SCIENCE & TECHNOLOGY Berkeley, recounted this oft-told story at the August national meeting of the American Chemical Society in Boston. His reflections on his career in fluorine chemistry were offered at a special symposium commemorating his 70th birthday (and the 40th anniversary of the discovery of the first noble-gas compound). The symposium, sponsored by the Division of Fluorine Chemistry, showcased many topics in inorganic and organic fluorine chemistry, not just noble-gas chemistry
STARRY REACTION UV irradiation of a 1:1 mixture of gaseous xenon and fluorine at McMaster University yields crystals of XeF2 that adhere to the inside of the reaction vessel.
GASES OF A LESSER NOBILITY After 4 0 years, noble-gas chemistry is much less in vogue but can still yield surprising discoveries RON DAGANI, C&EN WASHINGTON
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d i n n e r t i m e , Neil Bartlett, working all alone in his laboratory, made the discovery of a lifetime. His discovery, which shattered the textbook axiom that the noble gases were inert and could not be induced to react with other chemical species, grew out of earlier work he did with graduate student Derek H. Lohmann. They established that the gas platinum hexafluoride (PtF 6 ) has such an extraordinarily high electron affinity that it can oxidize molecular oxygen to form the compound 0 2 + PtF 6 ~. Bartlett realized that the first ionization potential of 0 2 was very close to that of xenon, suggesting that PtF 6 might be able to oxidize xenon as well. He ordered HTTP://PUBS.ACS.ORG/CEN
some xenon and waited impatiently for it to be delivered. Finally, on March 23,1962, Bartlett was ready to carry out the experiment. In his lab at the University of British Columbia, Vancouver, he had assembled a glass and quartz apparatus in which the two gases were separated by a fragile glass seal. At about 6:45 PM, he broke the seal and observed an immediate reaction as the deep red P t F 6 gas mixed with the colorless xenon gas and a yellow-orange solid precipitated. "I was greatly excited by this observation and left the laboratory to find a colleague or student who could witness the event," he later wrote. But no one else was in the building! Bartlett, today an emeritus professor of chemistry at the University of California,
SHORTLY AFTER making his groundbreaking discovery, Bartlett dashed off a brief, four-paragraph preliminary report in which he identified the yellow-orange product of the gas reaction as xenon hexafluoroplatinate(V),Xe+PtF6~ [Proc. Chem. Soc, 1962,218]. But as he told his listeners in Boston, it took many, many years to figure out exactly what he had made. The reaction turned out to be more complicated than it seemed at first because the initial oxidation product, XePtF 6 , reacts with a second molecule of PtF 6 to yield XeF+PtF6~ plus PtF 5 . And upon warming to about 60 °C, these two products react to form XeF + Pt2F n _ . Furthermore, the initial product, which is amorphous and has the approximate composition XePtF 6 , turned out not to be Xe+PtF6~ as Bartlett had thought. Rather, he thinks it's probably an XeF+ salt of polymeric (PtF5~)„—a Pt(IV) species. According to chemistry professor Gary J. Schrobilgen of McMaster University, Hamilton, Ontario, the variability of the XePtF 6 product's composition, the possibility of mixtures, and the inability of researchers to obtain single crystals for Xray structure determination were serious obstacles on the path to characterizing the product. In fact, he says, the most definitive paper on the nature of the original X e P t F 6 compound was published by Bartlett and his colleagues only two-anda-halfyears ago [Coord. Chem. Rev., 197,321 (2000)]. In 1962, Bartlett and other scientists were reluctant to spend too much time on unraveling the intricacies of the Xe plus PtF 6 reaction when there was a whole new world of noble-gas chemistry to explore. At Argonne National Laboratory, which had pioneered the synthesis of PtF 6 , researchers quickly confirmed that it reacts with xenon. They then went on to react xenon with fluorine to prepare xenon tetrafluoride (XeF^. Soon, other xenon compounds, such as XeF 2 , XeF 6 , XeOF 4 , and X e 0 3 , were synthesized at a number of labs, including Argonne. C&EN
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SCIENCE & TECHNOLOGY Before too long, researchers also showed that krypton and radon react with fluorine to form KrF 2 , RnF 2 , and related compounds. In the burst of activity in the 1960s, the reactivity of xenon, krypton, and radon were explored very rapidly, according to Lawrence Stein, a retired inorganic chemist who investigated radon fluoride and its salts at Argonne. But the feverish activity eventually cooled off, and some labs, including Argonne, abandoned noble-gas chemistry
| have used frozen-gas matrices to I prepare and stabilize a host of oth< er noble-gas compounds including * HXeH,HXeCl,HXeBr,HXeCN, I HXeSH, HXeOH, HKrCl, HKrl CN,andthelatest-HKrF. £ These successes leave only two £ noble gases—helium and neon— S for which no neutral, covalent com£ pound has been demonstrated (ex5 eluding inclusion compounds in E which, for example, the noble gas * is encapsulated inside a fullerene ™ cage). Of the two unmarried noble * gases, helium has the higher ionization energy and therefore would seem to be the least likely to form a chemical bond.
THREE DECADES after Bartlett's discovery, "new results in this area had slowed down to a trickle, and in MILEST0NE When platinum hexafluoride, a red the minds of most chemists, the gas (left), is allowed to mix with a large molar excess chapter on noble-gas chemistry had of xenon, the immediately formed product is a yellow NEVERTHELESS, Rasanen specubeen completed," according to an solid with the composition XePtF6 (right)—the first lated on the possibility of forming article by inorganic chemist Karl O. recognized compound of a noble gas. HHeF. This molecule was predictChriste of the Air Force Research r ed on theoretical grounds to be Laboratory at Edwards Air Force 3 neutral and metastable by associBase in California [Angew. Chem. Int. 1 ate chemistry professor Ming Wah Ed, 4 0 , 1419 (2001)]. But in the I (Richard) Wong of the National past few years, he writes, a "burst of Z University of Singapore. A study by startling discoveries" has shown o chemistry professor R. Benny Gerthat noble-gas chemistry "is still ° ber of UC Irvine and the Hebrew full of surprises." £ University of Jerusalem and colli' leagues also concluded that it Some of those recent surprises would be metastable, with a lifewere recapped at the Bartlett symtime of picoseconds as a free molposium. For example, chemistry ecule. But according to Gerber, his professor Konrad Seppelt of the team's computations also suggestFree University Berlin discussed his ed that HHeF would be stable indiscovery, with graduate student definitely in pressurized solid heliStefan Seidel, of the first bulk comum. So Rasanen is hopeful that pounds with direct gold-xenon HHeF, if it can be induced to form, bonds. The first to be reported was would be amenable to spectro[AuXe4}2+([Sb2F11]")2, which conscopic study in a frozen helium tains a square planar arrangement REACTIVE ELEMENT Fluorine (shown in green) of four xenon atoms around a gold is a key component of several new molecular species matrix. atom [Science, 290,117 (2000)]. Not prepared and crystallographically characterized in HHeF is not the only noble-gas only is this the first example of mul- Schrobilgen's lab at McMaster University. Clockwise, species that inorganic chemists tiple xenon ligands, Christe notes, from upper left, are the Br03F2 anion, C(0TeF5)3+ have set their sights on. Schrobilbut it also represents the first strong cation, FXeOXeF, and FXeOXeFXeF+AsF6 . gen, for example, discussed his metal-xenon bond. Seppelt's group group's efforts to make the trifluothen revealed more gold-xenon chemistry roxenate(II) anion, XeF 3 . This anion ported the preparation and detection of by isolating salts of the cations cis- and trans- argon fluorohydride (HArF), the first would be formed if XeF 2 could accept a [AuXe 2 ] 2+ , [XeAuFAuXe] 3+ , and transfluoride anion. But unlike its cousins XeF 4 ground-state argon compound with cova[AuXe2F]2+ [Angew. Chem. Int. Ed, 41,454 lent bonds to argon [Nature, 4 0 6 , 874 and XeF 6 , XeF 2 has not been shown to be(2002)]. have as a fluoride-ion acceptor. However, a (2000)]. They made it by photolyzing hy19 F nuclear magnetic resonance specdrogen fluoride in a frozen argon matrix. Also on hand at the Boston symposium troscopy study by the McMaster group has The compound is expected to be quite rewas chemistry professor Markku Rasanen demonstrated that fluoride ion exchanges active with almost any species and appears of the University of Helsinki, who recentwith XeF 2 , with the likely intermediate beto be stable in the matrix only at very low ly drew worldwide scientific attention by ing XeF 3 ~. Also, mass spectrometric evitemperatures. ensnaring another noble gas in a molecule. dence for the anion has been reported. Two years ago, Rasanen and coworkers reBesides HArF, the Helsinki researchers
The recent spate of discoveries in noble-gas chemistry "may signal the beginning of a renaissance in this field." 28
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The reaction of the xenon teflate cation Schrobilgen's group made a number of with CBr 4 produces the expected CBr3+ attempts to prepare salts of the anion by salt, but with an extra wrinkle: The byreacting XeF 2 with fluoride-ion donors such as CsF, all without success. The XeF3~ product, BrOTeF 5 , reacts stepwise with the carbocation to replace all three anion's reluctance to form, he remarked, bromines with teflate groups. The final is consistent with the low fluoride-ion product of this series of reactions is the affinity of XeF 2 . CCOTeF^ salt, which is produced along Schrobilgen and coworkers had better with the CBr 3 + salt. Schrobilgen and luck in their quest for what he called "the coworkers have obtained the X-ray crystal missing oxide fluoride of xenon(II)"—the structure of the CCOTeF^ salt. V-shaped molecule FXeOXeF. Michael Gerken, formerly a graduate student in The group also is studying "mixed" Schrobilgen's group, previously had precations that contain chloro, fluoro, and/or pared a crystalline, red-orange salt of the teflate groups in different combinations. "chain" cation FXeOXeFXeF + by reacting For example, CFC1 2 + , the first perhalo+ XeF AsF6~ with water in a hydrogen fluogenated carbocation containing fluorine, ride solvent at - 7 8 °C. The X-ray crystal has been characterized in solution by 19F structure of this "rather exotic" zig-zagand 13C N M R spectroscopy, according to shaped cation revealed that the central Schrobilgen. "This is significant," he notXe-F bond is much longer than the moleed, "because fluorine actually serves to cule's three other Xe-F bonds. It appeared destabilize carbenium centers." For this as though XeF 2 was coordinated to the reason, he said, the CF 3 + cation, which xenon end of the FXeOXe+ cation, Schroeluded Olah's group, will be very difficult bilgen explained, "so we tried to displace to make. it." Sure enough, when graduate students Noble-gas compounds also have been Matt D. Moran and Bernard E. Pointner instrumental in the study of high-valent treated the chain cation with nitrosyl fluoride (FNO) at - 7 8 °C, XeF 2 was kicked out and the list of known xenon oxide fluorides gained a new entry: FXeOXeF. Although Schrobilgen's explorations of noble-gas and fluorine chemistry are fundamental in nature, some of his discoveries have synthetic applications. At the Boston meeting, for example, he unveiled some preliminary work on using a xenon "teflate" salt to prepare stable, isolable carbocations having three halogen substituents. Carbocations Bart Lett Schrobilgen such as CC13+, CBr3+, and CI 3 + were generated and observed in solution more halogen chemistry For example, xenon and than a decade ago by George A. Olah and krypton compounds have helped chemists coworkers at the University of Southern to attain the heptavalent state of bromine, California. But these carbenium ions were its highest oxidation state. In the late not isolated as salts, and their X-ray crystal 1960s, Evan H. Appelman and his colstructures have never been reported. leagues at Argonne synthesized the first Br(VII) species: perbromate (Br0 4 ~) and perbromyl fluoride (Br0 3 F). A few years NOWp BY USING the combination of the later, Schrobilgen, working as a graduate teflate cation XeOTeF 5 + and the weakly student in Ronald J. Gillespie's lab at coordinating Sb(OTeF 5 ) 6 ~ anion in McMaster, used krypton fluoride salts to S0 2 C1F solution at - 7 8 °C, Schrobilgen oxidize BrF 5 to give the hexafluorohas shown that stable trihalomethyl cations bromine(VII) cation, BrF 6 \ can be prepared and isolated in crystalline form. For example, the XeOTeF5+ salt reToday, Schrobilgen and graduate stuacts with CC1 4 to yield C10TeF 5 plus Xe dent John F. Lehmann are trying to exand CCl3+Sb(OTeF5)6~. The X-ray crystal tend Br(VII) chemistry even further, usstructure of this CC13+ salt confirms that ing perbromyl fluoride as a convenient the cation is planar and Y-shaped, Schrostarting material. The McMaster chemists bilgen said. This is the first crystal strucwondered, for example, whether B r 0 3 F ture of a perhalogenated carbenium ion, would be willing to donate a fluoride ion hetoldC&EN. to SbF 5 (a good fluoride-ion acceptor) to HTTP://PUBS.ACS.ORG/CEN
furnish the heptavalent Br0 3 + cation, which is unknown. As Lehmann noted in his own presentation at the Boston symposium, when he attempted to carry out this reaction at 0 °C, it took an unexpected course: The Br0 3 F substrate was reduced —losing oxygen gas as well as transferring fluoride —to yield Br0 2 + SbF 6 " Although the hoped-for Br0 3 + cation did not materialize, Schrobilgen later told C&EN, "we did squeeze out the crystal structure of Br0 2 + , which hadn't been done before." Lehmann has had more success exploiting perbromyl fluoride as a fluorideion acceptor. He found that Br0 3 F readily reacts with a variety of fluoride-ion donors to yield a new Br(VII) species, Br0 3 F 2 ", which he isolated as a salt and characterized by X-ray crystallography. BY CONTRAST, Lehmann said, the analogous reaction with C10 3 F just doesn't go under the same reaction conditions. C10 3 F has a weak fluoride-ion affinity and shows no inclination to form the difluoro anion. Schrobilgen pointed out that the E § B r 0 3 F 2 ~ anion is the first new I Br(VII) species to be synthesized 1 in more than 25 years. Perhaps g that's not surprising, considering ° that, as he put it, "all these compounds are a struggle" to make. The same could also be said of noble-gas compounds. Working with noble-gas and fluorine compounds can be hazardous and challenging, Schrobilgen noted, and not as many researchers are active in this field as once were. The Air Force Research Lab's Christe thinks that the recent spate of discoveries in noble-gas chemistry "may signal the beginning of a renaissance in this field." But Bartlett worries that practical difficulties in doing this sort of chemistry may make it difficult to sustain the momentum. Furthermore, he said in an interview with C&EN, "Who is interested in this sort of exotica these days? The funding agencies want to have some assurance that there's some practical outcome. And frankly, that's deadening for this kind of research because one can't easily foresee the benefits." Schrobilgen agrees. In the quest for immediate payoffs, he believes, "we're losing sight of the value of fundamental research" and the possibility of spin-offs. "I think we have to be more patient. And we have to let dreamers dream. They produce some amazing things if they're allowed to pursue their creative instincts." • C&EN
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