In the middle of the periodic table of elements, on the block that bridges the two jutting sides, is a series of elements known as transition metals. The electronic composition of transition metals makes them great catalysts for some of earth’s most life-enabling reactions. They mediate key reactions, for example, in photosynthesis and help convert nitrogen in the atmosphere so it can be used as a nutrient to sustain life.
“You can have these huge, greasy protein scaffolds made up of thousands of atoms of which at most maybe 10 are a transition metal,” says Kyle M. Lancaster, Chemistry and Chemical Biology, “and somehow those very few atoms are able to drive the most important reactions on the planet. What we do is try to understand what it is about an environment surrounding a transition metal—copper, iron, nickel, etcetera—that allows them to mediate very difficult reactions.”
To achieve this goal, the Lancaster lab uses every tool available—from x-ray to laser-based and ultraviolet spectroscopies, to kinetic diagnostics and biochemical techniques, to protein synthesis, and more. These tools throw all the hammers at the nail, and he describes his lab as a one-stop shop for helping labs around the world understand their compounds. “We elevate the chemical story,” he says.
Lancaster and his team are probing the role of transition metals in essential biological reactions and zooming in to understand their reactivity. In biological and nonbiological areas, their discoveries are overhauling long-held notions of how these fundamental elements work.
Converting Nitrogen to Nutrients—A Quest to Understand the Cycle
One area where Lancaster’s lab is making breakthroughs is in understanding the nitrogen cycle, how nitrogen from the atmosphere is converted to nutrients (and pollutants) and released back into the atmosphere. Many different classes of bacteria play key roles in this process. In soils, nitrifying bacteria’s fuel is ammonia, a nitrogen compound that comes from plant and animal waste but largely from commercial fertilizers. The bacteria extract energy from the ammonia and excrete what’s left in the form of nitrites, which other species of bacteria can further use.
Lancaster’s interest in this process began as a quest for models of how to control the release of energy. “When you eat a bowl of Wheaties, the energy gets released slowly. It doesn’t light on fire in your stomach, and it’s the same with these bacteria. They take the energy out in a controlled fashion,” Lancaster says. “If we can control these kinds of reactions, we can potentially store and transport the energy. But how do we control energy release from an energetic molecule? It’s a very tough problem. So we’re looking to the reactions in nature that do this.”
The bacteria’s ability to use ammonia relies on a metalloprotein, an enzyme with iron, a transition metal, at its core. Using many tools and techniques, a closer look at this enzyme immediately yielded surprising results. “We discovered very quickly that our understanding of what the reactions are has been incorrect,” Lancaster says. “Most recently, we’ve solved something that has been misunderstood for about 40 years.”
In the nitrification process, the bacteria first convert ammonia to an intermediate compound called hydroxylamine. The next step was thought to be that the iron-containing enzyme, hydroxylamine oxidoreductase, then helps finish the job, converting hydroxylamine to nitrite, the final product for these bacteria. “We wanted to understand how the iron takes the hydroxylamine and does this controlled release of energy,” Lancaster says, “but we could never make nitrite, not unless there was oxygen around.”
It turns out that there is another intermediate compound; the iron helps convert hydroxylamine to nitric oxide. Another enzyme entirely, which Lancaster’s group plans to report on soon, converts the nitric oxide to nitrite.
Nitric oxide is interesting for many reasons, one of which is that it seems to be everywhere in the nitrogen cycle, Lancaster says. “Now we know it plays this additional role. What it all points to is that nitric oxide plays a very central role in the environment and the speciation of nitrogen,” Lancaster explains.
His group has also found, significantly, that the nitric oxide made by these bacteria can react with another iron enzyme, cytochrome P460, and an additional molecule of hydroxylamine to make nitrous oxide. Nitrous oxide is a greenhouse gas with 300 times the warming potential of carbon dioxide and is also an ozone depleting agent.
“We think the nitrous oxide formation is a way these organisms can alleviate the burden of the nitric oxide toxicity. It’s like a bleed valve,” Lancaster says. With an overabundance of nitrogen from fertilizers in the soils, the organisms produce more nitric oxide and nitrous oxide, placing a heavy burden on the atmosphere.
Fully understanding these reactions is the first step to revising recommendations for agricultural practices and preventing increases in emissions. “We like the fact that this has societal consequences, and at the same time we can advance our understanding of very fundamental chemical concepts,” Lancaster says. “It’s been very exciting, and we have a tremendous amount of work left to do. There’s a lot of great chemistry that remains untouched.” Lancaster’s group is currently probing deeper into the enzymes involved in this process and what specific mechanistic role the transition metals play.
“I wanted to build a lab where we’re not limited…I want my students to be able to explore and have fun and to learn that science is about thinking outside the box and challenging established dogma.”
These projects are also spawning collaborations with soil scientists, microbiologists, and other experts, within and outside Cornell. “The Chemistry and Chemical Biology Department here is one of the most civil, collegial places you could possibly imagine, and what I’m discovering now is that this collegiality extends outside our department,” Lancaster says. “When you are part of such a rich and diverse institution as Cornell, we can really have some impact.”
Transition Metals and Oxidation States
Lancaster’s group also has a nonbiological side of the lab that tunnels all the way down to single electrons and an elemental debate. “One of the central concepts inorganic chemists love to argue about is the concept of an oxidation state,” Lancaster says.
Transition metals have been described as reservoirs of electrons in chemical bonding, losing electrons and entering into positive oxidation states. “But this is really not the whole picture,” Lancaster says, “because a transition metal is almost always supported by other atoms that we call ligands.” For example, the iron in hemoglobin—the protein that transfers oxygen in red blood cells so we can breathe—is supported by organic molecules called porphyrins.
Lancaster and others have argued that these supporting structures change the character of the metal, making the standard descriptions of oxidation states inaccurate. Many inorganic chemists recognize that metals themselves are not always coughing up electrons—instead the ligands around the atom are, making them something called redox non-innocent. Lancaster’s group has been showing that the list of innocent ligands—ligands that don’t give up electrons—is probably very short.
Lancaster’s group provided a surprising example in a compound called tetrakis trifluoromethyl copper monoanion, which had been referred to as a unicorn because it was thought to have a rare, highly-oxidized copper(III)—or a copper atom missing three electrons—but was nevertheless stable. “We used x-ray spectroscopy at the Cornell High Energy Synchrotron Source to show that it’s not a copper (III); it’s a far more common copper (I),” Lancaster says. “The ligands around it have coughed up the electrons. We’re trying to understand now whether that is just a curiosity, or is this something special that can be used to predict or design reactivity for metals like copper, nickel, and cobalt?”
What unites the synthetic and biological inorganic chemistry in Lancaster’s lab is the multifaceted approach, using many different tools. This style is inspired by Lancaster’s mentors and heroes in the field, as well as a desire to give his students the training he had with a wide range of experiences and tools to solve any problem.
“Sort of along the lines of the Ezra Cornell philosophy of building a place where anyone can study anything, I wanted to build a lab where we’re not limited,” Lancaster says. “We’ve got access to pretty much everything an inorganic chemist could want, and I want my students to be able to explore and have fun and to learn that science is about thinking outside the box and challenging established dogma.”