Zn(III)? Not so fast

Chemists love rogues, oddballs which seem to defy the rules and bond, react, and exist on their own terms. The rogues are valuable because they push the boundaries, teach us about new principles of structure and reactivity and challenge coveted preconceptions. One of the most striking rogues in the history of chemistry was the compound xenon hexafluoroplatinate which shattered belief in the non-reactivity of noble gases. Another immensely productive rogue was the first stable carbocation, a species that was considered too unstable to isolate until George Olah surmounted the energy barrier. Some rogues are literal rogues in the sense that they need to be incarcerated in order to prevent their unruly bonds from going haywire; witness the classic taming of cyclobutadiene by Donald Cram. There's no doubt about it; rogues' galleries are shining gems in the chemical establishment.

A couple of months ago it seemed that another minor rogue had made its appearance. Students around the world know that the most stable oxidation state of zinc is +2. The rationale is simple; unlike transition metals like iron and nickel, zinc contains a fully filled d orbital with an electronic configuration of d¹⁰s². It is therefore quite happy to lose its outermost 2s electrons and remain stable. Thus it came as a surprise when a paper detailing theoretical calculations on zinc compounds predicted the existence of a complex with Zn as Zn(III). That fact would have raised the eyebrows of a million freshmen memorizing transition metals trends for their final exams. The existence of the unusual zinc was based on quantum chemical calculations done by Puru Jena's group at VCU and their basic rationale was that a ligand that was oxidizing and electronegative enough would essentially force the metal to donate extra electrons. They seemed to find such a complex in Zn(AuF₆)₃. On the face of it that would make sense, but it's worth keeping in mind that nature loves filled orbitals; for instance, even a ligand as oxidizing and electronegative as CF3 does not force copper to adopt the +2 state (copper's most stable oxidation state being +1, resulting from its d¹⁰s¹ configuration).

Now another paper suggests, with ample evidence it seems, that the suspected Zn(III) rogue might be an upstanding chemical citizen after all. Sebastian Riedel's group in Berlin has examined the purported Zn(III) complex and found that not only does zinc adopt its good old +2 oxidation state in the compound, but that the compound would not be thermodynamically stable. They use a quantum chemical method more sophisticated than the previous one and include more relativistic effects. Where does relativity enter the picture, you ask. It turns out that for heavy metals, the speed (in a crude sense) of the d s-electrons can be fast enough to warrant relativistic calculations; indeed, relativistic quantum chemistry has shed light on many commonplace and yet unusual phenomena, like the color of gold and the liquid state of mercury at room temperature. Relativity is not the exclusive domain of physicists.

I am not enough of an expert in quantum chemistry, but it clearly seems from the latest paper that the more accurate calculations which account for finer details of electron correlation and relativistic effects indicate something striking; the purported complex with Zn in the +3 oxidation state would undergo an exothermic reaction. Put simply, it would not be stable and would decompose to a compound with good old Zn(II). For me, a more troubling fact was the structure of the compound published in the original paper. The structure was supposed to be a low-energy minimum but it has two fluorines practically colliding with each other in an unholy, 197 pm, sub-Van der Waals radius squeeze (illustrated above). The present calculations find another low-energy structure existing only 10 kJ/mol above the previous structure that is consistent with a Zn(II) state. Other calculations indicate that the fluorines in the previously proposed structure are best described as radical anions bridging two AuF₅ units.

There's other theoretical evidence in there too, along with citations of experimental facts that point to the recalcitrance of Zn to adopt a +3 state. The study seems to be as solid a piece of evidence against Zn(III) as can be obtained using current levels of theory. But the real reason I want to point out this paper is because it illustrates one of the key roles of theory and computation in chemistry; as an invitation to experiment. Just because the complex is thermodynamically unstable does not mean it cannot be isolated under any conditions. One of the great lessons of chemical science in the last forty years has been the fact that given the right conditions almost anything can be synthesized, stored and isolated (Cram's cyclobutadiene again being a case in point) and that the words stable and unstable are highly relative constructions created by an impoverished chemical lexicon.

This tussle between two predictions illustrates the function of theory in chemistry as the ultimate teaser; if I were an experimentalist I would be rubbing my hands in glee, sending my armies of postdocs and graduate students into the lab to try to synthesize Zn(AuF₆)₃. As Excimer commented on another site, "Enough theory! Someone make the damn thing already".
Image source: ACS publications.

Three lessons

When you have been in graduate school for two or three years, it's a good question to ask which are the important lessons you have learnt from your time there. These are not the "lessons of life" I am talking about- they are perhaps even more important- but technical lessons which spring to your mind. In my case, three lessons strikingly come to my mind. All of them are related to work that our group has done, but also to other people's work and general thought. All of them deal with material published in the literature that is unfortunately "fiction", and articles in journals don't seem to acknowledge this fact. Here they are:

1. Oxidation states for transition metals greater than 1 (eg. Au3+, Cu2+) are fictional and non-existent (related post and references)
2. Single "average" structures derived from NMR for flexible molecules are fictional and "virtual" (related post and references)
3. 3-fluoropiperidines exist in solution almost exclusively as the axial-F conformer. (related post and references)

All transition metals exist only in +1 state (or a close approximation thereof)

Yes, I know; reading this statement might make people think that megagallons of ethanol are coursing through my arteries. It sounds quite counterintuitive. But I have become convinced that this is the case. I have not explored the matter in exhaustive detail, and some questions still linger. The explanation can get long-winded and I will reserve it for sometime when my own thoughts are clearer. But here's the summary.

There is a very simple reason why all transition metals will prefer only the +1 state in their complexes; Pauling's electroneutrality principle.
More simply put, consider this; won't it be improbable for a charge of +2 or +3 to accumulate on a metal ion? There will be repulsive destabilization. We know that charge tries to disperse itself and stabilize any system. So this positive charge will try to redistribute itself and get delocalized, thereby making the system close to electrostatically neutral. Where can it do this? Well, we don't want to forget 'the other half'- the ligands. What this means is that it's actually the ligands that are getting oxidised by helping to delocalize the positive charge. So even in a compound like AuCl3, the Au has a +1 charge, and it's really the chlorides that are getting oxidised.

The whole root of the problem is that chemists' definition of oxidation state actually is quite arbitrary, based on qualitative notions of electronegativity. We think that since chlorine is more electronegative than gold, it will carry all the negative charge. But Pauling's electroneutrality principle tells us otherwise.

However, this is not something we usually think about. Consider the simple ammonium cation. Organic chemists depict it with the positive charge on nitrogen. But it's a well-known fact that the positive charge in fact is distributed over the hydrogens. Any system will try to stabilize itself by dispersing charge as much as it can. That's what the ammonium cation does, nicely dispersing the positive charge over the hydrogens.

We might say; but how can halogens get oxidised? They are supposed to be among the most strongly oxidising elements. But note, that this statement applies to diatomic halogen molecules, not to halogen ions in solution.

You can take my word that even the fiercely electronegative CF3- anion in the complex [Cu(CF3)4]- is oxidised, and that the copper exists as Cu(I) in this state. Even a look at the formula suggests otherwise; the Cu must be existing as Cu(III). But you don't have to take my word for it. Consider this publication in Angewandte Chemie. There was also a small debate, published in another issue of the same journal, which further combs some fine points and definitions and clarifies things. High level calculations strongly suggest that the Cu exists as Cu(I).

High level calculations, but more logically, the electroneutrality principle and basic concepts of electrostatics.

The point is, many things in chemistry are defined either arbitrarily, or with respect to some standard. Based on the existence of another chemical species, many things can change for a given species. Given the right conditions, almost anything can get oxidised or reduced. Given the right conditions, anything can be made stable. After all, nobody had thought that stable carbocations could exist. But George Olah opened up an entire new field of carbocation chemistry with his pioneering work. In chemistry, there are no absolutes. Which again proves why it is so interesting.

Actually, it's simple to think about. Usually we assume that the metal is changing its oxidation state. Instead of that, let's adopt the other perspective; the metal makes the ligand change its oxidation state. The result is the same. We still see AuCl3 in solution. But it's really the Cl- that has changed, and not the Au.

Some questions still linger; then why do all these metals seem to exist in different oxidation states. More importantly, how can one prove this contention by experiment? After all, ours is an experimental science.

But even in the absence of definite experiments, if one ad hoc concept is taken for granted, then another ad hoc concept surely cannot be worse. Especially if it sounds more logical and convincing. Like all metals existing only in +1 oxidation state (or something close to that).

Important addendum: I should mention that when I say +1, the actual number is better defined as "close to +1". Calculations show that it can change from +0.7 to about +1.2 (this bears on the question d_orbital asked in the comments) depending on the ligand environment. Thus, in hemoglobin for example, in the oxidised and reduced form, Fe won't exist as exactly +1 both of the times. Its oxidation state will fluctuate between decimal places close to one.
And, as Mitch quite rightly says, in most of our daily life, metals have an unequivocal oxidation state of 0!!