Let’s do some pure chemistry today, because an interesting paper has come out about a reaction that every student learns about in their sophomore organic chemistry course: the Birch reduction. It’s a powerful technique that will do some things that very few other reactions will do for you (such as break up the aromaticity of benzene derivatives). That strength comes from the nature of the reagent itself. It’s hard to get more hard core about adding electrons to a compound than soaking it in a bath of solvated electrons themselves.

I’ve written about that one a few times over the years, mainly in terms of how to avoid running it under the classic conditions. Those are a mild pain in the rear to run, because you need to round up a tank of ammonia and a dry-ice condenser to rig on top of your reaction flask. You’ll be opening up that tank gradually to run a stream of ammonia gas across the cold finger, whereupon it will start dripping clear liquid ammonia (boiling point -33 degrees C) into a receiving flask. Unfortunately, if you want the best yields and reproducibility, you’re going to want to distill that first run of ammonia into your actual reaction flask (details here in a very good guide to running these reactions). There’s often some water in the system, and there can be traces of iron from the inside of the cylinder, which can mess up a Birch at catalytic levels. I have not always taken the trouble to do this, but then, my Birch reductions have not always gone as well as I would have wanted them to, either, so there is that.

Ammonia’s a pretty good solvent, a lot of organic compounds dissolve in it just fine. Birch reactions are generally run with an alcohol co-solvent (you need a proton source), and that helps even more. Once you have your starting material dissolved in this mixture, down in your dry ice bath, you start the enjoyable part: flinging in small bits of metal. The reaction can be run with lithium, sodium, potassium, and even calcium – actually figuring out which of those will work best is not so easy to do from first principles, which can at times be another complication. Most people use lithium or sodium, because they’re a lot easier to handle than potassium, from a not-producing-beautiful-magenta-flames standpoint, and straight calcium metal is relatively rare on the lab shelves.

As those bits of reactive metal go in, they dissolve in swirls of deep, vivid blue that trace through the clear ammonia solution and disappear. Here’s a video at YouTube showing just that. The blue is the color of the solvated electrons themselves, and it disappears as they react with your starting material. The classic way to tell when your reaction is done is persistence of the blue color (nothing left to react with!) But there’s an interesting phenomenon that you can see if you go wild and keep flinging metal into such a solution: the blue gets deeper and darker, but then you start making something new: a bronze-colored solution, which can be a separate layer (with the lithium reactions, I’ve seen it floating around on top of the blue phase). You can turn the whole thing to bronze phase with enough metal. But what is that stuff, and why does it look and act like that?

Well, it looks like a metal – a liquid metal. And it turns out to be an excellent conductor, too, far better than the blue solution. Now there’s a new paper that went to the trouble of doing photoelectron spectroscopy on the stuff to see what’s really going on. (Here’s some interesting background on the PhD student involved, Ryan McMullen). This is not such an easy experiment to do. The apparatus used cold microjets of the ammonia solutions spraying across a synchrotron-derived X-ray beam, which is a setup that I will assume took just a tiny bit of troubleshooting along the way. Basically, you’re using that vicious X-ray beam to blast electrons out of the surface of a sample when they absorb all that excess energy. The energy that they have as they escape depends on a lot of important characteristics – their original electronic state/energy level, for sure, as well as their rotational and vibrational state. The peaks you get are quite diagnostic and the technique is also extremely sensitive (and earned its developer, Kai Siegbahn, a Nobel Prize in 1981.

What you see when you blast the blue stuff is a peak at the “vertical displacement energy” (named for how these things are straight-up lines on an energy diagram) of about 2 electron volts. That is from the solvated electrons, which at anything short of extreme dilution are spin-paired dielectrons hanging out together. That peak is proportional to the concentration of alkali metal, and is pretty much the same no matter which metal you use, both of which point to them just being the electrons without any participation from the metals themselves.

And as you keep increasing the amount of alkali metal, that peak changes. It shades into something asymmetric, with a sharp cutoff on the low-energy end. That is a Fermi edge, and it’s just what you see with a bulk metal. I’m sure if you showed that photoelectron spectrum to someone in the field without telling them what it was, that’s exactly what they would guess as soon as they saw the trace. You also get a couple of “plasmon peaks”, which are in the visual range and account for the shiny bronze color – the same thing that accounts for the appearance of bulk metal samples as well. You can think of metals as being a “free electron gas” floating around through the metal atom lattice, and that’s what is being seen in these bronze-phase sample as well.

The gradual transition of the peaks means that the two situations (dielectrons surrounded by ammonia molecules, versus conducting delocalized electron gas) must exist at the same time and that the transition between them is smooth and not a sudden jump. No one knows yet just how that’s done – if there are tiny domains where each one of these obtains, separate from the other, or what. Those electron pairs must sort of coalesce as the concentration goes up, and you can imagine the sort of microdroplet mixture process that’s seen in some liquid-liquid phase transitions. So well before you start seeing bronze stuff in your Birch solution, you’ve already got a different situation developing down in the flask.

This work may seem pretty esoteric, and it is, in a way. But I think it’s representative of a lot of 21st-century science, in that we’re getting down to gritty microdetails about what’s really happening in physical samples and systems. You can watch crystals form in a concentrated solution, for example, but just how exactly do those molecules come together and line up so well? Do individual ones just ping around until they hit a spot they like, or do they come together as groups (how small, how large?) and assemble as larger building blocks? You can tell, similarly, that calcium ions move back and forth across living cell membranes – but just how exactly do they do that? What proteins do it, and what are their features that allow this to happen – is there a kind of “river” of ions flowing down the middle of such a protein channel, or do individual ones sort of Tarzan along from one favorable protein interaction to another? We’re really getting down to the nanoscale details, where chemistry, physics, and biology can all start to run together. Things are different down there, and we need to know what those differences are, why they exist, and how we can maneuver them to do useful things for us.