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A Basic Rule of Chemistry Can Be Broken, Calculations Show

A study suggests atoms can bond not only with electrons in their outer shells, but also via those in their supposedly sacrosanct inner shells
illustration of compounds formed through inner-shell electron bonding between cesium and fluorine



Maosheng Miao

Most of us learned in high school chemistry class that chemical bonds can only form when electrons are shared or given away from one atom’s outer shell to another’s. But this may not be strictly true. A chemist has calculated that under very high pressure not just the outer electrons but the inner ones, too, could form bonds.

Inside atoms, electrons are organized into energy levels, called shells, which can be thought of as buckets of increasing size that can each hold only a fixed number of electrons. Atoms prefer to have filled buckets, so if their outer shell is missing just one or two electrons, they are eager borrow form another atom that might have one or two to spare. But sometimes, a new study suggests, atoms can be incited to share not just their outer valence electrons, but those from their full inner shells. “It breaks our doctrine that the inner-shell electrons never react, never enter the chemistry domain,” says Mao-sheng Miao, a chemist at the University of California, Santa Barbara, and the Beijing Computational Science Research Center in China. Miao predicted such bonds using so-called first-principles calculations, which rely purely on the known laws of physics, and reported his findings in a paper published September 23 in Nature Chemistry. Such bonding has yet to be demonstrated in a lab. Nevertheless, “I’m very confident that this is real,” he says. (Scientific American is part of Nature Publishing Group.)

His calculations show that two possible molecules could form between cesium and fluorine atoms under extremely high pressure—about 30 gigapascals (higher than the pressure at the bottom of the ocean, but less than at Earth’s center). Cesium, all the way on the left side of the periodic table, has one superfluous electron in its outer, or sixth shell. Fluorine, on the other hand, is toward the far right of the table, just next to the column of noble gases with completely full shells (which is why noble gases are notoriously unreactive—they have little incentive to gain or lose electrons) and is one electron short of a full outer shell. “Under normal pressure, cesium gives an electron completely to fluorine and they bind together,” Miao says. “But under high pressure, the electrons from cesium’s inner shells start to form molecules with fluorine.”

Miao identified two compounds that could form and remain stable up to very high pressures: cesium trifluoride (CsF3), where cesium has shared its one valence electron and two from an inner shell with three fluorine atoms, and cesium pentafluoride (CsF5), where cesium shares its valence electron and four inner-shell electrons to five fluorine atoms. “That forms a very beautiful molecule, like a starfish,” Miao says. Both the shape of the resulting molecules and the possibility of their formation are “very surprising,” says chemist Roald Hoffmann, a professor emeritus at Cornell University, who was not involved in the calculations. “This is the first clear case of an alkali metal not only losing its single easily ionized valence electron in bonding, but also ‘breaking into the core’ in its bonding with several fluorines.”

The reason these reactions may occur has to do with enthalpy—a measure of the total energy of a system at constant pressure. Chemical reactions tend to move toward products with lower enthalpy. Miao calculated the enthalpy of cesium fluoride (the basic one-to-one bond of two atoms that forms naturally) and the enthalpy of the possible compounds cesium trifluoride and cesium pentafluoride. He found that above certain pressure thresholds, these larger molecules had lower enthalpy—and therefore were likely to form. “Pretty much everything we see in terms of structure and bonding is merely a manifestation of the system finding ways to minimize the potential energy, balancing the energy gain from the energy cost,” says Purdue University chemist Paul Wenthold, who was not involved in the study. “Although inner-shell electron oxidation is not normally expected for something like cesium, if you put it in close enough proximity of something with a high enough propensity to accept electrons, sure enough, it will do it.”

So far, no one has attempted experiments to make these molecules in a laboratory, but Miao says it should be possible, although fluorine is difficult to work with. The pressures necessary are well within reach of modern equipment. “Experimental studies on these systems will provide an excellent opportunity to calibrate theoretical chemistry of systems at high pressure,” says chemist Lee Sunderlin of Northern Illinois University, who also was not involved in the study. “This is an excellent example of a theoretical approach that can guide experimentalists in choosing systems that should exhibit unprecedented properties.” Now that this basic rule of high school chemistry has been overwritten, who knows what other molecular surprises could be in store?

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