The chemistry of thenoble gases
N. Bartlett
The conclusive establishment, in r 962, of the existence of compounds of the noble gases has
overthrown a long-established chemical belief. This article reviews the preparation andchemistry of these compounds. The molecular structures of the simple xenon compounds areseen to correspond to those of iodine compounds having similar stoichiometries and elec­
tronic configurations. The available experimental evidence indicates that the nature of thebonding is similar to that found in other, less exotic, compounds. The article concludes withan assessment of the range of compounds the various noble gases can be expected to form.
Preparation of the first true noble-gas compounds

The first true compounds of the noble gases were prepared in
I962 [ I -3], more than sixty years after the discovery of thegases themselves by Sir William Ramsay. Although manyclaims to have made compounds in which noble gas atoms
were chemically bonded to other atoms had appeared earlier,none had been conclusively established. Clathrate compoundsof the noble gases had, however, been known for a number ofyears; indeed, P. Villard prepared the first, a hydrate of argon,
shortly after the discovery of the gas itself [4]. However, there
is in such compounds no chemical bonding of the noble gas,
the molecules of which are simply caged by the host lattice.
Apparently Moissan, a friend of Ramsay, tried without suc­cess to bring about a reaction between argon and fluorine [5].
In his later years, Ramsay was evidently convinced of thechemical inertness of the gases, because he rejected a suggestionby G. Oddo [ 6] that krypton and xenon ought to form halides.
Similar predictions were made by A. von Antropoff [7], whowent so far as to predict a possible octavalence for the heavier
gases; he and his co-workers attempted to make compounds,
but they had no unequivocal success. In 1933, L. Pauling [8]saw that the heavier gases were the more likely to form com­pounds, and that fluorine and oxygen were the elements most
likely to enter into combination with them. He suggested that
hexafluorides of krypton and xenon, and an octafluoride of thelatter, might be made. In the same prophetic communication
he also predicted the existence of an acid H 4Xe0 6 and salts ofoctavalent xenon, such as Ag4Xe06 • In the same year, un­
successful attempts were made to synthesize fluorides of argon
and k rypton by 0. Ruff and W. Menzel [9] and of xenon by
D. M. Yost and A. L. Kaye [ro]. The latter workers, experi­
menting with electrical discharges through mixtures of xenon
and fluorine, came closest to success. Unfortunately, they con­ducted their experiments in quartz apparatus. They concluded
that the etching of their apparatus might have been due to theaction of a xenon fluoride, and that the experiments did not
rule out the existence of such a compound. Had their apparatus
been efficiently cooled, or constructed of fluorite or alumina
instead of quartz, the xenon fluorides might well have been
discovered thirty years ago. In the event, the failure to confirm
the existence of any of the predicted compounds, and the successof the early electronic theories of valence, which stressed thestability of the inert-gas configuration, combined to entrench
the conclusion that the noble gases were incapable of chemical
reaction.

A clear indication that it should be possible to oxidize at least
the heavier noble gases to produce salts arose from the discoveryof the remarkable oxidizing properties of platinum hexa­
fluoride [rr]. N. Bartlett and D. H. Lohmann had establishedthat this gas combined with molecular oxygen in a r : r molar
ratio, forming a red solid. As prepared at low temperatures,
this compound proved to be isomorphous with potassium
hexafluoroplatinate (v), K[PtF6]. (The symbol (v) refers to thevalency of, in this case, the platinum). When formed at ahigher temperature, it is isomorphous with potassium hexa­
fluoroantimonate (v), KSbF 6 • This, together with the magnetic
and chemical properties of the oxygen compound, left no
doubtthatitwasdioxygenylhexafluoroplatinate (v), 0 2 +[PtF 6 )-.
Since the first ionization potential of molecular oxygen (r2 · 2 eV)is almost the same as that of xenon, this element also appeared
to be within the oxidizing range of platinum hexafluoride. The
lattice energy of a salt is generally inversely proportional to thesum of the ionic radii, but the lattice energy of the hexafluoro­
platinates is rather insensitive to cationic radii changes, because
the anion is large. A computation made for the hypotheticaloxidized xenon compound Xe +[PtF 6]- showed that the latticeenergy would be little more than IO kcal less than that of the
dioxygenyl salt. The predicted interaction of xenon andplatinum hexafluoride was confirmed in a simple and visually
dramatic experiment [r]. When deep red platinum hexa­
fluoride vapour was mixed with an excess of xenon gas at room
temperature, a yellow solid was immediately precipitated(figurer). The quantity of xenon consumed was commensuratewith a combining ratio of I : r, and accordingly the compound
was formulated as xenon hexafluoroplatinate (v), Xe +[PtF .J-.
This work aroused much interest at the Argonne National
Laboratory, because platinum hexafluoride and the otherpowerfully oxidizing hexafluorides RuF 6 and RhF 6 were firstmade there [r2]. The Argonne workers confirmed the reaction
between xenon and platinum hexafluoride, and then investi­
gated the reaction of xenon with other hexafluorides. Theyobserved that ruthenium hexafluoride reacted with xenon, butthat more than two moles of the hexafluoride reacted with eachmole of xenon [ 3]. Further, some reduction of the ruthenium.
seemed to occur. This suggested that the hexafluoride was,at least in part, acting as a fluorinating agent, and that afluoride, or some fluorides, of xenon existed. This supposition
was very quickly proved correct.

N. Bartlett, B.Sc., Ph.D., F.R.I.C.
Was born in 1930 and educated at King's College, Durham. After a year as
senior chemistry master at the Duke's School, Alnwick, Northumberland,
he became a lecturer in chemistry at the University of British Columbia,
where he is now an associate professor. His prime research interests are
in inorganic fluorine chemistry, particularly the structural aspects of thesubject, and in correlating molecular structure of new compounds with
their electronic configuration.

By heating five volumes of fluorine with one volume of xenon to,
400° C in a nickel can, followed by rapid cooling in cold water,
H. H. Claassen, H. Selig, and J. G. Malm [ 2] succeeded in
making the first fluoride of xenon, xenon tetrafluoride (figure2). 'I'\ ithin a few weeks, they and their collaborators at Argonne

Preparation and chemistry of xenon fluoridessFigure 1 The xenon-platinum
hexafluoride reaction. Left: Platinum
hexafluoride vapour separated fromxenon by a glass diaphragm. Right: The
product (XePtF.) of the xenon, platinum
hexafluoride reaction immediately after
the breakage of the glass diaphragm.
Figure 2 Crystals of xenon tetrafluoride,
seen against a blue background.

useful as such, because the reduction product, being xenon, isunlikely to contaminate the fluorinated material. The hexa­
fluoride is thermally less stable than the tetrafluoride, and much
more reactive. It appears that the tetrafluoride may be storedindefinitely in dry silica vessels, whereas the hexafluoride reacts
rapidly to produce xenon oxide tetrafluoride:
2XeF 6+SiO 2 ➔ SiF 4+2XeQF 4

2

National Laboratory had also prepared the difluoride and thehexafluoride. The existence of a difluoride was first indicated bymass-spectrometric studies [3] on impure tetrafluoride samples.
It was soon produced quantitatively by irradiating a xenon/
fluorine mixture with light from a high-pressure mercury arc,
the product being condensed in a cooled limb of the apparatus
as soon as it was formed [13]. The hexafluoride was preparedby heating xenon with a large excess of fluorine at 300 ° 0. That
independent reports of the preparation of the difluoride [14]and of the hexafluoride [15] appeared simultaneously within
three months of the announcement of xenon tetrafluoride
testifies to the great interest aroused by that discovery.
At room temperature, all three xenon fluorides are white
crystalline solids. The hexafluoride is the most volatile, having
a vapour pressure of about 30 mm Hg at 25° 0. It melts at 46° 0,
and the vapour is, surprisingly, pale yellow. The difluoride andtetrafluoride are less volatile, each having a vapour pressureof about 3 mm Hg at 25° 0. The xenon fluorides are soluble inliquid hydrogen fluoride, but only the hexafluoride solution isappreciably ionized, presumably as a result of the reaction [16]HF+XeF 6 ➔ XeF 0 +HF 2 +The tetrafluoride solution fluorinates bright platinum toplatinum tetrafluoride. All the fluorides ought to be good
fluorinating agents, and it is possible that they may prove to be
4

Compounds of radon and krypton
Efforts were also made at this time to prepare compounds ofother noble gases. Radon, in particular, would be expected toform compounds even more readily than xenon. Despite con­siderable experimental difficulties, P. R. Fields, L. Stein, andM. H. Zirin [1 7] have established the existence of radon
fluoride. Although radon fluoride is apparently much lessvolatile than any of the xenon fluorides, the very small scale ofthe preparation does not exclude the possibility of the entire
sample being either adsorbed on the nickel reactor surface, or
else chemically combined with the protective coating of nickel
fluoride. Large-scale preparations, which are technicallydifficult because of the radiation problem, are in hand andshould solve this problem.
A. V. Grosse and his co-workers [18] were able to isolate
krypton tetrafluoride by passing an electric discharge throughkrypton/fluorine mixtures at liquid-nitrogen temperature. The
compound is much less stable than its xenon analogue, rapidly
decomposing into the elements at 60 ° 0. It is a transparent
colourless solid, having a vapour pressure of 115 mm at 25 ° 0.
J. J. Turner and G. 0. Pimentel [16] photolysed a solid
mixture of fluorine and krypton in an argon matrix at 20° K andobtained a chemical species which infra-red evidence showed to
be k rypton difluoride. This method, known as the matrix
isolation technique, gave no evidence for the formation of anargon fluoride when a mixture of solid argon and fluorine wasirradiated alone.
The molecular structures of the fluorides
X-ray and neutron-diffraction studies of the solid xenon
difluoride and tetrafluoride have been made [19], and X-ray
analysis has revealed a 1 : 1 adduct of the di- and tetra-fluorides[20]. In the case of xenon hexafluoride, neither the X-ray nor
the neutron-diffraction studies are complete. Infra-red andRaman spectra of the two lower fluorides have been unam­
biguously interpreted [21], but the hexafluoride spectra,
although superficially similar to the spectra of molecules with
Oh symmetry, exhibit irregularities suggestive of lower mole­cular symmetry [16]. In both the solid and gaseous state, thedifluoride is a linear molecule. In the solid, the molecules lie
parallel to one another in a body-centred tetragonal latticearray; the Xe-F distance is 2·00 ± 0·01 A. Xenon tetrafluoride
molecules show square planar coordination of the xenon by thefluorine atoms, in both the solid and gaseous states. The angle
between adjacent fluorine ligands is 90 °, and the bond lengthis 1·953±0·002 A. Infra-red evidence [16] shows that kryp ton
difluoride has a molecular structure similar to its xenon ana­
logue, and it is probable that the tetrafluoride will prove to havethe same structure as xenon tetrafluoride.
Oxyfluorides and oxides
Xenon oxide tetrafluoride, which results from the reaction of thehexafluoride with glass, is best made [24] by simply mixing thehexafluoride with an equimolar quantity of water:
XeF 0+H 2O ➔XeOF 4+2HF
It is a clear, colourless, mobile liquid (m.p. -28 °C). The
infra-red and Raman spectra have been interpreted [16] on thebasis of a model having C 4 v symmetry, the oxygen atom beingat the apex of the square pyramid and the xenon atom being
almost coplanar with the four fluorine atoms.
With the fluorides well established, attempts were promptly
made to prepare the oxides. Although synthesis from theelements proved to be fruitless, and pointed to the oxides being
endothermic, there was an indication that a hydroxide oroxide was produced in the hydrolysis of the tetrafluoride [2].
An oxide was in fact obtained by the hydrolysis of this fluoride,and proved to be the trioxide [22]. When the first samples wereobtained there was some uncertainty about the nature of the
product. However, it was proved to be the trioxide by X-ray
analysis [23], which showed that the material closely resembled
iodic acid. Like 10 3 -, the molecule is trigonal pyramidal; theXe-O bonds are short (1 ·76 A) and so must have appreciable
double-bond character. The stoichiometry of the reaction isrepresented approximately by the equation
6XeF 4+12H 2O ➔4Xe+2XeO 3+3O 2+24HF
Although the mechanism is not understood, the xenon trioxide
clearly arises by disproportionation of solvated Xe (rv) species.
The disproportionation resembles those of the halogen oxy­anions, for example:
3CJO.- ➔Cl-+2ClQ 3 The trioxide is best made by the slow and complete hydrolysisof the hexafluoride [24]:
XeF 0 +3H 2 O ➔XeO 3+6HF
It is non-volatile and dangerously explosive, being comparable
in the latter respect to T.N.T. The explosive nature of thecompound is compatible with its being highly endothermic( !::,.Hf ° cryst. = +96 kcal mole- 1).
Some thermodynamic considerations
The thermal cycle

shows that the formation of a compound from atoms of anoble gas (N) and of a combining element (X), will be exo­
thermic if the average bond energy of the compound exceeds
the enthalpy of formation of the gaseous atom X. For thermo­
dynamically stable compounds, under ordinary conditions oftemperature and pressure, the average bond energies should
therefore exceed, for fluorides, 18·3 kcal; for oxides, 59·2 kcal
(double bond); and for chlorides, 29·01 kcal. From the thermo­
dynamic data obtained for the tetrafluoride, the hexafluoride,and the trioxide (Table r), the average bond energies for thetwo fluorides are seen to have the same value of 32 kcal, within
experimental error. The average energy of the oxygen-xenon
bond, which we can presume to be double, is less than this.
If the average bond energies in the oxyfluorides are giventhese values, then xenon oxide tetrafluoride should be an
exothermic compound with 6Hf° ( g l= -{[(4 x 32)+28]-[,p 18·3(6Hf 0 ( g l(F))+59·2(6Hf° ( g l(O))]}= - 24 kcal mole- 1
Table 1

Thermodynamic data for XeF4 , XeF 0 , and XeO,
Average°

Compoundl::,.Hf cryst.

l::,.Hsub.

/::,.Hf gaseous bond

XeF 4
XeF 6
XeO,

-as•

-701 ; -75 2

15 3
14.9•

128; 133
183
<84

°

+95 4

atoms

energy32-33 kcal
30·5
<28

1 L. Stein and P. Plurien [16].
'S. R. Gunn and S. M. Williamson [29].
3 J. Jortner, E. 0. Wilson, and S. A. Rice [16].
4 C. L. Chernick [16].
• S. R. Gunn [16].

However, xenon dioxide difluoride should be endothermic,since
6Hr 0 (g )= -{[(2 x· 32)+(2 x 28)]-[(2 x 18·3)+(2 x 59·2)]}= +35 kcal mole-1
This oxyfluoride, if it can be prepared, should prove to be athermally unstable compound.
It is interesting that a plot of average bond energy against
atomic number for the series of fluorides SbF 3, TeF 4, IF 5,
XeF 6 shows a smooth relationship\ as represented in figure
3. The values for the highest fluorides of antimony, tellurium,
and iodine lie on a similar smooth curve. Extrapolation showsthat the average bond energy of an octafluoride molecule
ought to be approximately 24 kcal, which would mean astandard heat of formation of the gaseous molecule of -45 kcal
mole - 1• It is possible that]. Slivnik and his co-workers [16],who have described a yellow product which could be pumped
from a bath at -78 ° C but was retained at -196° C and had
a fluorine to xenon ratio of 7·9 : 1, have already obtained thiscompound. The volatility is, however, greater than one would
expect for an octafluoride.
The difficulty of formation and the thermal instability ofkrypton tetrafluoride are consistent with an average bondenergy value of 18-19 kcal. This is obtained by extrapolationfrom values calculated for the fluorides of arsenic, selenium, and
1 The decrease in average bond energy, with increasing coordina­tion number, does not mean that the intrinsic bond energy isdecreasing; in fact, the reverse may be true. For the derivation of
intrinsic bond energies, the valence-state promotional energies (which
would be expected to increase with coordination number) for thereacting gaseous species would need to be known.
5

80
oi 60
u
,::t:

40

SbF,

AsF,

100

SbF,

SeF,

t

i

ABrF
BrF,; BrF,
Exothermic,
noble gas fluorides

IF,

KrF,
• [KrF ,] ; KrF,

[XeF,]

20 ,....,------------•·------------

Endothermic,
noble gas fluorides

Figure 3 Average bond energy (kcal) for some fluorides.
Square brackets indicate unknown compounds, open circles,
experimental values; closed circles, hypothetical values.
bromine (figure 3). In VIew of the greater bond energy ofbromine monofluoride (59·4 kcal) compared with the trifluoride
and pentafluoride (48·3 and 47.7 kcal respectively), it ispossible that the average bond energy for kryp ton difluoride
may be appreciably greater than 19 kcal, making this a ther­mally much more stable compound than the tetrafluoride. It
is of interest that Turner and Pimentel [16], by their matrix
isolation technique, obtained evidence for the existence of only
the difluoride.
For oxides and chlorides of the noble gases to be formed
exothermally, the average single-bond energy needs to begreater than ~ 30 kcal because t:,.Hf(g) 0 (0) = 59·2 kcal/mole andt::.Hr(g) 0 (Cl) = 29·01 kcal/mole. Consideration of the averagebond energies, estimated from the available thermodynamic
data for the oxides and chlorides of antimony, tellurium, andiodine (figure 4), shows, however, that the bond energies forxenon oxides and chlorides should be less than this. Con­
sequently, the xenon oxides and chlorides should be endo­
thermic compounds. This has been confirmed for the trioxide,
which exists because of kinetic stability. It is therefore notsurprising that attempts to make chlorides by synthesis fromthe gases failed, and an attempt [25] to prepare the chloride byinteraction of boron trichloride with xenon tetrafluoride at- 78 ° also failed:
3XeF,+4BCl. ➔ 3Xe<g>+6Cl 2 +4BF 3 (g)shape of xenon tetrafluoride is the same as that of the iso- or
pseudo-isoelectronic ions [ClF,J-, [BrF,J-, and [ICl,J- andthe shape of the difluoride the same as [ICl 2J-, it is to be
expected that XeF 6 will be similar to the [IF 6 J- ion. Unfor­
tunately, the structure of the latter is unknown. There is,however, evidence that the non-bonding electron pair issterically active, because the hexafluoroiodates (v) are notisomorphous with those hexafluorometallates that do not con­tain ·non-bonding valence electron pairs, for example KSbF 6•
Simple valence electron-pair repulsion theory predicts mole­cular shapes for the noble gas compounds that are consistent
with the experimental findings (figure 5) [16]. Thus thedifluoride is predicted to be a linear molecule having threenon-bonding electron pairs at the points of an equilateraltriangle, situated in the equatorial plane of the molecule. The
tetrafluoride is predicted to be square planar, having one non­bonding electron pair above and one below the plane, both on
the fourfold axis of the molecule. The hexafluoride should be aslightly distorted octahedron, having a non-bonding electron
pair at the centre of a face or at the mid-point of an edge. In
the xenon oxide tetrafluoride molecule, the doubly bonded
oxygen is predicted to lie on one side of the plane containing
the xenon atom (which is square coordinated by four fluorine
atoms) and the non-bonding electron pair on the other. The
fluorine atoms and the xenon are predicted to be almost
coplanar, because of the near equality in repulsive power of adouble bond and an electron pair. The trigonal pyramid shapeof the XeO 3 molecule is seen to be consistent with the presenceof a non-bonding electron pair at the apex of the pyramid, thispair repelling the bonding electrons (four in each Xe=O bond)to produce the pyramid.
In this treatment, which is not capable of being madequantitative, all the valence electrons are treated equally, and
it is implicit in the theory that d-orbitals are involved in thebonding. R. E. Rundle [26] andj. Hinze and K. S. Pitzer [16],however, have argued that d-orbital involvement should besmall. Indeed, the most popular theoretical model has been
the simple molecular orbital scheme; this uses p-atomic
orbitals and was first used by Pimentel [27] to discuss inert gas

The chlorides may not be kinetically stable, and hence it may
not be possible to prepare them under ordinary conditions oftemperature and pressure.
The nature of the bonding and of the molecular shape

There is no doubt that the chemical bonding in the noble gas
compounds is similar to that found in other compounds. In
particular, the bonding must be similar to that found in theinter-halogen compounds, the halogen oxides, and otherhalogen compounds. For example, the bond lengths, bondstrengths, and molecular shapes of the fluorides of xenon arecomparable to those of iodine fluorides. Since the molecular
100
80
'i 60
u
,::t:

I

F

F

F

,

F�

Noble gas oxides

l�F

Xe

>90 °

�F

,

F

or

Sb,O,

F

F

l'!;go;'Xe �020[XeCI,]

Figure 4 Average bond energies (kcal) for some oxides andchlorides. Square brackets indicate unknown compounds; open
circles, experimental values; closed circles, theoretical values.

F.,

I

F
°

p·�, �F

I

0

TeO,

ltoo�

'.,ii(·- Xe <90

F

40

6

\. �-

F- Xe -F

F

�F

;,
Xe

°

0

&~109 0 0

0

Figure 5 Molecular shapes predicted by simple valence
electron-pair repulsion theory [30].

compounds qualitatively. For the xenon fluorides, Rundle has
proposed linear, three-centre, four-electron bonds (F-Xe-F)
in which two electrons occupy a bonding molecular orbital,
4'µ, and two occupy a non-bonding molecular orbital, '-1'6· The
latter places the electrons mainly on the fluorine atoms, andconsequently the bond must be considered to have some ioniccharacter. Since this situation allows only one bonding pair ofelectrons for the two F-Xe bonds, the bond length is predicted
to be greater than for a conventional electron-pair bond. This
theory allows for the existence of the three xenon fluorides,
XeF., XeF 4, and XeF 6, and correctly predicts the molecularshape of the di- and tetrafluorides, Since the linear F-Xe-Fbonds should be approximately at right angles, it predicts anoctahedral XeF 6 •As has been pointed out by L. C. Allen [16], the contro­
versial theoretical situation should become clearer when the
structure of xenon hexafluoride is known. Since there is an
accumulating body of evidence to show that this compound is
not octahedral, the case for d-orbital involvement in thebonding grows stronger. The preparation of an octafluoride
would also prove this involvement.
Octavalent xenon

Although the octafluoride is not yet known, the anticipatedoctavalence of xenon [8] has been attained with the isolationof the perxenates. The reaction of xenon hexafluoride withsodium hydroxide was shown by J. G. Malm and his colleagues[16] to give white crystalline solids which contain xenon in the+ 8 oxidation state. Several hydrates were stated to exist, but theone that crystallizes first has recently been shown by W. C.
Hamilton, J. A. Ibers, and D. R. Mackenzie [28], by X-ray
structural analysis, to be Na 4XeO 6.8H 2O. The perxenate ion(XeO 6) 4- is nearly regularly octahedral, having a mean
Xe-O bond length of 1·875 A. This bond is 0·12 A longer than
the average Xe-O distance in XeO 3 [23], which compares wellwith the increase of o· 1 1 A in the 1-0 distances in theisoelectronic iodine-containing ions 10 3 - (1·81 A) and ro:­( 1 ·93 A). Presumably, all the electrons of the xenon outer octet
are involved in the [XeO 6]4- bonding.
Conclusion

Approximately one hundred scientific papers have been pub­lished since the report of the first xenon compound in 1962.
Never before has the picture of the chemistry of an element andthe structural properties of its compounds been so rapidly
developed. The limitations of the chemistry of the noble gases
can now be clearly seen. It is evident that compounds ofhelium, neon, and argon are unlikely to be produced other thanas highly unstable species; even the synthesis of a fluoride ofargon appears unlikely. Although more compounds of xenon
and radon may be prepared, those of xenon will probably
be confined to substituted fluoride molecules (for example,
CF 3XeF 5 and ClXeF 0); to adducts of the fluorides; to Xe (1)complexes related to Xe + [PtF 6]-; and to oxygen-contammg
salts. Radon may well form a chloride and even a more stable
oxide than does xenon.
References[1] Bartlett, N. Proc. chem. Soc., Land., 218, 1962.
[2] Claassen, H. H., Selig, H., and Malm,J.G. J. Amer. chem. Soc.,
84, 3593, i962.
[3] Chernick, C. L. et al. Science, 138, 136, 1962.
[4] Villard, P. C.R. Acad. Sci., Paris, 123, 377, 1896.
[5] Travers, M. W. 'Sir William Ramsay', p. 146. Edward Arnold(Publishers) Ltd, London. 1956.
[6] Oddo,G. Gazz. chim. ital., 63, 380, 1933.
[7] von Antropoff,A. Z, angew. Chem., 37, 217, 1924.
[8] Pauling, L. J. Amer. chem. Soc., 55, 1895, 1933.
[9] Ruff, 0. and Menzel, W. anorg. Chem., 213, 206, 1933.
[10] Yost, D. M. and Kaye,A. L. J. Amer. chem. Soc., 55, 3890, 1933.
[u] Bartlett, N. and Lohmann, D. H. Proc. chem. Soc., Land., 115,
1962.
[12] Weinstock, B. Rec. chem. Progr., 23, 23, 1962.
[13] Weeks, I. L., Chernick, C. L., and Matheson, M. S. J. Amer.
chem. Soc., 84, 4612, 1962.
[14] Smith, D. F. J. chem. Phys., 38, 270, 1963.
[15] Dudley, F. B.,Gard,G., and Cady,G. H. Irwrg. Chem., 2, 228,
1963.
Weaver, E. E., Weinstock, B., and Knop, C. P. J. Amer. chem.
Soc., 85, II1, 1963.
Malm,J.G., Sheft, I., and Chernick, C. L. Ibid., 85, 1IO, 1963.
Slivnik, J., et al. Croat. Chem. Acta, 34, 253, 1962.
[16] 'NobleGas Compounds'. H. H. Hyman (editor), University of
Chicago Press, Chicago. 1963.
[17] Fields, P. R., Stein, L., and Zirin, M. J. Amer. chem. Soc., 84,4164, 1962.
[18] Grosse,A. V., Kirshenbaum,A. D., Streng,A.G., and Streng,
L. V. Science, 139, 1048, 1963.
[19] Levy, H.A. andAgron, P.A. J. Amer. chem. Soc., 85, 241, 1963.
Agron, P. A. et al. Science, 139, 842, 1963.
Burns, J. H.,Agron, P.A., and Levy, H.A. Ibid., 1209, 1963.
Templeton, D. H., Zalk.in,A., Forrester,]. D., and Williamson,
S. M. J. Amer. chem. Soc., 85, 242, 1963.
[20] Burns, J. H. Conference on Noble Gas Compounds, Argonne
National Laboratory. April 1963.
[21] Agron, P.A., Begun,G. M., Levy, H. A., Mason, A.A.,Jones,
C.G., and Smith, D. F. Science, 139, 842, 1963.
Claassen, H. H., Chernick, D. L., and Malm, J. G. J. Amer.

z,

chem. Soc.,[22] Williamson, S. M. and Koch, C. W. Science, 139, w46, 1963.
[23] Templeton, D. H., Zalkin,A., Forrester,]. D., and Williamson,
S. M. J. Amer. chem. Soc., 85, 816, 1963.
[24] Smith, D. F. Ibid., 85, 816, 1963.
[25] Bartlett, N. and Jha, N. K. Unpublished work.
[26] Rundle, R. E. J. Amer. chem. Soc., 85, II2, 1963.
[27] Pimentel,G. C. J. chem. Phys., 19, 446, 1951.
[28] Hamilton, vV. C., lbers, J. A., and Mackenzie, D. R. Science,
141, 532, I 963.
[29] Gunn, S. R. and Williamson, S. M. Ibid., 140, 177, 1963.
[30] Gillespie, R. J. and Nyholm, R. S. Quart. Rev. chem. Soc., n,339, 1957.

7