Noble-gas compounds
Neil Bartlett
This year marks the tenth anniversary of the preparation of the first true compounds of noble-gas
elements. Our view of the range of attainable compounds and of the bonding is now more clearly
defined since the author first wrote on the subject in this journal eight years ago. Indeed it is now
possible, using a simple bonding model, to account for the observed stabilities of the known
compounds, and realistically appraise the bonding possibilities for compounds yet unknown.
The range of known noble-gas compounds

Over the past ten years, much work has been devoted toextending the range of noble-gas chemistry, the existenceof noble-gas compounds having been conclusively
established by the author and other workers in 1962. The
historical developments leading up to t�is discovery were
described by the author in an earlier article [1 ]. How­ever, the range of compounds that have so far been
produced remains limited, as Table 1 shows [2]. The con­straints may be readily discerned from an appraisal of
the known compounds, which have two features incommon: (1 ) the noble-gas atom is a heavier (morereadily oxidizable) atom (Rn, Xe, Kr) than the atomsof the groups (ligands) attached to it and (2) the ligands
are always highly electronegative. Both these featuresindicate that the removal of electron density from thenoble-gas atom to the ligands is e:;;sential for effective
bonding.
A sampling of the known chemistry will illustratethese remarks. In keeping with its lower ionization
potential ( ro·6 eV), radon is more readily oxidized thanxenon (I= 12·1 eV). Radon interacts spontaneously
at ordinary temperatures with fluorine or halogen
fluorides (except IF5) to give a fluoride (probably
RnF2), whereas xenon must be thermally or photo­excited in order to interact with fluorine. However, bysuitably controlling the fluorination conditions forxenon, XeF 2 , XeF4 , or XeF 6 may be prepared. Unlikeradon, xenon is not oxidized by CIF3 or BrF3 , although
it is oxidizable by iodine heptafluoride at 200° :Xe +IF 7 ➔ XeF 2 +IF5
The oxidation of krypton (I= 14·0 eV) is much moredemanding, and high-energy, low-temperature methods
are used to prepare the difluoride. Krypton difluoride isthermodynamically unstable (in contrast to XeF 2) andsince L'1Ha. t(KrF2 (g) ➔Kr(g)+2F(g)) = +23 kcal mole-1,
whereas L'1Ha.t(F2 (g) ➔2F(g)) = + 37 kcal mole-1, KrF 2
is a better fluorine atom source than molecular fluorine.
All efforts to prepare a fluoride of argon (I= 1y8 eV)
have failed so far.
Since chlorine is an electronegative ligand we might
expect that chlorides of radon and xenon would be
sufficiently well bound to be stable at ordinary tempera­
tures. Such has not proved to be the case (at least forxenon). However, microscopic quantities of the dichlo­ride have been prepared by electric discharge of xenonN. Bartlett, Ph.D., D.Sc., F.R.I.C.
Was born In 1932 and received his training in chemistry at King's College
Newcastle-upon-Tyne. In 1966 he was ap.pointed a Professor of Chemistryat Princeton University where he stayed until 1969, when he accepted his
present appointments as Professor of Chemistry -in the University ofCalifornia at Berkeley and Principal Investigator in the Inorganic Materials
Research Division of the Lawrence· B\lrkeley Laboratory. His prime re­search interest is In Inorganic fluorine chemistry and he was the first to
prepare compounds of the noble gases.

chlorine mixtures, followed by rapid entrapment in inert
matrices at 20 °K. The dichloride, dibromide, and tetra­
chloride have also been detected through Mossbauer
studies in which the xenon halides were derived from thef3-decay of their 129I relatives thus: 129IBr2 - �XeBr2 •These halides decompose at temperatures well below- roo ° C and have not been obtained in macroscopic
quantities.
Since the oxygen atom is highly electronegative andis comparable in size to that of fluorine, noble-gas
oxyfluorides and oxides were to be expected. Oxy­
fluorides XeOF4 , XeO2F 2 , XeO2F4 , and XeO3F 2 havebeen prepared, directly or indirectly, from XeF 6• Like
the oxyfluorides, the only oxides to be prepared arethose of xenon, and these are limited to the trioxide
and tetroxide. Both oxides are endothermic and are
dangerously explosive. They have been derived-the
tetroxide indirectly-from the thermodynamically stable
fluorides. It is of interest that XeO and XeO 2 have notbeen prepared as laboratory chemicals (an XeO species
has been detected in xenon/oxygen discharges) although
the fluoro-relatives XeF 2 and XeF4 are both readily
preparable.
A noteworthy feature of the xenon oxide system is thedisproportionation behaviour of the lower oxidation
states. Hydrolysis of xenon tetrafluoride does not yielda dioxide, but proceeds according to the equation
3XeF4 +6H2O-XeO3 +2Xe + 1·50 2 + 12HF
Furthermore, certain alkaline solutions of Xe(v1) dis­
proportionate to yield perxenate:
HXeO4-+2OH- ➔XeO6 - +Xe +0 2 +2H 2O
4although the perxenate is prepared in better yield byTable IA selection of noble-gas compounds* to illustrate known oxidation
states and ligands [2]Noble-gas andoxidation state

Ligands

Kr +2
Xe +2

Fluorides
KrF2
XeF2

+4
+6

XeF,
XeF,

+8
Rn +2(?)

RnF2

Oxyfluorides Oxides

XeOF 4
XeO 2 F2
XeO 2 F,XeO 3F2

XeO 3

Other:j:FXeORtXe(OR) 2
F,XeOR'
F,XeOR'XeO,* Available in macroscopic quantities.
t The molecular structure of FXeOSO 2 F is illustrated in Figure 1.
:j: -OR includes -OTeF,. -OSO 2 F, -OClO 3 , O,CCF3 andoso.cF•. -OR'= -oso.F.
107

Table II

Atomic radius and first ionization potential for each noble gas [2]He
Noble gas
1·3
Radius (A)First ionization
potential (eV) 24•586

Figure 1 The molecular structure of FXE}OS0 2 F [2].

passing ozone through a dilute solution of Xe(v1) inbase. Perxenate is the source for the tetroxide:
Ba 2XeOa +2H 2SO4 ➔2BaSO4 +XeO4
Although XeO4 is known and well characterized, allattempts to synthesize the octafluoride have failed.
There is some evidence [3], however, for the Xe(vm)oxyfluorides, XeO3F 2 and XeO2F4•
Since groups like - OSO2F and - OC1O3 are known
to be highly electronegative, it was anticipated that sub­stitution for fluorine could be effected by interactions
such asXeFx +yHOSO 2 F ➔XeFx_y(OSO 2F) y +yHF
Preparations of this kind have been indicated for all
three binary xenon fluorides and the xenon(n) com­pounds, particularly the compound FXeOSO 2F, shownin figure I, have been well characterized. (It should be
noted that the F-Xe-O arrangement in FXeOSO2F
is linear; similar geometry is anticipated for all Xe(OR) 2
compounds.) All of these XeFx_y(OR) y compounds arethermally less stable than the corresponding fluorides.
The nature of the bonding In noble-gas compounds

Since, as we have seen, only the heavier, more readily
ionizable, gases form compounds and only the mostelectronegative atoms or groups are satisfactory ligands
for the noble gases, appreciable bond polarity is to be
expected in noble-gas compounds: Xe8 +-Fx8 - for ex­
ample. We shall illustrate the several theoretical models
which readily allow for this using the XeF 2 case.
The simplest model to visualize is that proposed by J.
Bilham and J. W. Linnett [3]. They suggest that eachfluorine atom acquires a share in a xenon valence-shellelectron:each Xe-F bond being a single electron bond. In thisrepresentation, each fluorine ligand tends towards theelectron configuration of neon; the xenon atom, how108

Ne
1·6

Ar
1-92

Kr
1-98

Xe
2·18

Rn

21,56315•75913·99912·12910•747

ever, retains a (somewhat imperfect) octet configuration.
The acceptability of these postulates is supported byionization potential data, given in Table 11, showing thexenon electron configuration to be less stable than theneon configuration. Neglect of this long known infor­
mation was a primary reason for the slow progress ofheavy-noble-gas chemistry. If the single electron ineach Xe-F internuclear region is equally shared between
the two atoms, then the net charge distribution in themolecule becomes -½F-Xe+1-F-½.
An identical charge distribution is predicted by thevalence-bond representation advocated by C. A. Coulson[4] and by the three-centre molecular orbital approachsuggested independently by G. C. Pimentel [5] andR. E. Rundle [6]. The former model represents XeF 2
as a resonance hybrid of the canonical forms [(Xe-F)+F-J
and [F-(Xe-F)+]. (Note that (Xe-F)+ and F- are both'classical' octet species). If the positive charge is localizedon the xenon atom, the net charge distribution is again-½F-Xe+i-F-t.
The simple molecular orbital model, for XeF 2 , repre­
sents the bonding in terms of three-centre molecularorbitals, derived from a p-orbital from each of theparticipating atoms. Combination of the three atomic
orbitals provides three molecular orbitals; one bonding( + F - , - Xe +, + F - ) ; one non-bonding ( +F - ,- F +) ; and one anti-bonding ( + F - , +Xe - , + F - ) .
The three-atom system provides four electrons topopulate these orbitals. The lowest-energy orbital, whichis the bonding orbital, receives the full quota of twoelectrons, as does the non-bonding orbital which is ofintermediate energy. The highest-energy molecularorbital is the anti-bonding orbital and it remains empty;
consequently there is net bonding. Since the pair ofelectrons in the non-bonding orbital resides largely on
the fluorine ligands, the charge distribution in themolecule is effectively -½F-Xe+l-F-½, the bonding­orbital electron pair being assigned to the internuclear
regions. The molecular orbital model, therefore, corre­
sponds closely to the Bilham and Linnett model. All
three models are, in effect, equivalent, but each view­
point has unique advantages. The Coulson valence-bondapproach is particularly convenient for our discussion.
A wealth of physical data, some of which are presented
Table Ill

Mossbauer data for the xenon halides.* [7]Halide

XeF,
XeF 2
XeCI,
XeCl 2
Xe8r2

Splitting(mm/sec)e•qQ..c(MHz)

Sp Electrontransfer

Electrontransfer
per bond

41·04±0·07
39·0±0·1
25·6±0·1
28·2±0·1
22·2±0•4

2620
2490
1640
1800
1415

3·00
1•43
1-88
1·03
0·81

0•75
0•72
0•47
0,52
0•41

• The bonding model assumes that the only Xe orbitals partici­
pating In the bonding are the Xe 5p orbitals.

in Table m, support appreciable bond polarity Xe8 +-U­
in XeF 2 and the other xenon compounds. For the XeF2
case, a polarity of -½F-Xe +1 -F-½ is consistent with all
the experimental findings, but it is particularly impres­sive that the enthalpy of sublimation calculated by
S. A. Rice and his co-workers [8] in 1963, assuming thatpoint charge distribution, is 13·3 kcal mole-1, whereas
the experimental value, reported [9] in 1968, is 13·2
kcal mole-1•
It will be recognized that if a fluoride ion is removed
from the XeF2 molecule:
XeF2➔XeF + +Fthe resulting cation will be an electron-pair-bond species(: te: R: ) +. If, therefore, the single-electron bond repre­
sentations for the XeF2 molecule are valid, XeF+ salts
should exhibit a shorter Xe-F bond than in XeF2
itself. Vibrational spectroscopic and crystal structure
findings, given in Table 1v, show that the removal of afluoride ion from the XeF 2 molecule does indeed furnish
a bond which is shorter and stronger than in XeF2•
Furthermore, if we accept that the bond length of 2 ·o I A
reported for XeF 2 represents a bond order of 0·5 (that is,a one-electron bond), then by employing an empiricalequation, given by Linus Pauling [IO] which relates
fractional bond order to bond length, we can predict
the bond length for bond order unity. The predicted

length for Xe-F+ of 1·83 A is very similar to the shorter
Xe-F distance in FXeSb2 F11 , which may, therefore, be
formulated as the salt [FXe] + [Sb2F11J-.
The single-electron bond model applies to all of thenoble-gas halides. In valence-bond terms, the tetra­
halides are representable as the resonance hybrid of thecanonical forms (NL2)2 +(L-h, for example, (XeF 2)2 +(F-) 2• Similarly, xenon hexafluoride is the resonancehybrid of (XeF3)3+(F-)a canonical forms. Note that theion species represented in the canonical forms arealways classical octet species (that is, each atom possesses
a valence-electron octet). The increase in net positive
charge on the xenon atom, as its coordination number
grows, accounts adequately for the decrease in bondlength and increase in stretching force constant (given
in parentheses, in mdyn A-1 or 10-16N/m), in the seriesof xenon fluorides:
XeF2, 2·01A (2·6); XeF4, 1·95A (3·0); XeF6 , 1•89A (3·3)
Comparison of the Xe-O bonds with Xe-F bonds
reveals that the oxygen linkage is much stronger than
the fluorine linkage. Thus the bond length and theforce constant for the Xe-O bond in XeO3 are 1·76 A
and 5·66 mdyn A-1, respectively; whereas for XeF6 thecorresponding values are 1·89 A and 3·3 mdyn A-1. In
the molecule XeOF4, the bond lengths and force con­stants are, for XeO, 1•70 A and 6·6 mdyn A-1 and forTable IV

Bond length and bond order in XeF2 and its complexes with F--acceptors
Possible contributing
Canonical forms

Molecule

2·010(6)F�Xe-;r-F"'-180°./"F

2·14(3}"�

/Xe
1·90(3) /FED\ 11ac2>

Xe)

Bond length(A)

Bond order*

2·01

O·St

497

FXe) (Xe(i) +/
\+
(F
F)

(F

\+

(ii)

'f'F

Xe) Xe)
+/

(F
(iii)

F)
+/

(Xe

F-

1·90
2·14

588

(Xe
\+

FF)
(i)}>(ii)=(iii)
1-882(16) 2·186(17)F-�xe�F'178(2) 0 138(2)�
RuF,(F-Xe)+(RuF.)-; plus possibly,
F-(Xe-F)+.RuF,

1-88
2·2

2·35
1-84
F�Xe J°"CF~180 ° 147�
Sb,F, 0(F-Xe)+(Sb2F 11 )-; plus possibly,
F-(Xe-F)+, Sb 2 F,o

1 •84
2·35

XeF+

1·83

0·8,
0·2

604

621

1·0

* From the Pauling equation r(n/2)=r(1)-0·60 log n/2. [10]Assumed value.

t

109

Xe-F 1·90 A and 3·0 mdyn A-1, respectively. It isappropriate to represent the Xe-O linkage as an electron�
pair bond. On this basis the molecule XeO4 is represented
simply as

••
•o •
•
•• •••••: 0: Xe• 0 8....•••:••·.o.

This implies an appreciable pos1t1ve charge on the
xenon atom, since equal sharing of a xenon electron pairwith an oxygen ligand gives a net charge of -1 to thatligand. Evidence for greater polarity of the Xe-O
compared with the Xe-F linkage has been obtained from
X-ray photoelectron spectroscopy. Analysis of thechemical shift data for molecular XeF2 , XeF4, XeF6 ,and XeOF4 indicates that the oxygen ligand in XeOF4
withdraws approximately twice the electron densityfrom the xenon atom than does each of the fluorine
ligands [11]. The findings are consistent with the charge
distribution Xe +3Q-1(F-i)4 which the simple bonding
model suggests.
The simple bonding model also permits us to assess
the stabilities of compounds in a semiquantitative way.
Energy considerations

The marked decrease in ionization potential with in­crease in atomic number in the helium group is itself
suggestive of chemical activity for the heavier elements.
We can appreciate that an electronegative ligand, L,interacting with a noble-gas atom, N, is likely to remainbound if the energy required for electron removal from
N is compensated for by the electron acquisition by L
and the electrostatic attraction energy of N8 + and L6 -.
In order that we may make energy estimates using
readily available thermochemical data it is convenient
to use the cycle shown in Table v. We see that for the
XeF2 case, the ionization energy for Xe + formation(280 kcal mole-1) is appreciably compensated for by theattraction energy between Xe + and F-, (-166 kcal
mole-I). This is what one expects, since the electron is
not removed far from the xenon atom when it becomes
associated with the adjacent, small, fluorine atom. The
attraction energy and the electron affinity of atomic
fluorine together (-246 kcal mole-1) come close tomeeting the requirement; the contributions of anelectron-pair bond energy for Xe-F + (-48 kcal mole-1 )and resonance energy for the canonical forms (Xe-F)+F­
and F-(Xe-F) + generate a stable molecule for which
t.Hat(XeF 2 (g) ➔Xe(g) + 2F(g)) =65 kcal mole-I. A simi­
lar evaluation for KrF 2 immediately reveals that thegreater ionization potential of the noble-gas atom is notrecompensed by an increased attraction energy. Assum­
ing the same resonance energy as in the XeF 2 case, theheat of atomization of KrF 2 is determined to be +33
kcal mole-1, the experimental value being + 23 kcal
mole-1• By similar reasoning we can readily appreciate
why ArF 2 has proved to be so difficult to prepare.
It is also easy to understand why the other halides
should be so much less stable than the fluorides. The
atoms of chlorine and the other halogens are muchbigger than the fluorine atom. The attraction energy,
for the (LN)+L- ion pair, is therefore significantly less(being estimated at 138 kcal mole-I for (Cl-Xe)+Cl-).
Also, the electron pair energies for XeCl + and the otherspecies must be less than for XeF + since fluoride bondenergies are well known [10] to be stronger than thoseof the other halides (for example, I-F =67 where
I-Cl=58 kcal mole-1). We can also ascribe the low
thermal stability of the fluorosulphates and perchlorates
to the large ligand size (the electron affinities of -OSO2 F
and -OC1O3 are much more favourable than for F.)
The case of the xenon oxides is of particular interest,since the tetroxide and trioxide are known, whereas thedioxide and monoxide have been observed only as short­
lived species in discharges. The geometries of XeO4 andXeO3 being known, the electrostatic point charge

Table V

Estimation ofthe heat ofatomization ofa noble-gas dihalide, NL 2 (values in kcal mo/e-1)_;;,(L•N L· ion pair(N-L)'r--------�
l>H(electrostatic)l>H(electronI
pair bond)ResonanceN•,?TL -<o>Energy,t L---------N<•>

+

/

"'

2L<•> <c-----L-N-L<,>l>H(atomization)-<--------Experimental Quantities--------+ t.H(electrostatic)*
Moleculet.H(L+ e ➔L-)
t.H(N++ L-�(N-L)+) [13]/(N)

Resonanceenergyassumedconstantt.H(atomization) [2]observedfrom cycle

XeF2
KrF 2
ArF 2

280
312
342

-80
-80
-80

-47
-37
-38

-166
-176
-195

-52
-52
-52

65 assumed
33
23

xec1 2289-83-40(est.)

-138

-52

32

* The t.H(electrostatic) is estimated as the attraction energy (£=-e 2/estimated or observed N-L distance).
The bond stretching force constant for XeCl2 1·3 and that for XeF2 2·6 mdyn A- 1•

t

110

=

65
23
Molecule
not knowntenergies for the Xe x+ (O-)x configuration may be readily
determined. Table VI shows that the ionization energy
is largely compensated for by the electrostatic attraction
term, although the oxygen atom electron affinity also
contributes. To estimate the heat of formation from theatoms it is, however, necessary to know the energy to be
associated with the electron-pair-bond formation, that is,

fluorides, and KrF2 Raman spectra show that thecomplex KrF2 ·2SbF5 is very much like its xenon relativeand is best represented as the salt [KrF] + [Sb 2F11J-. Now
consideration of the thermodynamic cycle (values inkcal mole-I):
3

In the case of XeO4 the observed heat of atomization
requires that this term be - 152 kcal mole-I; the derived
average electron-pair bond energy ( - 38 kcal mole-I)
when applied to the other oxides allows us to determine
the heat of atomization, t:Jlat(XeOx➔Xe +xO). The
value of - 38 kcal mole-I for the Xe:0 electron-pairbond is compatible with the experimental value [ I 3] of-48 kcal mole-I for Xe:F+ . As may be seen from
Table VI, the calculated t:Jlat(XeO3) agrees with theobserved value within experimental error. The dioxide
is predicted to be very weakly bound and the monoxide
unbound to the extent of +33 kcal mole-I.
It is of interest that spectroscopic studies [ 12] indicatethat Xe-0 is bound with respect to an unspecified
singlet oxygen species, the reported dissociation energy
being 8 kcal mole-1. If the singlet species were ID(O)this would imply that XeO should be unbound by 35
kcal mole-I with respect to 3 P(O), thus providing re­markable agreement with the simple calculations. In
any case, the failure to prepare macroscopic quantities
of XeO or XeO2 is accounted for.
Although these simple evaluations ignore manydifficult points, there is comforting agreement between
observed behaviour and theoretical expectation.

-- 8
3
+
Kr+ + F-·KrF

15I

l E (KrF')

Kr+ F-KrF_:S.12shows that the electron affinity of the cation is equal tothe ionization potential of krypton, less the difference inthe bond energies of KrF and KrF, which is unlikely
to exceed 40 kcal mole-I. The electron affinity of KrF+
is, therefore, at least 275 kcal mole-I. The spontaneousoxidation of iodine pentafluoride is in keeping with thisvalue:+

KrF+Sb2F11 - +IF5➔IF6 +Sb2F11- +Kr
The increasing effective positive charge on the Xe atomin the series XeF2 , XeF4 , and XeF6 , suggests that theease of fluoride ion donation should decrease from XeF2
to XeF6 • The difluoride is a superior donor to XeF4,
but, surprisingly, XeF6 is the best donor of the three
fluorides.
The xenon hexafluoride molecule has been extensively
studied but we still do not know its ground-state geo­
metry. In the vapour phase the molecule is evidently
fluxional (a 'jellyfish'-like molecule): whereas electron
diffraction data show that it is much distorted fromoctahedral symmetry [ 14], molecular beam studies [ I 5]
indicate that any dipole moment must be less than 0·03
Debye. Crystalline XeF6 is best represented as the salt
XeF5 +F- [ I 6]. The XeF5 species seen in crystalline
XeF6 is similar in size and shape to that in the salt
XeF6 RuF 6 -, whose structure [17] is illustrated infigure 2. Crystal structure studies give convincing
evidence that the non-bonding xenon (v1) electron pairoccupies the sixth position of the pseudo-octahedral
XeF5 species. It is probable that the tendency of the
non-bonding 'pair' to steric activity in molecular XeF6
is hindered by the ligand crowding and manifests itself+

Fluoride ion donor properties of krypton and xenon
fluorides

The single-electron bond model for the halides indicatesthat donation of a halide ion, thus,
XeFx➔XeFx-I+ +F- or 2XeFx->-Xe 2F2x-I +F­++

should be assisted by formation of one electron-pair bondin the cation. Examples are known for all of the xenon

Table VI
An estimate of the relative stabilities of the xenon oxides (values in kcal mote-')�trost�i:;(' " ��lon cluster)

Xe<+

+

Xe

l
I
+ x O�XeO,
l'>H atomization)(

xO

I! 'f,1,-x I xx - 33

x x Xe: 0 electronpair-bond energy

Molecule:El/➔ X

l:iH(xO-)

l:iH(electrostatic) *Electron-pair-bond energytx(Xe:O)l:iH (atomization)
Cycle
Observed [2]

XeO,
XeO,
Xe0 2
XeO

2520
1507
768
280

-132
-99
-66
-33

-2320
-1339
-634
-176

-152
-114
-76
-38

=obs.
-45
-8
+33

-84
,::,,-50* Estimated as the point charge attraction energy using observed (or where necessary) estimated interatomic distances,t The electron-pair-bond energy for the Xe:O bond was obtained from the cycle for the XeO, case and the unit energy (-38 kcal
mole-') then used for all other cases.
111

reagents. The exceptional oxidizing properties of KrF 2
have already been remarked upon. Even Xe(n) is anunusual oxidizer as the ready generation of the S0 3F
radical indicates:
XeF 2 +2HS0 3F ➔Xet +2HF +2S0 3F::;S20 6F 2
Since the reduction product is a 'relatively inert' gas thenoble-gas compounds behave almost as ligand radical
sources and the noble-gas leaves the reaction system:
N-Lx + Reactant➔Reactant L + Nt
or
2N-Lx +xHA➔xAL +xHL +2Nt
The interaction of XeF 2 with aromatic hydrocarbons[18] promises to be a convenient source of aromatic
fluorocarbons, for example,
XeF 2 +C6H6 ➔Xet +C6 H5F+HF
The noble gas, expelled from the reaction may of course
be collected and recycled.
The greatest scope for new compounds may well lie inthe preparation of NL+ salts. Thus it is possible thatArF+ salts could be prepared (the cation is bound [13])but a major difficulty is to find a counter-anion that willwithstand the enormous electron affinity of the cation.
It is probable that a variety of XeL+ species (with L
ranging from aromatic hydrocarbons to inorganic
radicals like - SF 5) can be prepared. Such species would
be powerful electrophilic reagents and we may expectthat ingenious chemists will respond to the challenge of
their synthesis.

Figure 2 Structural features of FXeFRuF, and XeF,+RuF.- [17].
(Attempts to prepare a XeF,•RuF, complex have failed.)

not only in the fluxional behaviour but in aiding therelease of fluoride ion.
Relationship of noble-gas compounds to compounds ofthe preceding groups

It should be recognized that the noble-gas compounds
are close relatives of compounds of preceding elementsin the Periodic Table [1]. The perxenates (XeO/-) are
isoelectronic with the paraperiodates and the ortho­
tellurates; and Xe0 3 with 10 3-. Furthermore, XeF 2
can be simply correlated with IF3 and TeF4 ; KrF 2 with
BrF3 and SeF4 ; XeF4 with IF 5 and TeF6 ; and XeF6 with
IF 7• As expected, the bonding models, which work so wellfor the noble-gas compounds, account in a very satis­
factory manner for the properties of these relatives also.
Applications and future directions for research

The ability of radon, xenon, and krypton to form com­pounds [2] may well prove to be of value to the nuclear
energy programme. Radon is a health hazard inuranium mines. It is now feasible to scrub the air of such
mines with suitable oxidizers to remove the radon. Xenon
and krypton are gaseous fission products. The krypton
is dangerously radioactive and presents a considerable
handling and storage problem. Compound formation
could provide for convenient separation and storage.
Perhaps the largest potential for noble-gas compounds
lies in their exploitation as clean oxidative or substitutive
112

Acknowledgment
This work was supported by the United States Atomic Energy
Commission, under Contract Number W-7405-eng-48.
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