NOBLE-GAS COMPOUNDS
by Neil Bartlett

The noble gases, long regarded as inert,
turn out to have a chemistry after all.
Furthermore, this new-found reactivity
throws new light on chemical bonding

IN BRIEF: Two years have now passed since
the announcement that the noble gases, so
named because of their reluctance to asso­
ciate themselves with any other elements
in compounds, were really not so noble after
all. A dozen or more compounds of xenon,
krypton, and radon now sit in bottles onlaboratory shelves all around the world,
while scientists by the hundred probe thechemical, physical, and theoretical peculiar­
ities of these newest of molecules.
Oddly enough, the peculiarities are few;
the structures and reactivities of noble-gas
compounds seem, on hindsight, to be per­fectly straightforward and predictable.
Even the classical laws of valence bonding
still stand. Only one compound, xenon hexa­
fluoride, now appears to be a troublemaker;
its structure, as yet undetermined, may
prove whether the valence-bond theory or
the molecular-orbital theory of bonding iscorrect. No other test case has shown as
great divergence in predicted structure be­
tween the two methods.
The unexpected has been assimilated, thepatient still lives.-H. W.M.

■ The discovery in late 1962 that the heavier
noble gases have a chemistry, albeit limited,
was greeted with incredulity by some chem­
ists and with astonishment by the majority.
The reasons for this surprise, and for the con­tinued reluctance in some quarters to accept
the idea that chemical bonding in noble-gas
compounds is similar to that in "normal" com­pounds, are rooted in history and ingrained
tradition.
Throughout that history, the noble elements
have played a central role in the developmentof a theory of chemical bonding. And the sig­
nificant thing, in most minds, about the dis­covery that they can form bonds is the
new light it may throw on bonding.
Although there was almost universal sur­prise at the announcement, there was little
actual disbelief. Such was not the case in1894 when Lord Rayleigh and William Ram­say announced their discovery of a new ele­ment, argon, present in air in a concentrationof almost 1 % .
An element like argon had not been pre­dicted by Mendeleev's periodic table, formu­
lated 25 years earlier, and many saw no place
for it in the table. Even Mendeleev maintained
56

IXTERXATIOXAL SCIENCE AND TECHNOLOGY

that the new gas must be another molecularform of nitrogen, N3 , and predicted that it
would be very stable.
Ultimately, of course, argon was estab­lished as a new element, and as a member ofa new group in the periodic table, with its
relatives helium, neon, krypton, and xenon. As
Ramsay had proposed early in the argon work,
the new group was placed between the highly
oxidizing halogens and the strongly reducing
alkali metals. This location indicated that the
new gases might well be neither oxidizing nor
Feducing, and hence chemically inert. Ray­leigh and Ramsay conducted numerous exper­
iments on argon, endeavoring to produce
compounds of the element, but they were uni­
formly unsuccessful. They thus concluded, "Itis a most astonishingly indifferent body, inas­much as it is unattacked by elements of very
opposite character.... It will be interesting
to see if fluorine is also without action ...."
To follow up this speculation, Ramsay sup­plied argon to his friend Henri Moissan, who
discovered fluorine in 1886, but nothing cameof the experiments (see FLUORINE CHEMISTRY,
November 1962).
The other gases were subjected to a muchless searching chemical investigation than ar-

Shell Number. n

Noble Gas

1

Helium

2

Neon

2

8

Argon

2

8

8

Krypton

2

8

18

8

Xenon

2

8

18

18

8

Radon

2

8

18

32

18

2

3

4

5

Total

2
10
18
36

Fig. 1. The electronic configuration of the ele­ments produces an eight-electron highest energy
shell in all noble gases heavier than helium.

September 1964

6

54

8

86

r

5econd7as

Fig. 2. Historic apparatus-these glass reaction
tubes were used in the first experiments to show
actual chemical bonding in noble gases; glass
diaphragms in drawing were broken to allow
xenon and platinum hexafiuoride vapor to mix,
producing XePtF,. Crystals below are xenon
tetrafiuoride.

gon, and no evidence of chemical reactivity
was revealed by their discoverers. Since kryp­
ton and xenon were so rare, they were theleast investigated of the group.
Anatomy of a misconception

In the midst of this chemical ferment, J. J.
Thompson discovered the electron in 1897, andcorrelated his speculative atomic model withthe periodic table by arranging the electronsin concentric rings of eight. This tied in beau­
tifully with the new group of apparently zero
valent elements.
Mosely's discovery in 1913 that elements
bombarded with cathode rays emitted x rays
provided the tool which helped establish theatomic numbers of the noble gases as 2, 10,18, 36, 54, and 86, and thus their position in
Mendeleev's scheme.
Bohr's atomic theory, postulated the same
year, restricted the energy of the orbiting
electrons to an integral number of quanta, andlimited the maximum number of electrons ina given energy level. In this model, the noble
gases always contained eight electrons in theirhighest occupied energy level.
All the various pieces of physical evidence
seemed to define the concept that the electronic
configuration of the noble gases was the"ideal" toward which an anthropomorphic ele­ment might "strive" when undergoing reac-shared equally between the atoms. Thus, eachtion with another. This concept was formalized
in 1916 into two hypotheses, put forward by
Lewis and Kossel working independently.
Kossel pointed out, for example, that thoseelements that immediately preceded the noble
gases in the periodic table, the halogens, at­tracted electrons strongly and were univalent.
The alkali metals following the noble gases inthe table, on the other hand, gave up electrons
easily and were also univalent. Thus eachgroup seemed to strive to attain an outside
shell of eight electrons, like the noble g·ases,
even if it had to become an electricallycharged ion to do so. A similar situation held
for the groups of divalent elements removed
one additional step in each direction.
According to Kossel, this gave rise to mu­
tual attraction between oppositely charged
ions, and hence to chemical bonding. The
then new technique of x-ray crystallography
confirmed that this theory must be approxi­
mately true, at least for those elements adja­
cent to the noble gases.
Kossel's theory was somewhat less plausible
in nonionizable molecules such as methane,
CH4• Here, Lewis showed another way that an
atom could achieve a noble-gas structure-by
sharing its electrons. He represented the dia­
tomic hydrogen molecule, H2, with each of theelectrons of the two participating atoms

September 1964

INTERNATIONAL SCIENCE AND TECHNOLOGY57

Compound

XeF2
XeF,
XeF.
KrF.
XeOF2
XeOFa
XeOF,
XeOa
Xeo.-•

Formcolorlesscrystal

,,
,,
,,
,,
,,

Melting
Vapor pressure,point, ° C
mm at 25 ° C

140
114
46
60 (dee.)
?
?

3.8

3

29
115
?
?

colorlessliquid -28
29
white
solid
explodesexists as salt of barium (Ba,XeO.)
and sodium (Na,XeO, • 8H2O)

Table 1. Simple known compounds of noble gaseswith fluorine and oxygen.

hydrogen atom was visualized as approaching
the ideal noble-gas configuration of helium.
In methane and ammonia, NH 3 , the carbon
and nitrogen approach the neon configuration,
the hydrogen atoms approach that of helium.
The twin theories were so successful incorrelating the body of cherpical facts thatworking chemists tended to overemphasize
them, and to ignore the numerous cases where
the so-called octet rule did not apply.
The ever-present skeptic

Not everyone was convinced of the noble­
ness of the group, however. For instance,even the discoverers of argon made use ofargon's high solubility in water, by extracting
it from rain water, which proved to contain
twice as much dissolved argon as nitrogen.
The French chemist Villard showed that theelement formed a crystalline hydrate withwater, and later, hydrates of krypton andxenon were also prepared. Structural studies
have since shown that these "compounds" arereally clathrates, in which the noble-gas atoms
are physically caged in holes in frameworks
composed of hydrogen-bonded water mole­cules.
Water is not the only substance to enclath­rate the heavier noble gases either; �-hydro­
quinone and phenol also form clathrates in­volving some other molecules in the cage
besides the noble-gas atoms.Neon and helium
do not form clathrates; presumably the at­
tractive forces between the caged species andthe cage are too weak in such small, poorly
polarizable atoms.The existence of clathrates,
we should note, doesn't involve the valence
electrons in any chemical bonding. The octet
remains intact.
However, there was one class of noble-gas
compounds which grew up in the literature
which did involve valence bonding.These werethe compounds of noble gases, particularly
helium, with metals. In 1924 Manley described
the first such compound, a helide of mercury,produced by interacting low-pressure helium
58

IXTERNATJONAL SCIENCE AND TECHNOLOGYwith mercury in a glow . discharge. Later
workers claimed helides of tungsten, phos­
phorus, iodine, and sulfur.
Unfortunately, none of the work can besaid to have conclusively established the exist­
ence of definite chemical compounds of thegases with other elements. Waller, after un­
successfully attempting to repeat a number ofpreparations in 1960 asserted, "... it may beconcluded that the inert gases are truly inert."
In 1935, Booth and Wilson claimed to haveevidence of compound formation between
boron trifluoride and argon, but Wiberg andKarbe were unable to confirm these findings,
nor to detect any evidence of compound forma­tion between boron trifluoride and either
krypton or xenon.
Although Oddo must be credited as one of
the first to predict reactivity of the heavier
noble gases-he suggested in a letter to Ram­say that krypton and xenon should form hal­ides-he was deterred by Ramsay's pessimism.
Von Antropoff, coming to the same conclusion
as Oddo, even predicted octavalence for theheavier elements in 1924. In 1932 he and his
co-workers thought they had isolated a red
krypton chloride; a year later he showed it
to be a compound formed from nitrous oxideand hydrogen chloride, both impurities in thereaction mix.
Pauling made the most accurate predictionsof all in 1933: From considerations of ionicradii and periodic table relationships, he pre­dicted the formation of xenon hexafluoride,
XeF6 , and krypton hexafluoride, KrF 6 , andthe possible existence of xenon octafluoride,
XeF 8 • He also suggested that xenic acid,
H 4XeO 6, should exist, and that salts such as
silver xenate, AgXeO 6 , should be preparable.
Unfortunately, the experimental difficulties
of working with fluorine in 1933 were great.
For example, when Yost and Kaye took Paul­
ing's cue and studied the xenon-fluoride reac­tion at CalTech in 1933, they had to generate
their own fluorine from a sometimes-balkycell. Their total supply of xenon, "borrowed"
from Allen of Purdue, was 200 cc at less thanhalf an atmosphere of pressure. The reaction
was carried out in quartz apparatus, which(we now know) doomed it to failure. All they
got for their pains was etched quartz, in allprobability a result of interaction with thefluorine.Had their apparatus been sufficiently
cooled, or constructed of fluorite or alumina
instead of quartz, the xenon fluorides would
have been discovered thirty years earlier.
The failure of all attempts to establish con­clusively the existence of any compounds ofthe noble gases in which the elements were
either oxidized or reduced had, by the latethirties, entrenched the belief that the noble­
gas electronic configuration would apparently
withstand all chemical assaults.
The chain of events that was to shatter the
belief in this invulnerability began in 1961.

September 1964

D. H. Lohmann and I were fluorinating plati­
num and platinum compounds at the Univer­sity of British Columbia, working in glass
apparatus. One of the compounds we obtained
was a red solid, which (after three man-yearsof analysis) turned out to have the empiricalformula PtO2 F 6• Depending on the conditionsof preparation, it resembled K [PtF 6 ] or K
[SbF 6 ]; the potassium (K) ions in both com­pounds, of course, are univalent, and theplatinum (Pt) and antimony (Sb) ions penta­
valent. Weinstock, Claassen, and Malm (keepthose names in mind) had reported making
platinum hexafluoride, PtF 6 , a volatile redsolid, four years earlier at Argonne NationalLab, but monovalent ions of molecular oxygen(0 2 +) were unknown-molecular oxygen holds
its electrons so firmly it takes 281 kcal to re­move the first electron from 32 g (1 mole).
The confirming test of the existence ofmonovalent dioxygenyl ions would be if PtF6
would oxidize oxygen itself to yield the com­pound we had found. In brief, it did; PtF 6 andoxygen reacted rapidly, and the ionic formula,
0 2 [PtF6 J-, took the place of the simple em­
pirical formula.
From calculations of the heat cycle involvedin the reaction, it was apparent that platinum
hexafluoride was the most powerful oxidizing
agent ever known; it obviously had such a
powerful electron affinity that it could create
cations never before attained, from com­pounds normally considered stable. I was
turning over in my mind various possibilities
to try, when quite by chance I ran across achart of ionization potential plotted against
atomic number in a book I was checking foranother matter entirely.
Now ionization potential, the energy re­quired to strip away an electron from anatom or molecule in oxidation, rises and falls
through the periodic table (see Fig. 3) withmaxima at the noble gases and minima at thealkali metals. The maxima, however, gradually
decrease as you go up in atomic number.
Then the thought: We had contemplated theoxidation of relatively inert chemical species;
why not the "inert" noble gases?
+

The first ionization potential of xenon was
the same as that of molecular oxygen, about
12.2 electron volts; radon was also possibly
within range, with a first ionization potentialof 10.7 Ev. Although other factors tend toaffect these energy calculations, we decided totry to confirm experimentally the predicted
reaction with xenon.
And then it happened

It turned out to be a simple and dramatic
experiment. A container of deep red platinum
hexafluoride vapor of known pressure wasseparated by a glass diaphragm from a second
flask containing the same volume of xenon at
a greater pressure. When we broke the dia­phragm, the two gases immediately combined
to produce a yellow solid at room temperature;
the quantity of xenon remaining was com­mensurate with a combining ratio of 1:1. The
yellow solid sublimed when heated under vac­uum, and the sublimate yielded xenon whentreated with water. Clearly the xenon wasbound tightly enough for it to persist in avapor state species. The compound was for­mulated as Xe + [PtF 6 J-, the first known noble­
gas compound in which valence electrons ac­
tually entered into the bonding.
We published these results immediately, andClaassen, Selig, and Malm, at Argonne,
quickly confirmed the reaction. They thenwent on to study the interaction of xenon withruthenium hexafluoride, RuF 6• They notedthat roughly three molecules of RuF 6 wereconsumed per atom of xenon and that some
reduction of the ruthenium seemed to occur.
This suggested that RuF 6 was, at least in part,
acting simply as a fluorinating agent, produc­
ing fluorides of xenon.
To check this heretical idea, the group ran aquick, qualitative test in which they mixed
essentially equal parts (by pressure) of xenon
and fluorine in a metal can and heated themfor an hour at 400 ° C, then cooled to -195 ° C.
When the reaction mix was warmed to room
temperature, there was not a sign of thexenon gas! In spite of the fact that they had
hoped for direct reaction, Claassen wrote in

Heliumo��___..______._________'-'----------L---------------__.___
2

10

18

36

Atomic number

54

86

Fig. 3. Chart of ionization potential against atomic number shows regula;r increase
to a noble gas, then drops. Xenon and oxygen have same value.

September 1964

INTERNATIONAL SCIENCE AND TECHNOLOGY59

Fig. 4, Average bond ener­
gies of related compounds
plotted againstgroup inperiodic table predict bondenergiesofhypotheticalnoble-gas compounds. Com­
pounds in parentheses are
not yet known; closed circles
are predicted values, open
circles are experimental.

�
�

>il

40.,...

his laboratory notebook that afternoon, "Go
home somewhat puzzled. Where is the Xe?"
When the experiment was repeated quanti­tatively, they found that about four parts offluorine combined with one part of xenon, tomake colorless crystals of xenon tetrafluoride,
XeF4• Within a few weeks, they and theircollaborators at Argonne had also prepared-the difluoride and the hexafluoride, the difluo­
ride by irradiating a xenon-fluorine mixture
with light from a high-pressure mercury arc,
the hexafluoride by heating xenon with a large
excess of fluorine at 300 ° C.
A number of other xenon compounds havesince been prepared; they are listed in Table
1 with their properties.

dence for an ogygen-containing salt of kryp­
ton derived in this way. They have not yet
determined its formula, however.
Turner and Pimental at Berkeley irradi­
ated a solid mixture of fluorine and krypton
in an argon matrix at 20 ° K with an intenselight flash, and obtained a chemical species
which infrared evidence showed to be krypton
difluoride (KrF2 ). This method gave no evi­
dence for the formation of an argon fluoride
when a solid argon/fluorine mixture was ir­
radiated alone.
So far, no conclusive evidence has beenfound to support the existence of any argon,
neon, or helium halides or oxides.

The rest of the nobility

What about the theoretical aspects of com­pound formation in these elements? From athermodynamic standpoint, noble-gas chemis­
try turns out to be basically the same as otherchemistries. In every reaction, two factors areimportant: the strength of the bonds formedin the product compared with the bonds
broken in the reactants, and the change indisorder (entropy) as you go from the re­actant system to the product system. If thebonds in the product are stronger than thebonds that must be broken in the reactants,
the reaction will tend to proceed.
The change in entropy in reactions forming
noble-gas compounds is almost always nega­tive. Thus, the total energy (enthalpy) of the
product system must be large enough to offset
the loss, caused primarily by a reduction in
the total number of molecules. For example, it
requires 36.6 kcal/mole to split molecular fluo­rine (F 2 ) into two atoms of fluorine (2F).

Judging from the ionization potentials,
radon should be easier to oxidize than xenon,
and thus should have a more extensive chem­
istry. However, all the isotopes of the gas areshort-lived radioactive species, and this causes
considerable experimental difficulties. How­ever, Field, Stein, and Zirin of Argonne havemanaged to establish the existence of a fluo­
ride. Grosse and his co-workers at Temple
University made the first krypton compound,which they claim to be krypton tetrafluoride(KrF4), by passing an electric discharge
through krypton/fluorine mixtures at liquid
nitrogen temperature. A colorless crystalline
material, it is much less stable than its xenon
analog; it decomposes rapidly into its elements
at 60 ° C.
By analogy with XeF4, you would expect
KrF4 to yield kryptates when hydrolyzed, andindeed tne Temple workers have found evi60

INTERNATIONAL SCIENCE AND TECHNOLOGY

Some thermochemical considerations

September 1964

GROUP (Group No. ==G)
VII
Via

IVa

Valence

Va

VIII(XeF 8 )?

G
G-2

G-4
G-6

Valence

GROUP (Group No. =G)
Va
Via
VII

IVa

G

SeF6

G-2

SeF4

BrF5

VIII

G-4
G-6

.s
"C

IVa

Va

GROUP
Via·

VII

VIII

Silicon.
(Si)

Phosphorus(P)

Sulfur(S)

Chlorine
(Cl)

Argon

!Germanium(Ge)

Arsenic(As)

Selenium(Se)

Bromine(Br)· Krypton(Kr)

Tin. (Sn)

Antimony. (Sb)

Tellu..-ium(Te)

Iodine
(I)

Xenon
(Xe)

Lead
(Pb)

BismuthPolonium(Po)

Astatine(At)·Radon(Rn)

I(13.i)

�.-��d=.

'�-�- .-......;;.,.t=.�

(A)

Fig. 5. Plotting just the existence of compounds
in groups ahead of the noble gases predicts theexistence of those that we now know are real.

Thus, 18.3 kcal/mole can be designated the en­ergy of formation of each fluorine atom. Ex­
periments on both xenon tetrafluoride andhexafluoride show the energies per Xe-F bondto be about 32 kcal. This is about 14 kcal morethan the enthalpy of formation of the fluorine(see margin) ; voila, fluorides form easily andare stable.
On the other hand, the average energy as­sociated with each Xe-0 bond in xenon triox­
ide is less than 28 kcal. This is 31 kcal lessenergy than is required to form atomic oxy­
gen from molecular oxygen, and is in accord
with the explosive nature of the trioxide. It
also accounts for the lack of success in ob­
taining the compound from the elements.

Thermochemical calculations can be used to
predict the stability of yet uninvestigatedcompounds. For example, calculations showthat xenon oxide tetrafluoride, XeOF4, should
be an exothermic, stable compound with anenthalpy of -24 kcal/mole. Xenon dioxide di­fluoride, Xe0 2F2, on the other hand, should beendothermic and unstable, with an enthalpy of+35 kcal/mole. In fact, the first has alreadybeen found; the second has not.
Another method of estimating the averagebond energies in unreported compounds is to
extrapolate from values for related com­pounds. Thus, in Fig. 4, we have plotted aver­age bond energies for compounds in the same
period of the periodic table. From this it ap- ·pears that the average bond energy in xenon
octafluoride, XeF 8, should be about 24 kcal/
mo!. This would give a standard heat of for­mation of about -45 kcal/mole, which is cer­tainly exothermic enough to be stable, al­though the entropy for this reaction is highly
unfavorable. A group of Yugoslavian scien­tists headed by J. Slivnik, has already re­ported isolating this compound. Although thereported volatility is more than you would ex­pect for an octafluoride.
The difficulties of preparing KrF4 and its
instability also coincide with the extrapolated
curves. Krypton difluoride may be more stable
than the tetrafluoride. Oxides and chloridesof xenon and krypton should be endothermic,
and hence thermodynamically unstable. In
fact, only the trioxide and tetroxide of xenon
have been isolated.
We must assume that the combination ofelements other than those discussed so far
with the noble gases will be still less favor­
able. Only the more electronegative elements
will form strong bonds, and only those ele­ments with low ·heats of atomization will befavorable. Nitrogen, with its high heat of
atomization, (113 kcal/mole) would almost
certainly render a noble-gas compound endo­
thermic and unstable.
Keeping up with the Joneses

Another approach which can supply infor­
mation about noble-gas chemistry is compar­
ing tl'le chemistry of neighboring atoms withina given period of the periodic table. For ex­
ample, if you chart the existing fluorides ofthe elements in the period containing xenon(tin, antimony, tellurium, iodine, and xenon)
against their respective group numbers, thevalence of the cation turns out to be numeri­
cally equal either to the group number, or to
the group number minus an even number.
Fluorides of xenon are a perfectly normal ex­tension of this tradition, as shown in Fig. 5.
When you do the same thing with the chlo­rides, however, you get quite a different pic­
ture. Chlorine is less able to excite the higher
valences; from the trend of the curve, the ex­istence of any xenon chlorides appears un-

September 1964

INTERNATIONAL SCIENCE AND TECHNOLOGY61

likely. A chart of the fluorides of the elementsin the period containing krypton also indi­cates little likelihood of krypton fluorides
higher than the tetrafluoride.
It is more difficult to attain valences as highas the group number in the radon period, andanalogs of XeF8 and XeO6-4 would be lesslikely. In spite of this, however, several com­pounds appear to be possible, including thefluorides RnF2, RnF4, and RnF6, as well as achloride, perhaps RnCl�, and an oxide, RnO 3•

MOLECULE AND ION SHAPES
Shape

62

INTERNATIONAL SCIENCE AND TECHNOLOGY

RelatedNonb(?nding
Compounds Bonds
pairs

linear2

3

42

square-basepyramid5

1

tetrahedral

4

0

3

1

6

0

squareplanarShapes and more shapes

As far as their molecular shapes are known,
the noble-gas compounds are identical to thoseobserved in isoelectronic ions-those groups
which on becoming ions gain or lose an elec­tron or two and achieve an electronic configu­
ration identical to those postulated for thenoble-gas compounds. For example, the added
electron in the iodine dichloride ion, IC1 2 - ,compensates for iodine's having one less elec­tron than xenon; thus the ion is isoelectronic
with xenon difluoride.
Furthermore, the shapes may all be ration­
alized on the basis of the mutual repulsion ofpairs of valence electrons, a useful rule ofthumb in predicting shapes. One major pointof this rule is that the pairs of electrons in avalency shell of an atom are always arranged
basically in the same way, irrespective ofwhether they are shared with another atom
(in which case they are bonding electrons) or
not shared (nonbonding). Their basic ar­rangement depends only on their number. Sec­
ond, the electrons of a double or triple bondare repelled as a set. Thus, they occupy only asingle position in the arrangement, but since
they embrace multiple electrons they repel
more effectively than a single bond pair. Fi­
nally, the nonbonding electron pairs will becloser to the parent atom than the bonding
electrons, and have a greater repulsive power.
In xenon difluoride, for example, one elec­tron in each Xe-F bond is contributed by the
xenon and one by the fluorine. Thus two xenon
electrons of the valence set of eight are in­volved in bonding, while six are grouped inthree nonbonding pairs. Minimum electron re­pulsion for such a set is achieved by placing
the pairs at the points of a trigonal bipyramid,
as shown in Fig. 6. To minimize repulsion be­
tween nonbonding pairs, the nonbonding elec­
trons take positions 120 ° apart, in an equilat­eral triangle array. The fluorines locatedabove and below this plane each experience
repulsion from the nonbonding electron pairs,
which will lengthen the XeF bonds. This
lengthened bond has been verified. ·The other cases can be derived in a similarway, and it is noteworthy that the predictions
are correct in all known cases. The one big
problem that still exists is xenon hexafluoride,
for which we do not yet know the structure.
The electron-pair repulsion rule predicts a

Configuration

XeO3

X

trigonalpyramid103

octahedral• =xenonO =fluorine•=oxygen,Fig. 6. Shapes of noble-gas compounds comparedwith those of neighboring elements, according toelectron pair repulsion rule.

nonoctahedral structure, since the xenon will
be surrounded by seven electron pairs (as in
iodine heptafluoride, IF 7, which raised a few
questions of its own when it was discoveredin 1927), six bonding and one nonbonding.
XeF 6 is isoelectronic with IF 6·, and there isevidence that indicates that this ion is non­octahedral. The infrared and Raman spectra
which have been obtained for XeF 6 cannot beinterpreted in terms of an octahedral mole­cule. Even the lower volatility of the com­pound compared to IF 7 is against such an in­
terpretation. On balance, then, I believe thatthe available evidence favors nonoctahedral
symmetry; evidently the compound crystal­lizes differently from the other hexafluorides,
which are without exception octahedral.
Theoriesof bonding

It should be obvious by now that the noble
gas compounds must owe their existence tothe same kind of chemical bonding found incompounds of other elements. Before we dis­cuss the electronic bonding theories advanced,
however, let's take a quick look at the basic
concepts of atomic structure. The most signif­
icant difference between our present view and
Bohr's is that electrons cannot be pinpointed,
but have only a probability of being in a par­ticular location at a given time. These proba­
bilities are expressed as wave functions, which

September 1964

are usually characterized by four constants,
the so-called quantum numbers. The principal
quantum number, n, roughly defines the en­ergy of an element, l defines the shape of theorbital described by the probability wave
function, and m 1 defines the orientation of theorbital in space. In addition, s gives the spinof the electron. Electrons having l
0 areoften called s, electrons with l
1 are called
p, while those with l
2 are called d electrons.
For atoms in which n
3, l can equal 1, 2, or
3, and s, p, and d orbitals are available; theirshapes are shown in Fig. 7 and the sketchin the margin.
In the isolated hydrogen atom, these differ­ent orbitals have the same energy. In more
complicated atoms, orbitals of like n but dif­fering l are not equivalent. Indeed in xenon,
where n5, l has the values of 0, 1, 2, 3, and4, but only the n = 5, l = 0 and 1 states areoccupied. The eight electrons so contained con­stitute the outer octet of xenon. The next­
highest energy level available to a promoted
valence electron is not the n5, l
2 orbital
set, but the n = 6, l = 0. This is why cesium,the element following xenon, is an alkali
metal; its valence shell contains only one elec­tron n6, l0, which is readily lost.
All electronic theories of chemical bonding
are based on these atomic orbitals, and all fall
into one or the other of two different ap­
proaches. In the valence-bond method, chemi­
cal bonds are discussed in terms of the over­lap of atomic orbitals. The greater the overlap,
or interpenetration, of the electronic orbital"clouds" associated with the interacting orbi­
tals, the higher the electron density between
the nuclei, and the greater the attraction-thebond. Thus, a hydrogen molecule is formed
from two spherical hydrogen atoms, whose
orbitals overlap as shown in the margin.
In hydrogen, the interacting atoms use pureatomic orbitals; in other cases, however, ex­perience demands that rather different orbi­tals should be used. For example, carbon con­tains one 2s and two 2p orbitals. Overlappingthese orbitals, however, does not give the tet­rahedral symmetry observed experimentallyin carbon compounds. It is therefore necessary
to postulate that the carbon orbitals are
mixed, or hybridized, to give four equivalent
orbitals, called sp3 hybrids. This involves the"promotion" of one s orbital electron.
A hybridization scheme may also be invokedto account for the octahedral molecule of sul­
fur hexafluoride, SF 6 , which involves two
high-energy 3d orbitals to give the required: hybrid set of sp3 d 2 orbitals. Extending thisapproach to the noble gases, xenon difluoride,
as a linear molecule, would use two sp hybridsof the xenon to overlap with p orbitals of thefluorine atoms. The six remaining nonbonding
xenon electrons would then be housed in three"hybridized" xenon orbitals, involving either
the 6s, 4/, or 5d.

=

=

=

=

=

=

=

64

=

=

INTERNATIONAL SCIENCE AND TECHNOLOGY

Fig. 7. Three p orbitals (marginal) are equiva­
lent in energy, except for orientation, as are thefive d orbitals (above). Because of their direc­
tionality, however, they react differently in thepresence of other atoms or molecules.

Simila-rly in the tetrafluoride, which is squareplanar in structure, the valence-bond approach
would allow the use of a high-energy d orbital
to attain an sp 2 d hybrid set of bonding orbitals; a second high-energy orbital and theremaining p orbitals of the valence set wouldhouse the four nonbonding electrons. Xenon
hexafluoride, with six bonding electron pairs
and one non-bonding electron pair, would in­volve three orbitals of higher energy than thevalence set. If, as the electron-pair repulsion
theory predicts, the nonbonding pair is ef­
fected by spatial or steric considerations, thenthe molecule would not be octahedral, and a
hybridization scheme for the xenon orbitals
such as sp 3 d3 or sp 3 f3 would be appropriate.
Many chemists are unhappy with the use ofhigh-energy orbitals in bonding such molecules
as SF 6 , and lately a simple molecular orbital
approach has found support. A molecular or­
bital, or M.O., is the molecular counterpartof an atomic orbital: true M.O.'s belong to theentire molecule. They are generally less pre­cisely defined than atomic orbitals, and infact are generated by a suitable linear combi­
nation of the mathematical expression for theatomic orbitals. The number of molecular or­
bitals is equal to the number of participating
atomic orbitals.
In forming a hydrogen molecule, for ex­
ample, the two ls orbitals give rise to two

September 1964

8·Hzmolecular orbitals, one bonding (lower-energy
configuration) and one antibonding (higher­
energy) as shown in the margin. Since eachmolecular orbital, like each atomic orbital, cancontain only two electrons, the two electronsof the hydrogen molecule reside in the bond­ing orbital. In contrast, the atom of helium
contains two electrons already; thus both thebonding and antibonding orbitals are filled,
and helium remains monatomic-forming an
He2 molecule would offer no advantage inenergy over two separate atoms.
Now a true M.0. extends over all atoms ina molecule; it is multicentered. A bonding
M.O. containing two electrons will therefore
tend to hold all of the atoms in a molecule
together. The bonding stability will, however,depend on the number of such bonding M.O.'s
in the molecule; if there should be only one
bonding orbital in a molecule with n atoms,(and there is no filled antibonding orbital) theinteratomic linkages will be only about 1/n
as strong as an electron-pair bond. In thelinear iodine dichloride ion [Cl-I-Cl] -, physi­
cal evidence shows that essentially only the porbitals of the iodine can be involved in bond­ing. Similar evidence shows that p orbitals arealso responsible for the bonding in the squareplanar [IC1 4] - ion.
Molecular orbital theorists have widely as­
sumed that similar situations will obtain in
the molecules of XeF2 and XeF4, which arepseudo-isoelectronic with the iodine chlorides,
and have used the theory without significant
modification to describe these compounds. In
XeF2 , they assumed that a p orbital of thexenon valence set combines with two fluorine
p orbitals to give rise to one bonding M.O.,
one nonbonding M.O., and one antibonding
M.O. The two XeF bonds therefore are both
dependent on a single bonding electron pair;consequently the bonds should be rather longand ionic. The electron pair repulsion theory
also predicts a long bond. M.O. theory ac­
counts for xenon tetrafluoride in the sameway, with two xenon p orbitals interactingwith pairs of fluorine atom orbitals.
In the hexafluoride, all three xenon p orbi­tals are involved in creating three-center
M.O.'s (two F, one Xe). The p orbitals aremutually at right angles; thus the theory pre­dicts the tetrafluoride to be square planar andthe hexafluoride to be octahedral.
The symmetries observed for the difluoride
and the tetrafluoride confirm those predictedby both theoretical approaches. However, thehexafluoride, still unmeasured because of ex­
perimental difficulties, is indicated to be non­octahedral on the balance of experimental evi­
dence. This would be compatible with valencebond theory and would, of course, imply the
involvement of high-energy xenon orbitals.
Exper,imental verification of the shape is thus
being pressed with great interest.
From the preceding discussion, it should
66

INTERNATIONAL SCIENCE AND TECHNOLOGYeobe obvious that the range of stable compoundsof the noble gases will probably never become � ��extensive. In all probability the fluorides or �
fluorine-containing compounds of kryP.t6ns,
xenon, and radon, and perhaps chlorides andoxides of the latter will be the only thermodynamically stable noble-gas compounds. We�
may be able to prepare kinetically stable oxygen-containing salts of radon and krypton,
however, and as in the case of the xenates andperxenates, these should be powerful oxidizing agents.

ry'0

Where to from here?

It is unlikely that the noble gases wilt form
even metastable compounds with elements lesselectronegative than chlorine, but it may prove
possible to link complex groups of high elec­
tronegativity to xenon atoms already linked
to fluorine. Thus it will be worthwhile expend­
ing some effort to prepare compounds suchas F 5 XeNF2 , F 5 XeCF 3 , arid even F 5 XeSF 6 •There appears to be no reason why xenon
or the other noble. gases should form com­pounds in which the noble gas acts as an elec­tron acceptor, to become a negative ion. True
helides, xenides, and the like, for example,
seem most unlikely. However, in view of theclaims by chemists, respected for their skill
and veracity, for compounds of the gases withmercury and other metals, the whole subject
should be reinvestigated.
Presumably the xenon fluorides will find use
as fluorinating agents in preparing fluoridesof other elements in high-valent states. Suchreactions would be "clean"; the reduction
product of the fluorinating agent will be rela­
tively inert xenon, which would simply bubble
off to be recovered. The explosive oxides may
find an application where oxygen is needed
rapidly in high concentration; decomposition
products would be gases.
Much more certainly, a number of uses will
be found for the easily stored, thermally stablealkali perxenates. They are strong oxidizing
agents; their reduction products are merely
the alkali oxides of hydroxides and xenon.
Radon compounds could be interesting assources of alpha particles. They could be de­posited at any site in a properly designed ap­
paratus, and be built up in a probe to any de­sired level. When necessary, the radon could
be easily removed from the site by cliemicallyreducing the compound.
Undoubtedly the greatest immediate effectof the discovery of simple chemical com­pounds of the noble gases has been its effecton the attitude of the average chemist. It has
removed inhibitions and has given an injec­tion of healthy scepticism to a generation
which had become overly theoretical in its ap­
proach to what is still very much an experi­mental science.
For what little additional information there
is on the subject, see To Dig Deeper, page 103.

September 1964

DO SCIENTISTS
NEED A'COMPREHENSIVEINDEX TOALL OF SCIENCE?
When we asked scientists thisquestion recently, most of them
replied yes - a few said no. If
you are among the few who are
perfectly satisfied with the "selec­
tive" approach of conventional
indexing systems, stop right here.
However, if you, along with the
majority, feel that improved meth­ods and comprehensiveness areneeded in information retrieval anddissemination, read on.

Every research scientist has, on
one occasion or another, been
stumped by the complex problemsof searching the scientific litera­ture. Today's search requirements
are hobbled by the limitations of
traditionally organized indexing
systems.
At the Institute for Scientific Infor­
mation, we have tested and devel­
oped a new dimension in indexing.
To this we have added compre­
hensive coverage, resulting in aunique tool for information dis­covery. We call it Science Citation
Index.

The Science Citation Index instantly
identifies the most recent publi­
cations referring to a particular
work since its publication. The
Science Citation Index provides thescientist with an improved, novel
starting point: the specific paper orbook of a specific author • • • andfrom there the searcher is brought
forward in time to current papers
relating to the earlier work.
Revolutionary? Invaluable? We
think so and we think you will too
after you've had an opportunity toreview descriptive material on thisnew technique. Write us now and
we'll send details by return mail-without obligation, of course.
Dept. 27-2

INSTITUTE FOR SCIENTIFIC INFORMATION
325 Chestnut Street

106Phi/,adelphia Pa 19106

INTERNATIONAL SCIENCE AND TECHNOLOGY

found in "Magnetohydrodynamics-Future Power Process?"by Sporn and Kantrowitz, Power, Nov. 1959. To gage theenthusiasm of the Russians for MHD power generation, see
the article "Power Stations of the Future" in Svetskaya
Litva, Mar. 1961. The possible coupling of nuclear heatsources to MHD generators is discussed in "MHD Power
Generation" by Gundson, Smith, Tsu, and Wright, Nu­cleonics 21, 7 (1963). A more detailed account of MHD
work going on in the authors' laboratories is presented in
Brogan's "MHD Power Generation," IEEE Spectrum, Feb.
1964. Readers interested in the process of nonequilibrium
ionization should scan Rosa's "Nonequilibrium Ionization in
MHD Generators," Proc. IEEE 51, 5 (1963).
The day-to-day literature of the field is heavy going, but
recent papers may be found in Proc. of the 5th Symposiumon the Engineering Aspects of Magnetohydrodynamics(April 1964 at MIT); write Prof. W. Heiser, Dept. of
Mechanical Engineering, MIT, Cambridge 39, Mass. The
world-wide interest in this field is well illustrated by papers
given at the recent (July 6-11, 1964) International Sympo­
sium on Magnetohydrodynamic Electrical Power Generationin Paris. For a booklet of lengthy abstracts and the 20-page
summary of this meeting prepared for ENEA by Satkowski
of ONR write to H. B. Smets, Research Office, European
Nuclear Energy Agency, OECD, 38 Blvd. Suchet, Paris 16.
There is, as yet, no single book on MHD power; however,two useful volumes are Cambel's Plasma Physics and Mag­
netoff,uid-mechanics (McGraw-Hill, 1963, $11.50) and Spitz­
er's Physics of Fully Ionized Gases (Interscience, 1956, $5).
NOBLE-GAS COMPOUNDS

56

The literature of the reactions of these elements is nat­urally sparse, except in recent review articles, with one
notable exception. A book called Noble-Gas Compounds,edited by Hyman, reproduces a symposium held at the Ar­
gonne National Labs about one year after the initial dis­coveries; and it's good (U. of Chicago Press, 1963, $12.50).
A shorter review is by Chernick, "Chemical Compoundsof the Noble Gases" (Record of Chemical Progress 24, No. 3,1963). If you'd like to read about some frustrations ac­
companying the discovery and characterization of the noble
gases, try A Life of Sir William Ramsay by Travers (Ed­ward Arnold Publ., Ltd., London, 1956, $10). It's complete
and recent. Finally, a two-volume treatise on the gases them­selves has recently come out, edited by Cook (Interscience,
2 vols., 1961, $17.50 each).
PICTURE CREDITS
6-Top: Jacques G. Richardson; cen­ter: Robert Lyons
7-Howard Mattson; Joe LaCava
19-Hoffrel Instruments
20-Clinical Center, NIH; Sam Adams
II, Eastman Dental Dispensary
21-Top: Robert Phillips; bottom left:
Sam Adams II; bottom right: Clini­
cal Center, NIH
22-Top & center: Hirotake Kakehi;bottom: V. Parsonnet, G. Myers
23-ITT Federal Laboratories
24-Left: NIH; center: Sean Mullan,
U. of Chicago; right: Realites­
Opera Mundi; bottom: Antoine
Remond25-Left: Borg-Warner; right: NIH,
Inst. R&D; bottom: Sperry Prod­
ucts-Grace New Haven Hospital
26-IBM-Tulane; NIH
28-National Institutes of Health29-Top: ACM Comm., U of Pa.; bottom left: NIH; bottom right: MIT

September 1964

31-Left & center: General Electric;
right & bottom: Weizmann Inst.
32-Wellcome Museum
33-40-Ron Appelbe, London
42-Douglas Aircraft
43-44-M-B Electronics Div., Textron
Electronics
48-Douglas Aircraft; Electrical Test•
ing Laboratories
49-Top: The Martin Co.; bottom:
Wyle Laboratories; right: Associ­ated Engineering Test Laboratories
SO-Tenney Engineering
52-Electrical Test Engineering
57-Left: Howard Mattson; Argonne
National Laboratory
70-77-Drawings by Rowland Wilson
Avco-Everett Research
84-Bottom:
Laboratory
85-Avco-Everett Research Laboratory
89-Left: International Research & De­
velopment Co., Ltd., Newcastle/
Tyne
99-Heka; Chase Ltd., Photo