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The Story of O

Roald Hoffmann

The Ring

So now we have oxygen chemistry in the atmosphere, in the cell, in a decontamination process. But have we exhausted all of oxygen's secrets?

Figure 3. In <em>a</em> . . .Click to Enlarge Image

For a few intriguing small molecules you can draw a perfectly good Lewis structure. But they don't exist. Beginning chemistry students, nicely unsullied by the burden of knowing the answer, often give us one—ask them for a Lewis structure of ozone, O3, and they are as likely to write down the cyclic molecule in Figure 3a as the known form in Figure 3b, for which one has to draw two resonance structures. So what's wrong with the ring in Figure 3a? Or four-, or five-, or six-membered cyclic variants? The usual explanation is that the rings have too many unshared electron pairs close to one another—note the 12 dots, symbolizing electrons, in Figure 3a.

Figure 4. Hypothetical tungsten complex . . .Click to Enlarge Image

Now the situation grows interesting. A theoretical analysis shows that the interconversion of the normal bent form (the O-O-O angle in ozone is 117 degrees) and the cyclic form of O3 is what R. B. Woodward and I called a forbidden reaction. Which means that just the breaking of what one might have thought to be a single weak bond should have a substantial energy barrier. So there may yet be hope for stable cyclic O3.

Indeed, the best calculations today confirm the metastability of this ring. Cyclic ozone lies about 130 kilojoules per mole above normal O3 but has a barrier of no less than 95 kilojoules per mole preventing conversion to the open form. There is an even bigger barrier to falling apart to O2 + O.

Sulfur is like oxygen, in some of its chemistry. So what happens for the sulfur analogues of ozone? SO2 forms a three-sided ring that is at much higher energy (computed at about 400 kilojoules per mole) than the open structure (the geometries resemble those of ozone in Figures 3a and 3b). Nevertheless, cyclic SO2 also has a relatively large barrier to breaking a bond and opening up an angle—84 kilojoules per mole. S or O, there is no escape from the constraints quantum mechanics puts on reaction barriers. In S3 the ring is calculated to lie only 33 kilojoules per mole above the open form, again with a substantial barrier to opening.

Cyclic O3 or S3 has not been detected. Should one give up on them? By no means. One of the beauties of organometallic chemistry is that an appropriately chosen MLn fragment (M=transition metal, L=ligand) can bind to and stabilize a molecule that, by itself, is unstable. So . . . to find the elusive cyclic ozone, one could try to make a molecule such as shown in Figure 4. Beate Flemmig, Shen-shu Sung and I are trying to predict what ligands on a metal (tungsten looks best) might be able to do this.

Actually, there has been one observation of cyclic ozone on an annealed magnesium-oxide surface. The technique used to image the molecules, transmission electron diffraction, showed O3 rings that appeared to be centered over underlying Mg ions.

Incidentally, the solid state is a fine place to look for oxygen in other weird forms. A mystery remains (one of many) about the chemical nature of oxygen in the cuprate superconductors. Oxygen doping is necessary for high-temperature superconductivity (high Tc) in systems such as La2CuO4. During the synthesis of this composite, the oxygen enters as O2, but it surely does not reach its equilibrium positions in the lattice in this form. To get at the problem in another way—all the signs are that the holes (missing electrons) in high Tc cuprates are partially on Cu2+, partially on (formal) O2-. But a hole, or one electron less, on O2- makes it O-. And this radical (here they are again!) will not sit still in the lattice. Large or small motions of the oxygens, static or dynamic formation of O21-, O22-, O23- or even a bigger aggregate—what will it be?

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