at this point I was without any coworkers having the necessary glassworking skills, and experience in handling fluorine. My progress was therefore dictated by what I could do with my own hands…

As we have seen in Chap. 2, the investigations into the nature of XePtF₆ led to preparation of XeF₄, XeF₅PtF₆ and the investigation of XeF₂ chemistry. Indeed we soon had experience in all areas of xenon chemistry and when Professor Kai Siegbahn, who was a visitor to Berkeley at the same time as myself, in late 1967, suggested that we provide samples of the xenon fluorides for ESCA gas phase studies, I readily agreed to the collaboration. The findings, which are set out in the book^a which Siegbahn wrote with his collaborators, provided a new insight into the bonding in the set of molecules XeF₂, XeF₄, XeF₆ and XeOF₄. That early study (based on core-xenon electron binding energies) clearly showed that the F ligands bore appreciable negative charge, and that this charge was essentially the same for each fluoride and XeOF₄. In addition the charge carried by the oxygen ligand in XeOF₄ was seen to be roughly twice that carried by the F ligands. These findings fit the simple three-center four-electron bonding model for XeF₂ and the other xenon fluorides, and the notion that a Xe–O bond (as in XeOF₄) was a dative electron-pair bond Xe: → O. Thus, crudely viewed, each Xe–F bond in the neutral molecules was a single-electron bond Xe·F, and an Xe–O bond, an electron pair bond Xe:O. The more complete examination of the xenon fluorides and XeOF₄, that was made later, in collaboration with T. D. Thomas and his coworkers (see Ref. 57), extended the ESCA data to include F_1s and O_1s binding energies. This work confirmed the earlier study of Siegbahn and his coworkers in all of its essentials, the semiionic bonding of the fluorine ligand to Xe being confirmed…

Within a year of the discovery of XePtF₆ and as a result of worldwide activity, it was clear that the chemistry of the noble gases would be limited to the heavier elements as set out in my Noranda Lecture (see Ref. S2). Because of the dangerous radioactivity associated with all of the radon isotopes, this meant that the bulk of noble-gas chemistry would be that of xenon. The chemistry of krypton appeared to be limited to KrF₂ and compounds that could be derived from it. In all cases, it was clear, the range of accessible noble-gas chemistry was dictated by lower ionization potentials at the noble-gas atom, and high electronegativity and small size of the ligand atoms, as discussed in Ref. 45…

When I discovered the remarkable oxidizing power of PtF₆, I was immediately interested in the possible existence of even more potent oxidizers. Since Weinstock and his coworkers had already reported the syntheses^a of RuF₆ and RhF₆, and shown that these molecules lost fluorine more readily than PtF₆ (and RhF₆ most readily), I concluded that they could be even more potent one-electron oxidizers than PtF₆, and RhF₆ perhaps the most potent of all. I was particularly interested in finding a molecule that would remove an electron from krypton, and thought it possible that RhF₆ would do that. We therefore spent much effort in 1962, carrying out the then difficult preparation of RhF₆ (but see Chap. 8 for a new approach). Unfortunately, RhF₆ proved to be unequal to the task of oxidizing Kr, but it did oxidize Xe, with essentially the same results as for PtF₆ (see Ref. 15, Chap. 2). We then turned our attention towards the possible synthesis of AuF₆…

Early on in my graduate education (under the influence of K. H. Jack) I witnessed the value of X-ray powder diffraction in its provision of valuable structural information. This was especially valuable in the fluorine chemistry that I was then pursuing, since opportunities for single crystal production (at least in the first twenty years of my work) were minimal. I quickly set about learning to index X-ray powder data, and to solve simple structures using well-measured intensity data. Indeed, my first structure was that of PdF₂ (see Ref. 4) and was worked out with one of K. H. Jack's students, R. Maitland, acting as my mentor…

In 1983, Bougon and Lance^a announced the synthesis of what they claimed was AgF₃. (They had used the remarkable oxidizer KrF₂, in solution in aHF, in this synthesis.) Many years before, my coworker P. R. Rao, in a failed attempt to prepare a higher fluoride of gold, had obtained single crystals of AuF₃, from which the then unique hexagonal spiral structure (handed) was obtained (Ref. 27, Chap. 6). From the X-ray powder diffraction data given by Bougon and Lance it was clear that their "AgF₃" was not related structurally to AuF₃. Nor did their "AgF₃" have the reasonable alternative trifluoride structure, appropriate for a high-spin d⁸ electron configuration, adopted by^b RhF₃. In addition, the weak Curie-law paramagnetism of the "AgF₃" indicated that it was magnetically dilute and hinted that the solid might be a mixed-valence material. I became much interested in the real nature of this material. In the "Introduction" to Ref. 110, I have described the genesis of the new approach to AgF₃ synthesis…

The successful generation, by precipitation out of liquid anhydrous HF, of fluorides thermodynamically unstable with respect to loss of elemental F₂, gave a forceful reminder of the remarkable stability of that solvent towards oxidation. Soon after the preparation of AgF₃, in attempts to find evidence for cationic derivatives (e.g. [AgF₂]⁺), it was discovered (see Ref. 98) that even divalent silver, as a cation, Ag²⁺ (solvated by HF), had the capability to oxidize Xe. This made chemical sense, since a cation, having an electron deficit, should have higher electronegativity than a related oxidation-state in a neutral or anionic species. Indeed, this had three important consequences…

In evaluating the significance of the formulation (see Chap. 1) I was faced with the need to estimate the lattice energy. I was familiar with A. F. Kapustinskii's simple equation, based on thermochemical radii,^a and from the cesium chloride like arrangement of the ions in , assessed the necessary radius sum (r_c + r_a), of 6:6 coordination, required for his empirical formula. This gave me the needed lattice energy with which to estimate the minimum electron affinity of PtF₆. Even for octahedral ions like and dumbbell cations like , however, the notion of a thermochemical radius is patently false. But over the years, as our fund of structural information built up, it became clear that in most of the solids with which we were familiar, the molecules and ions were usually in close-packed array, and formula unit volumes of complex salts (at least complex fluorides) were commonly equal to the sum of the volumes of their component molecules or ions…

The discovery of Xe(SO₃F)₂, and the finding that it provides a clean safe source of the remarkable oxidizer (SO₃F)₂, led us to explore the one-electron oxidizer properties of the SO₃F radical of which it is the source. This led (see Chap. 4, Refs. 68 and 118) to the finding that (SO₃F)₂ oxidatively intercalated both graphite and boron nitride. The blue boron nitride salt was an electrical conductor (BN is insulating) and was the first authenticated example of a BN salt. This stimulated an interest in graphite intercalation, especially by fluorides…

The following sections are included:

• References