If you ever wanted to name a single technology that transformed society while remaining nearly invisible, catalysts would be an excellent answer. Transforming one stable compound (say, nitrogen gas) into another stable compound (nitrogen-based fertilizer, for instance) requires an enormous amount of energy unless a catalyst is used. Then it suddenly becomes worthwhile. Indeed, that one reaction series—nitrogen gas to ammonia to fertilizer—is responsible for feeding the majority of people on the planet today.

The difficulty is that we don't really have a great understanding of how most catalysts work. And a recent paper adds considerably to that confusion by showing that single atoms can act as very efficient catalysts.

Why are catalysts necessary?

When we make new compounds, whether they're plastics, gasoline, fertilizers, or pretty much any modern material, we require a catalyst at some point. The reason for this is simple. We take raw materials from the world around us. These materials are stable. If they weren't, they would have reacted long ago to form a different material that was stable. Stable materials like to stay just as they are.

The materials that we want to form are also stable. (Yes, even most explosives are relatively stable—the last thing you want is an explosive that will go off in your face because you dropped it.) The goal of most industrial chemistry is to transform a set of stable materials into a different set of stable materials. In order to do that, you need to add a lot of energy, much more than we are able to produce. If that was the end of the story, the latter part of the industrial revolution would have died in its infancy because its products would have been far too expensive.

The trick of a catalyst is to reduce the amount of energy required at critical steps in the process. This is what makes most modern materials affordable.

How do catalysts work?

No one is too sure how most catalysts work (there are some proteins that are pretty well understood). One thing is sure: all reactions take place at the surface of a catalyst. Increasing the available surface area of a catalytic material is one sure way to increase the efficiency of a catalyst, but we also know that not every surface atom in a catalyst material behaves as a catalyst. Imagine that you have a crystal of palladium atoms, roughly cut into a sphere. To make a sphere, you have a lot of edges and corners to crystalline planes. The atoms at the edges and corners are the ones that do the catalysis.

The hypothesis is that when something like a nitrogen molecule attaches to the corners and edges of a catalyst, it gets a bit distorted, making it a little easier to break the bonds between nitrogen atoms. In addition, the catalyst should be able to donate or accept electrons to aid the bond breaking process, so not just any material with crystalline edges can do the job. But frankly, this all a little difficult to observe and confirm.

If this general picture were the whole story, the following experiment should not have produced the result that it did. Researchers from China and the US chose to study the production of hydrogen and carbon dioxide from water and carbon monoxide, a starting point for fuel production. One catalyst for this process is iridium, which is very expensive. In the ideal case, you want every iridium atom to play the role of the catalyst and use the smallest amount possible.

To achieve this, the researchers doped iron oxide with iridium at such low concentrations that most of the iridium atoms were alone in the iron oxide matrix. They used a bunch of very high-resolution imaging techniques to determine that at the lowest dopant concentrations (0.01 percent), the iridium atoms were almost always separated from each other by several iron oxide molecules. At higher concentrations, the iridium atoms started to clump together. These different samples provide a varying ratio between single iridium atoms and groups of iridium atoms.

By comparing the reaction rates between these different samples, the researchers were able to figure out how effective the single iridium atoms were. It turns out that even at the highest iridium concentrations (2.4 percent), single iridium atoms were responsible for 75 percent of the observed reaction products.

A catalytic reaction is a very complicated beast because it must proceed in steps that often involve molecules moving along the surface from one site to another. That movement is usually very slow, so the per-active-catalyst-site reaction rates are on the order of a single reaction per second. In this case, though, the single iridium atoms manage just over two reactions per second.

As far as I'm concerned, a more important fact is that at these concentrations, very few of the iridium atoms end up at corners or edges, so these structural features can only play a minor role. Indeed, for single atoms, it's very difficult to imagine how structure could play a role. Of course, flat surfaces without any edges will also catalyze reactions, but they are very slow and very ineffective. In this work, though, iridium atoms that were essentially part of a flat surface were both very effective and very fast, which means that I now have even less of a clue how this reaction proceeds.

Well, that's not entirely true. The iridium is not the same as the iron it replaces, so the crystal structure is a bit distorted at those points. Structure could still play a role, but these sorts of defects usually have to be at edges or corners to be as effective.

In any case, my confusion is a good thing. By highlighting the shortcomings of a classic explanation for catalytic activity, new ideas will be generated.

Journal of the American Chemical Society, 2013, DOI: 10.1021/ja408574m