Chemical & Chemical Engineering News (80th Anniversary Issue), Vol. 81, No. 36, 2003, Sept. Edited by X. Lu Introduction

NEIL BARTLETT, UNIVERSITY OF CALIFORNIA, BERKELEY

The Periodic Table of the Elements, as set out by Dmitry Mendeleyev or Lother Meyer, had not allowed for the discovery of a group of elements between the highly electronegative halogens and the electropositive alkalis. So the discovery of the unreactive monatomic gas argon by Lord Rayleigh and William Ramsay in 1895 came as a total surprise. James Dewar even indicated that the new gas could be an allotrope of nitrogen, N₃, a suggestion seconded by Mendeleyev! Within three years, however, Ramsay (and coworker Morris Travers) had also discovered helium, neon, krypton, and xenon. They established their monatomic and unreactive nature.

With the advent of the Rutherford-Bohr atom (1913), electron configurations of these unreactive monatomic elements soon came to have a central role in the emerging electronic theories of chemical bonding. In their 1916 papers, both G. N. Lewis and W. Kossel pointed to the electron configurations of these elements as especially stable. In each theory, the chemical properties of atoms of other elements were tied to the gain or loss of electrons from the configuration of the nearest monatomic gas. So successful were these theories in accounting for a wide range of chemical properties of the elements that the monatomic-gas electron configurations came to be thought of as chemically inviolate. This was fostered by early failures to make compounds of the gases (including an attempt by Henri Moissan to prepare an argon fluoride in 1895). Nevertheless, in his classic 1916 paper, Kossel made an astute observation relevant to the chemical reactivity of these elements.

On the basis of the first ionization potentials of the gases, Kossel noted that xenon was most likely to have the capability of forming fluorides and oxides. He also allowed that a krypton fluoride might be made. Similar predictions were made later, by Andreas von Antropoff (1924) and by Linus C. Pauling (1932), based on chemical trends in the periodic table. These predictions led D. M. Yost (with student A. L. Kaye) to attempt (in 1933) a xenon fluoride synthesis. That attempt failed. So matters rested until 1962.

It was the discovery of the remarkable oxidizing properties of platinum hexafluoride in making the salt O₂⁺PtF₆^– that led (via the recognition that O₂ and Xe have nearly the same first ionization potentials) to the oxidation of xenon by PtF₆. Later in 1962, Howard H. Claasen, Henry Selig, and John G. Malm, at Argonne National Laboratory prepared XeF₄. Syntheses of XeF₂, XeF₆, XeOF₄, XeO₂F₂, XeO₃, and perxenates (XeO₆^4– salts) were quickly reported from there and elsewhere. Even the highly unstable tetrahedral tetroxide XeO₄ was made (J. L. Houston, 1964). A fluoride of krypton, prepared and correctly identified as KrF₂ (George C. Pimentel and J. J. Turner, 1963), was first reported as KrF₄ (Aristid V. Grosse and coworkers, 1963), but no compound above Kr(II) has ever been established. Although the easier ionization of radon leads one to expect the most extensive chemistry for that element, the high instability of even the most stable isotope has severely limited studies of it. L. Stein, of Argonne, established (in 1962) the existence of a fluoride--probably RnF₂--but he and others were unable to confirm the existence of oxides or relatives of the perxenates.

The first ionization potentials of the noble gases provide a measure of how firmly the outer electrons are held by the effective nuclear charge. This hardness or softness of the valence electron set correlates well with the physical properties of the gases. The first ionization potentials are as follows: He, 24.6; Ne, 21.6; Ar, 15.8; Kr, 14.0; Xe, 12.1; and Rn, 10.7 eV.

Clearly, helium has the least polarizable electron cloud. This accounts for its low melting and boiling points and its low solubility in aqueous media, and hence its application in mitigating the bends as a diluent for oxygen in deep-water diving. In contrast, highly polarizable xenon has high solubility and is an excellent anesthetic. The low ionization potentials of the heavier gases also account for their chemistry.

In all of the known chemical compounds of the noble gases, the noble-gas atom has a net positive charge. We can take the difluorides as representative. In each, the noble-gas atom can be viewed as having lost an electron to the two F ligands. One way to picture this is to make the pseudo halogen Ng⁺ (transferring the electron to an F to make F^–), the Ng⁺ combining with the second F to make the classical electron-pair-bound [Ng–F]⁺ (Ng = noble gas). Each difluoride is then represented as a resonance hybrid of the canonical forms {F^–[NgF]⁺} and {[FNg]⁺F^–}. Bond energies in the isoelectronic relatives of the NgF⁺ species, ClF, BrF, and IF, do not vary greatly, so it is reasonable to assume the same to hold for the NgF⁺ species.

So, in making the difluoride from the constituent atoms via such species, the energy term that changes most from one Ng to another is the ionization potential. In support of this, we note that the heats of atomization (in kilocalories per mole) of XeF₂ (65) and KrF₂ (23) differ by almost the same energy as the first ionization potentials, about 44 kcal per mol. On this basis, since the first ionization potential of Ar is 42 kcal per mol higher than that of Kr, it can be supposed that the heat of atomization of ArF₂ would be less than that of KrF₂ by roughly that amount. Since the heat of atomization of KrF₂ is only 23 kcal per mol, this implies that ArF₂ cannot be made.

Because of the high electronegativity of the positively charged noble-gas centers, any ligands for those centers must themselves be electronegative. The ligands must also be small to provide high Coulomb energy. The known chemistry is in accord with this. For compounds that can be prepared and manipulated at ordinary laboratory temperatures, ligands are, so far, few. They are, for xenon: oxygen (S. M. Williamson et al., 1963; N. Bartlett, and coworkers, 1969), nitrogen (D. D. DesMarteau, 1981) and carbon (D. Naumann; H. J. Frohn, 1989); and for krypton (G. J. Schrobilgen, 1988–89): oxygen, and nitrogen. In all, the nitrogen and carbon ligands, to be effective, require linkage to other electronegative centers. The large size of the chlorine ligand means that XeCl₂ can only be made and used at cryogenic temperatures. The high bond energy of O₂ also leads to high thermodynamic instability of all oxides. XeO is only bound with respect to singlet oxygen, most likely 1D [O]. (E. H. Appelman of Argonne used XeF₂ in aqueous solution to prepare the first examples of perbromates, and this could have involved attack of bromate by XeO as a singlet oxygen carrier.)

Because of its lower effective nuclear charge, xenon is more easily oxidized than krypton. It is also a better Lewis base. Krypton is not oxidized by PtF₆, and in the solvent anhydrous HF (aHF), forms no complexes with transition-metal ions as xenon does. The HF-solvated Ag²⁺ ion (silver has the highest second ionization potential of any transition metal) oxidizes Xe at ordinary temperatures (Zemva et al., 1990) to make XeF₂. With the solvated Au²⁺ ion, however (Konrad Seppelt, 2001– 03), no oxidation of Xe occurs. Xenon is a sufficiently good Lewis base, however, to coordinate with Au²⁺ to form a variety of coordination complexes. It similarly complexes with Au³⁺ and with other metal ions. It is possible that Kr and Ar could also act as Lewis bases, but evidently they cannot compete with HF.

Many recent findings, including the first evidence for an argon compound, have come from matrix-isolation studies at the University of Helsinki in Finland (Markku Räsänen and coworkers). These studies have established the existence of a large variety of novel compounds, all stable up to 40 K. Included are HXeOH, HXeCCH, HKrCN, HKrCCH, and HArF. The last requires comment, because of the nonexistence of ArF₂.

In all of these compounds, the vibrational spectroscopic findings indicate that the canonical form ([HNg]⁺Y^–) contributes importantly to the binding of the molecules. The tiny proton is highly electronegative, and it bonds covalently to Ng in these molecules. The proton affinities of the noble gases are the following: He, 1.8; Ne, 2.2; Ar, 3.0; Kr, 4; and Xe, 6 eV.

Attachment of a proton to the more polarizable gases therefore gives significant energy toward bonding. Of course, this requires a small electronegative coligand, of which F has no superior; but OH, CN, and CCH are also small ligands with high electron affinities.

All of the noble-gas compounds are easily reduced, and the compounds can effectively carry their ligands as radical species, readily available to more reactive elements. For example, KrF₂, which is less bound than F₂ itself, will oxidize Xe to XeF₆! More electronegative cationic species such as KrF⁺ are even more potent (for example, the synthesis of BrF₆⁺, Ronald J. Gillespie, and Schrobilgen, 1974). Some of the potential reagents are very fragile and will need to be used at cryogenic temperatures, but the elimination of an atom of an almost inert gas can simplify the chemistry.

Neil Bartlett is an emeritus professor in the department of chemistry, University of California, Berkeley, and an emeritus senior scientist in the Chemical Sciences Division of Lawrence Berkeley National Laboratory. Bartlett succeeded in preparing the first noble-gas compound in 1962.The Royal Society (London) awarded him its Davy Medal in 2002.

COURTESY OF NEIL BARTLETT START OF SOMETHING NEW Bartlett's oxidation of xenon by PtF6, shown here, sparked the field of noble-gas chemistry.

THE NOBLE GASES AT A GLANCE

Name: Noble gases, inert gases. The six noble gases are found in the far right column of the periodic table. Helium's (He) name originates from the Greek helios, meaning sun; neon (Ne) from the Greek neos, meaning new. Argon (Ar) derives from the Greek argon, meaning inactive; krypton (Kr) from the Greek kryptos, meaning hidden. Xenon (Xe) comes from the Greek xenos, stranger, while radon's (Rn) name hails from the element radium.

Atomic mass: He: 4.00; Ne: 20.18; Ar: 39.95; Kr: 83.80; Xe: 131.29; Rn: (222).

History: Helium was discovered on Earth in 1895 by Scottish chemist Sir William Ramsay, though some credit the discovery of helium (in the sun's spectrum) to Janssen and Lockyear in 1868. Ramsay discovered most of the remaining noble gases--argon in 1894 (with Lord Rayleigh) and krypton, neon, and xenon in 1898 (with Morris M. Travers). Radon was discovered in 1898 by Fredrich Ernst Dorn.

Occurrence: Helium comes from natural gas deposits within Earth and can be isolated from air. Helium has the lowest boiling point of any element--4.2 K. Neon, argon, and krypton are obtained from fractional distillation of liquid air. Radon comes from radium decay.

Behavior: These elements were considered to be inert gases until the 1960s, because their oxidation number of zero prevents the noble gases from forming compounds readily. All noble gases have the maximum number of electrons possible in their outer shell (two for helium, eight for all others), making them highly stable.

Uses: Helium is used in balloons and in deep-sea diving to dilute oxygen that divers breathe. Neon, argon, and krypton are used in lighting. Radon is radioactive and hazardous.