Among the many novelty periodic tables on the internet, one can find a jokey version that describes itself as “The Periodic Table: as seen by an organic chemist.” Carbon, beloved by organikers, looms large, blotting out some of the other elements. Meanwhile, the noble gases, on the far right, are given a far less flattering treatment, labeled as “Lazy elements.”
To call the noble gases lazy is unfair, even for an organic chemist. These elements work hard. Helium cools our magnetic resonance imaging scanners and nuclear magnetic resonance spectrometers. Neon signs illuminate the best dive bars. Argon provides chemists with an inert atmosphere in their glove boxes. Krypton flashes help photographers capture images at high speed. Xenon-fueled ion thrusters propel spacecraft like NASA’s Dawn to the far reaches of our solar system.
Instead, one might describe the noble gases as aloof. Because they’re reluctant to share electrons from their filled outer electron shells, noble gases are generally considered unreactive. But it is possible to wrestle reactivity from these elements, as the late chemist Neil Bartlett showed in 1962, when he made the first noble-gas compound, Xe[PtF6], by mixing xenon with platinum hexafluoride.
Making noble-gas compounds is not for the faint of heart, however. Because the electrons in the noble gases’ outer shells are comfortable where they are, it requires extremes—like reactive reagents, low temperatures, or high pressures—to get them to budge. When compounds do form, the results are seldom practical: most noble-gas compounds are too fleeting or unstable to be useful. But the few chemists who take on the challenge of coaxing reactivity from these recalcitrant elements say the true rewards are finding new insights into the nature of reactivity and chemical bonding.
Gary J. Schrobilgen, a chemistry professor at McMaster University, has been making compounds with xenon and krypton for more than 40 years. These, he points out, are the only noble gases that form stable compounds in quantities of a few milligrams or more—chemicals like XeF2, which can etch silicon, and KrF2, a strong but rarely used oxidizer.
Synthesizing such xenon and krypton compounds, Schrobilgen says, is deeply rooted in inorganic fluorine chemistry. “As such, it is among the most challenging areas of synthetic and structural chemistry,” he says. Many noble-gas compounds must be synthesized at low temperatures, requiring specialized techniques. And once made, they are sensitive to moisture and heat, and they’re strong oxidants, making them difficult to work with.
“Most of the equipment required for routine work is unavailable as off-the-shelf items and must be custom built,” Schrobilgen says. Many of the aggressive reagents and solvents used to make noble-gas compounds attack glass, so most of this chemistry must be carried out using metal vacuum lines, high-pressure reactors made from materials like stainless steel, or fluoroplastic reaction vessels.
Recently, however, chemists in Schrobilgen’s lab have been exploring noble-gas compounds with surprising stability. They’ve created complexes of XeO3 surrounded by coordinating ligands, such as crown ethers, which can lend stability to the compounds (Angew. Chem., Int. Ed. 2018, DOI: 10.1002/anie.201806640). On its own, XeO3 will detonate violently with the slightest provocation. But when XeO3 coordinates with 15-crown-5, the crown ether sits atop the xenon like a royal tiara, and the overall complex is stable at room temperature.
With crystals of the complex, “you can literally put them on a steel plate and whack them with a hammer and they won’t explode,” Schrobilgen says. If scientists could make this, or a similar XeO3 complex, in sufficient quantities and with a stable shelf life, it might find use as a clean oxidant, able to neatly pull electrons off molecular species during reactions. Elemental xenon, which can be recovered, would be the by-product.
Learning to make noble-gas compounds provides chemists with an excellent skill set, Schrobilgen says. Training them in complex techniques “gives them the ability to change as the landscape shifts beneath their feet.” And the landscape is changing. For example, Schrobilgen notes that it has become difficult to get undiluted F2 gas, a critical reagent for making noble-gas fluorides, which are often the precursors to other noble-gas compounds. The material is dangerous, and Schrobilgen suspects that companies selling it fear litigation.
“Although fluorine is produced in large quantities, companies now refuse to ship to individual researchers no matter what their level of experience may be,” Schrobilgen says. “This basically puts anybody who relies on 100% F2 out of business.”
Because of such difficulties procuring reagents, as well as high start-up costs and challenges securing funding, the number of chemists making noble-gas compounds is dwindling. “There’s such a paucity of researchers in the field at the moment, and those that are still around are rapidly aging and are going to retire,” he says. “They’re not being replaced.”
Some researchers interested in experimental noble-gas chemistry have moved away from making milligrams of isolable material to using spectroscopic methods to study noble-gas compounds under extreme conditions. Markku Räsänen, a professor at the University of Helsinki, thrilled the chemistry community nearly 20 years ago by making the first argon compound, HArF, by photolyzing HF in an argon matrix at a chilly 7.5 K (Nature 2000, DOI: 10.1038/35022551). Researchers in his lab have since made myriad other noble-gas compounds by cooling matrices of noble gases to temperatures as cold as 3 K.
“Noble-gas compounds are some of the best examples of understanding chemical bonding and knowing exactly how a bond forms and how strong the bonds are,” Räsänen says. For instance, “It’s interesting to think that if you ionize noble gases, then they’ll look like halogens.” If argon is stripped of one electron, he points out, it should behave like chlorine in the ways that it bonds to other atoms. “We know chlorine has a very rich chemistry.”
There are still unexplored frontiers in experimental noble-gas chemistry, Räsänen says. Experimental findings have attracted theoretical chemists to the field, he says, and their collaborations have “greatly deepened our understanding of the bonding and properties of these elements.” Räsänen thinks that the new species these chemists predict computationally will stimulate experimentalists to verify them.
Artem Oganov of Skolkovo Institute of Science and Technology has been combining computational work on noble gases with high-pressure experiments. “The rules of classical chemistry completely break down under pressure,” he says. For example, with enough pressure, potassium behaves like a transition metal rather than an alkali metal, and oxygen under pressure has superconducting properties.
When Oganov and his colleagues began exploring how noble gases would behave under pressure, they started with an element that no one had observed as a compound—helium. Computations suggested that a reaction with fluorine was most likely, so Oganov asked his graduate student Xiao Dong to give it a try. It didn’t work. Helium also didn’t combine with oxygen. So Oganov told Dong to abandon the project.
“Quietly, in secret, he kept trying to marry helium with other elements,” Oganov says. “He went from the right side of the periodic table, where I suggested he start, all the way down to the left part of the periodic table.” When Dong hit sodium, he found it forms a stable compound—Na2He—at 113 GPa, about a million times Earth’s atmospheric pressure (Nat. Chem.2017, DOI: 10.1038/nchem.2716).
“He came to me and he said, ‘Do you remember the project on helium chemistry?’ ” Oganov recalls. “I said, ‘Yes, I told you to stop this project because it wasn’t promising.’ He said, ‘Well, I didn’t listen to you.’ ” Oganov says his initial reaction was anger, but when he learned that Dong had found a compound, he was thrilled. That helium would bond with sodium, he says, shows that “even the most rational and well-rooted ideas of classical chemistry can become absolutely invalid when you come into the world of high pressure.”
The finding was completely counterintuitive, Oganov says. Helium’s behavior “came as a total surprise,” he continues. “Helium is the second most abundant element in the universe, and its chemistry, until very recently, was an absolute desert.” This work suggests that with enough pressure, helium could form compounds with other elements. Oganov and colleagues predict that the compound Na2OHe might be stable at pressures as low as 15 GPa, which is in the neighborhood of the amount of pressure used to convert graphite to diamonds. While they haven’t made that compound yet, Oganov thinks it’s possible and that it, or a similar compound, might one day be used to store helium in the solid state—something chemists once thought unimaginable.
Contradicting chemistry’s dogma is what keeps this small group of chemists trying to spur the aloof noble gases into action, even in the face of challenges to the field. Oganov says, “When you find something that contradicts traditional knowledge, it’s a chance to learn something fundamentally important about chemistry.”
This story was updated on May 30, 2019, to correct the credit on the image of the KrF2 synthesis apparatus.