658

CHEMISTRY IN BRITAIN JULY \987

Twenty-five years of noble gas
chemistry
John H. Holloway

This year we celebrate the 25th anniversary of the discovery of noble gas
chemistry. The discovery toppled a belief that had grown for almost 70
years-that no member of the family of 'inert gases' could form athermodynamically stable chemical compound. Now, a quarter of acentury on, it is appropriate to look back at the origins, development andmore noteworthy recent advances in the field.
In March 1962 Bartlett and Lohmann
reported the preparation of O2+[PtF6r.1
The production of O 2+, by simply mixing
oxygen with an equimolar quantity ofplatinum hexafluoride at room temperature, showed PtF6 to be 'an oxidizer of
unprecedented power',2 capable of oxidising molecular oxygen. Because thefirst ionisation potential of molecularoxygen (12.2 eV) is slightly more than thatof xenon (12.13eV) Bartlett thought thatxenon gas might be similarly oxidised.
Within a month he had confirmed his
suspicion and, in June he published areport of his observation of the spontaneous and rapid oxidation of xenon by
the deep red PtF6 vapour to give a solid,
yellow-orange, quinquevalent platinumcomplex. 3
The papers outlining the discoveries
appeared in the Proceedings of the
Chemical Society,1.3 the predecessor ofChemistry in Britain. The second of theseshort papers, which was less than 250
words in length (Fig. 1), initiated what
was to become a flood of activity inan entirely new area of chemistry.
Before the end of August 1962, Howard
Claassen, Henry Selig and John Maim
had prepared the first binary fluoride of anoble gas, XeF4,4 and by the end of theyear, Rudolf Hoppe and his group hadannounced the synthesis of XeF 2,5 Jozef
Slivnik and his coworkers the preparationof XeF6,6 and Paul FieJds, Lawrence Steinand Moshe Zirin the first evidence of theexistence of a radon fluorideJ By early
1963* the existence of compounds ofkrypton, xenon and radon was firmly
established. 11Time has confirmed the validity ofthese early observations. Stable compounds are formed only with the heavier
gases, and the only stable compounds

Fig. 1 . Bartlett's short paper announcing the discovery of Xe1PtFef and carrying his
signature. Reprinted from Proceedings of the Chemical Society, June 1962, p 218.
Xenon Hexaflaoroplatiate(v) Xe~.PtF.J­
By NEIL BARTLETT(DEPARTMENT OF CHEMISTRY, THE UNIVERSITY OF BRITISH COLUMBIA,
VANCOUVER 8, B.C., CANADA)

A RECENT Communication l described the compound
dioXYlCllyl hexaftuoroplatinate(v), Ot .. PtF.-, which
is formed when molecular Ol\;ygen is ol\;idised byplatinum hexaftuoridc vapour. Since the first i~nisa­tion potential of molecular ol\;y~n,2 I L! ev, is comparable with that of xenon,z 12·\3 ev, it appearedthat xenon might also be ol\;idised by the hexafluoride.
Tensimetric: titration of xenon (AI RCO .. Reagent
Grade") with platinum hexaftuoride has proved theexistenc:e of a I: I compound, XePtF•. This is anoran~yeUow solid, which is insoluble in carbon
tetrachloride, and has a nqiiaible vapour pressure atroom temperature. It sublimes in a vacuum whenheated and the sublimate, when treated with water
vapour, rapidly hydrolyses, xenon and oxygen being
evolved and hydrated platinum dioxide deposited:
lXePtF. + 6H.O -+ lX. + O. + lPtO. + UHF

The composition of the evolved gas was establishedby mass-spectrometric analysis.
Although inert-gas c1athrates have been described,
this compound is believed to be the first xenon
charge-transfer compound which is stable at room
temperatures. Lallice-encrgy calculations for the
xenon compound, by means of Kapustinskii's equation,3 give a value -- 110 kcal. mole-I, which is only
10 kcal. Olole- I smaller than that calculated for the
dioxygcnyl compound. These values indicate that if
the compounds are ionic the electron affinity of theplatinum hcxaftuoride must have a minimum valueof 170 kcal. mole-I.
The author thanks Dr. David Frost for mass
spectrometric analyses and the National ResearchCouncil, Ottawa, and the Research Corporation for
financial support.
(R,cei",d, May 4th, 1962.)

Butlett and Lobmann, Proc. CM"'. 5«.. 1962, 115.
• Field and Franldin, "Electron Impact Phenomena," Academic Prest, Inc., New York, 1957, pp. 114-116.
• Kapustinsltii, QIllUI. RI .... , 1956, .8, 284.
I..

that can be made by direct synthesis are
the fluorides. Most compounds that havebeen obtained consist of noble gas atoms
bonded to the most electronegative
elements, fluorine and oxygen, but
species bonded to chlorine, nitrogen andcarbon have also been reported. Xenon
chemistry is the most extensive-Xe isknown in oxidation states ranging from
+2 to +8, while krypton and radon species
are known only in oxidation state +2.
Stable compounds of helium, neon andargon are not known, but experimental
estimates of the bond strength in the[ArFt cation 12 suggest that continued
efforts to lock this into an adduct areworthwhile.

Xenon fluorides 11 •13
Apart from XeF, which has been obtained
as an unstable free radical, there is noevidence for the existence of fluorides in
odd-numbered oxidation states. The
fluorides, XeF 2, XeF4 and XeF6 can beobtained by reacting xenon directly withfluorine, but early reports of XeFa havenot been confirmed.
• The difluoride is prepared by heating
an excess of xenon with fluorine or byirradiating mixtures of the gases withultraviolet light or sunlight.
• The tetrafluoride is prepared by heating xenon and fluorine under pressure ina ratio of 1 :5 by volume.
• The hexafluoride is prepared by heating xenon and fluorine under pressure ina ratio of 1 :20.
All the xenon fluorides are colourlesssolids except the hexafluoride, which isyellow in the liquid and gaseous states.
The di- and tetrafluorides have low
volatilities, and linear and square planar
structures in the solid and vapour states.
Of the 16 known hexafluorides, xenon
hexafluoride has the highest boilingpoint. Th~s suggests that, unlike theother hexafluorides, this compoundmight be polymeric in the solid state.
However, uncertainty about its struc. ture remains, mainly because it is readilyhydrolysed to the explosive Xe03 andits strong fluorinating ability damages
most container and window materials,
making it difficult to study. Unlike otherhexafluorides it is not octahedral in thevapour but is probably a non-rigid (floppy),
distorted octahedron. 14 When molten or• Early in 1963 it emerged that three other groups.8-10
working independently. had all prepared XeF, atabout the same time .

660

CHEMISTRY IN BRITAIN JULY 1987

~)

2

~)

Fig. 2. The structure of liquid XeFs.

i

3
[111]

solution and powerfully oxidising, converting for example 2Cr to C1 2, Cr(lIlkto
Cr(VI) and even [Br03r to [Br04r.17 The
tetra- and hexafluorides are progressively
more powerful oxidisers and difficult tocontrol. Tt"\eir instant hydrolysis by water,
however, has provided interesting oxides,
oxide fluorides and powerfully oxidising
xenate and perxenate salts .
• Hydrolysis reactions. The difluoride,
XeF 2, is stable in acid or neutral solution
but decomposes immediately in base,liberating xenon, oxygen and HF. The
tetrafluoride is instantly hydrolysed and,depending upon conditions, up to a thirdof the xenon may be retained in solution(rather than the half expected in a simple
disproportionation, 2Xe(lv) - Xe(vl) +Xe(II)), and XeF 2 may also be liberated. It
has been suggested 20 that the initial
product may be XeOF 2, which decomposes to give XeF 2 if the water supply
is limited. When larger amounts of water
are present Xe02 and XeO may beformed, which can then give Xe03'
The more vigorous reaction of XeF 6
with water can be controlled by passing
dry nitrogen over crystalline XeF6 tosweep the vapour into water where thefollowing reaction occurs:
XeF6 + 3H 20 - Xe03 + 6HF

Fig. 3. The crystal structure of the cubic phase of XeFe (-80 0 C). One unit cell
contains 144 'XeFs' units, 24 tetramers and eight hexamers. 1s
in solution the mono mer is in equilibriumwith tetrameric rings consisting of [XeFsr
and F ions, but at low temperatures themonomer disappears. The fluorines inthe ([XeF srF)4 tetramers take part in acomplex scrambling mechanism in which
each of the four units is connected as a
cogwheel to its neighbours (Fig. 2).'S In
the solid, although at least four forms
exist, only the cubic structure-comprising [XeFsr and F ions grouped as
tetramers and hexamers within the unit
cell-has been described in detail (Fig.
3).'S The fact that both hexamers andtetramers occur suggests that the way inwhich the [XeFsr ion is bridged is notcritical. Furthermore the bridging F ions
are not close to the four-fold axis of the[XeFsr groups. This suggests that in keeping with the Gillespie and Nyholm valenceshell electron-pair repulsion theory,
which has provided the most accurate
and simple rationalisation of the shapesof noble gas compounds, a lone pair ofelectrons may occupy this position.
Reactions of the xenon fluorides
With the exception of the reactions of
powerful oxidative fluorinating agents
such as PtFe, the chemistry of xenon hastastemmed from reactions of the binary
fluorides.
• As fluorinating agents. The difluoride
is a versatile mild fluorinating agent. 17•18
It provides routes to some aromaticcompounds which are otherwise difficult
to make and reacts with carbon-carbon
double and triple bonds, giving addition
products. The cleanliness of these reactions is of special significance with someof the products, which are of medicinaland biological importance. One example
is the ability of XeF 2 to convert uracil to
5-fluorouracil, one of the first antitumour
agents.
Xenon difluoride is also a common
inorganic oxidative fluorinating agent,
see equations (aHd).'7
CH31 + XeF 2 HF

2Sb + 3XeF 2 -HF

SbF3 + XeF2 -CH31F2 + Xe(a)

2SbF 3+ 3Xe(b)

SbFs + Xe(c)

CFzCI.CFCl z
Re~CO),o + 3XeF 2~

Re(CO)sF.ReFs + 5CO + 3Xe (d)Recently it was used in the preparation of
CrOF3.19 Soluble also in water, thedifluoride is stable in neutral or acid(e)

Thus the hydrolysis of both XeF4 and
XeF6 give the trigonal pyramidal Xe03 insolution. Aqueous solutions of Xe03
(known as 'xenic' acid) are quite stable
but solid Xe03, which is easily obtained
from the solutions, is a sensitive andviolent explosive.
• Reactions giving oxides, oxide fluorides
and oxygen-containing salts. In aqueous
solution Xe03 is a powerful, but kinetically slow, oxidising agent. The addition of alkali yields 'xenate' ions,[HXe04L and, although the salts can beisolated, solutions containing xenate
disproportionate readily to give the'perxenate' ion, [Xe06]4-, and xenon:
2[ HXe04r + 20H- - [Xe06r- + Xe + O2+. 2H 20 (t)The Xe(VIII) solutions are among themost powerful oxidising agents known.
For example, the oxidation of Mn(lI) to[Mn0 4L unlike that by Xe03, is immediate and is accompanied by the evolutionof oxygen.
The addition of solid Ba2XeOS to cold
concentrated H2S0 4 produces a secondoxide, Xe04, which is an unstable andexplosive gas. Understandably, it has
been studt.ed little, but infrared spectroscopy and electron diffraction haveshown it to be a tetrahedral molecule.
Of the known oxide fluorides, XeOF2,.xeOF4, Xe02F2 and Xe03F2, the moststable is the liquid, XeOF 4. Colourlessand volatile, with a square pyramidal
structure, it derives from the controlledhydrolysis of XeFs:
XeFs + H20 - XeOF 4 + 2HF(g)

This reaction, equation (g), readily goes
further to produce Xe03, but the expected intermediate, Xe02F2, has neverbeen isolated from the reaction. But,

662

CHEMISTRY IN BRITAIN JULY 1987

Xe02F2 can be produced via the interaction of xenon trioxide with the oxide
tetrafl uoride.
Xe03 + XeOF 4 - 2Xe02F2(h)

The related and possibly even morehazardous reaction with Xe04 yields
Xe03F2:
Xe04 + XeFs - Xe03F2 + XeOF 4(I)• Substitution reactions. Reactions involving the replacement of fluorine inxenon fluorides with other groups beganwith studies of reactions of XeF 2 withanhydrous oxygen-containing acids: 17
XeF 2 + HL -F-Xe-L + HF
and XeF 2+ 2HL - L-Xe-L + 2HF

(J)
(k)

.In all cases, the linkage to xenon is
via oxygen (L = OTeF 5, OSeF 5, OS02F,
OPOF 2, ON0 2, OCI0 3, OIF 40, OCOCH 3
or OCOCF 3). Only those species containing the highly electronegative -OTeF5
and -OSeF5 units are stable underordinary conditions. The -OS02F and-OIF 40 derivatives decompose readily
and the remainder are explosive.
The high electronegativity of the-OTeF5 group means that Xe(OTeF 5h issimilar to XeF 2 and an interesting reaction of Xe(OTeF 5h is that with halogenoalkanes such as CF 2=CFCI, CF 2=CCI 2 and
CF 2=CFH, which results in the addition of-OTeF5 groups across the double bond. 21
The development of substitution reactions involving Xe(IV) and Xe(vI) compounds has required the use of thereagent B(OTeF 5h, an obvious relative of
BF 3. Reactions involving this molecule
have given the only Xe(IV) compound
to date, Xe(OTeF 5)4, while Xe(VI) gives
Xe(OTeF 5)s, OXe(OTeF 5)4, 02Xe(OTeF 5hand, under certain conditions in solution,
the related partly substituted derivatives.
Like many of the Xe(II) compounds, mostof these are also unstable.
Because noble gas chemistry is mostly
associated with Xe-F and Xe-O bondformation the key to the successful bonding of xenon to other atoms to yield stablecompounds lies in producing suitablehighly electronegative groups. Such· agroup is -N(S02Fh which has provided
the route to the first Xe-N bond: 22

F-Xe-N(S02Fh(I)

The solid, once formed, is reasonably
stable. It contains a linear F-Xe-N
group and a planar configuration about
the nitrogen. 22.23 Its reaction with AsF5
yields the solid, [(FS0 2hN-Xe-F-XeN(S02 FhnAsFsf, containing a V-shaped
cationic relative of [Xe2F3f.
Although, in general, the noble gases
form bonds with only the most electronegative elements, fluorine and oxygen,the success in making an Xe-N bond incompounds has increased speculation as
to whether an Xe-C bond is possible. A
plasma-induced reaction of XeF 2 vapour
with CFj radicals may have provided theanswer. The volat,ile, waxy, white solid

product of the reaction, which decomposes to give XeF 2 and carbon fluorides
at room temperature, is thought to beXe(CF3h·24• Reactions with fluoride-ion acceptors
and donors. 25 Reactions of fluoride-ion
donors and acceptors with the binary
fluorides and oxide fluorides of xenon
give a variety of cations and anions(Table 1). More unusual entities such as[Xe2f and [Xe03Xr (X = F, Cl or Br) also
occur. The representation of many of thecomplexes as salts, however, is an oversimplification, especially for the xenon
difluoride adducts. Also, some adducts of
XeF 2-for example, those with XeOF 4
and the halogen pentafluorides-are notdonor-acceptor compounds but simple
molecular adducts in which the component molecules preserve their molecular identities.
The enthalpies of ionisation of thebinary fluorides:
XeFs(g) - [XeF 5r (g) + F-(g)tlH=9.24eV(m)

XeF 4(g) - [XeF 3r (g) + F(g)tlH= 9.66eV (n)XeF 2(g) - [XeFr (g) + F-(g)tlH= 9.45eV (0)imply that complex formation might beprogressively easier in the sequence
XeF4 < XeF 2 < XeF s. The chemical
evidence supports this. Xenon difluoride
reacts wjth AsF 5, SbF 5, BiF 5, and arange of transition metal pentafluorides, to give adducts that are often
written [XeFnM 2Fll L [XeFnMFsr and[Xe 2F3nMF sr but in which the cation
units are linked to the counter ion byfluorine bridges; the V-shaped [Xe2F3f
ion can be regarded as two [XeFf units
linked by a fluorine bridge. The mostionic adducts are those derived fromthe strongest Lewis acids. The related
derivatives of MOF4(M = W or Mo), which
are weaker Lewis acids, have lesS ioniccharacter.
One of the most interesting reactionsof [XeFf is its reduction with somecompounds to give the bright green
paramagnetic cation [Xe2r.2S This cation
can also be obtained by oxidising xenon
with dioxygenyl salts. Furthermore,
Bartlett has shown that the first noble
gas compound, Xe+[PtFsL has thevariable composition, Xe(PtFs)x (x =1-2).
Whether this compound contains Xe+,
[Xe 2f or [Xe2Ff or can be described interms of [XeFnPtFsr and [XeFnPt 2Fll r is
still uncertain.

In accord with the enthalpy of ionisation data, xenon tetrafluoride forms
complexes with only the strongest fluoride ion acceptors-SbF 5 and BiF5-but
xenon hexafluoride yields adducts with anumber of pentafluorides and also withsome tetra- and trifluorides. Solutions of
XeFs in HF also have a higher conductivity
than those of XeF 2 and XeF4 because
both [XeF 5f and [Xe2Fllf are readily
formed in solution:([XeF 5fF-)4 + nHF ~2[Xe2Fllf + [(HF)nFr(P)[Xe2Fllf+ nHF ~2[XeF 5f + [(HF)nFr(q)

These same two ions are found in
the adducts. The [XeF 5f ion has thesquare pyramidal form found in XeFs and[Xe2Fllf is best described in terms of anionic model, [XeF 5fF [XeF 5f, but withsome covalent character in terms of([XeF 5fXeF s) to account for the nonlinearity of the Xe-F-Xe bond.
A relatively poor fluoride ion donor,
xenon oxide tetrafluoride, forms adductswith only the stronger Lewis acids.
Thus, with SbF 5, [XeOF 3nSb2Fll r and[XeOF 3nSbFsr are obtained. The related dioxide difluoride, Xe02F2, yields[Xe02Ff[Sb2Fl1r and [Xe02FnSbFsL
but these are less stable. Reaction of
FXeOTeF 5 with AsF5 gives the complex[XeOTeF 5nAsF sr and reactions of thiscomplex give rise to a number of as yet
uncharacterised xenon cations.
Fluoride ion acceptor behaviour islimited to XeFs. With alkali metal fluorides
it yields M2XeFS (M =Cs, Rb, K and Na)
and MXeF 7 (M = Cs, Rb), which lose XeFs
when heated. The thermal stability ofthese adducts increases with increasing
molecular weight and the low decomposition temperature of the sodium complex, which is below 100 °C, provides aconvenient way of separating XeFs fromXeF 2 and XeF 4 with which no reactions
occur. Solid state structures have notbeen determined, but the related complexes derived from NOF and N02F havebeen prepared and [NO]i[XeF s]2- hasbeen shown to contain a slightly distorted
square antiprismatic [XeF s]2-. Valenceshell electron pair repulsion theory suggests that a ninth coordination position inwhich a lone pair can be accommodated
should be present in this anion. However,there is no evidence for it, and this
implies a rare violation of the theory innoble ga'S chemistry.

Table 1 . The cationic and anionic derivatives of binary fluorides and oxide fluorides
ofxenon. 25
Fluorides

Cationic derivatives

XeF 2
XeF.

[XeF 3f

Anionic derivatives[XeFr; [Xe2F3r[XeF 6r; [Xe 2 F"r[XeOF 3r

[XeF 7L [XeF e]2-

XeOF.
Xe02F2[ Xe0 2 Fr

[Xe0 2 F3r[XeOF 3r

XeFe

XeOF 2

r: [Xe30 3F13r[XeOF 6I

664

HEM I TRY I N BRITAIN J

Scheme 1.

used as fluorinating agents in industrial
processes and radon chemistry being
applied in areas of radon monitoring andatmospheric cleansing . Had it not beenfor Bartlett's astute chemical awarenesS,
however, it is possible that we might still
be talking of the 'inert gases'.

The reaction of XeOF4 with CsF.

CsF + XeOF 4 (excess)oDC, vacuum)- XeOF4

LY 1987

cSF.xeoF4 1
20 °C, vacuum- XeOF4

Dr John H. Ho 110 way is a senior lecturer
in the department of chemistry, University of Leicester, Leicester LET 7RH.
References

Stable white solids containing thenM +[Xe0 3 F]~ anion consist of chains ofpseudo-octahedral xenon atoms linked
by angular bridges. These are depositedwhen
solutionsofstoichiometricamounts of Xe03 and HF are concentrated . The haloxenates, CsXe03CI andCsXe03Br have also been obtained.
Complexes of XeOF 4 with CsF, RbF,
KF and with NOF have been known forsome time . The former are obtained bytreating the alkali fluorides with an excess of XeOF 4 and pumping to constantweight under appropriate conditions (seeScheme 1). The structure of CsF'XeOF 4
contains an [XeOF 5f anion of distorted
octahedral geometry with a lone pairof electrons apparently occupying anoctahedral face . The CsF'3XeOF 4 adductcontains the [F(XeOF 4bf anion in which
the XeOF 4 units, of close to squarepyramidal shape, are linked via three
Xe---F bridging bonds to a common
fluorine which is out of the plane of thethree xenons. Trigonally bonded fluorine
is rare but, where it does occur (ie (XeF 6)6
and NaSb3FlO), the arrangements about it
are closely similar.
Radon compounds7.13.25.27
Radon chemistry has assumed increasing importance since · it emerged thatradon can be oxidised and complexed bycertain liquid fluorides and fluoride salts.
Tt. ;" offers the possibility of removing
radon from mine atmospheres and ofdeveloping radon -monitoring devices.27
The high radioactivity of radon has
meant that the chemistry of radon has
been explored by tracer methods. Whenmicrocurie amounts of radon are heatedwith fluorine a difluoride appears to beformed. The same species is also gen erated spontaneously when millicurie orlarger amounts of radon are mixed withgaseous or liquid fluorine in a small
volume; the latter reaction is brought
about by the intense a- radiation . Radondissolves in halogen fluorides as an in volatile species. Electromigration studies
have shown that in pure solvents theradon is present as a cation and ionisations involving the following equilibria
have been suggested :RnF2 ~[RnFf ~

[RnFf + FRn2+ + F-

(r)
(s)

The existence of the [Rn- Ff cation
seems certain in view of the spontaneous
reaction of radon with a variety of solids
such as [02nSbF 6f containing oxidising

cations. By analogy with xenon chem istry, the reaction with [02nSbF 6f is
probably:Rn(g) + 2([02nSbF 6f}(s) [RnFnSb2F11ns) + 20 2(g)(t)

Krypton compounds 13 .25
Krypton chemistry is more limited thanthat of xenon. Although sma ll amounts ofthe violet free radical, KrF, have been
observed in the y-i rr adiati on of KrF2'
most of the chemistry derives from thedifluoride. No oxides or oxide fluorides
have been isolated and ea rly reports on
the preparation of KrF4 have not been
confirmed.
The volatile, colou rl ess, crystalline di fluoride is prepared by subjecting mixtures of krypton and fluorine to an electric
discharge at low temperatures or byirradiating liquid krypton- fluorine mixtures with ultraviolet light. 28 The linear
molecular structure of KrF2 is similar tothat of XeF 2 but, because of its much
lower bond energy, it is a very powerful
fluorinating agent capable of oxidisingiodine to IF 7, xenon to XeF 6, low-valenturanium, neptunium and plutonium
compounds to the hexafluorides,29 and
AgF2 to AgF3.30 KrF2 is also rapidly
decomposed by water.
The adducts of KrF2 are analogous tothose of XeF 2 and are limited to cationicderivatives formed with fluoride ion
acceptors. Species such as [ Kr 2F3nMF 6(M =As, Sb, Ta), [KrFnMF6f (M =As, Sb,
V, Ta, Pt and Au) and [KrFnM2Fllf(M =Sb, Nb and Ta) are known in the so lid
state and in solution and, more recently,
compounds such as [KrFf[MOF 5f (M =Cr, 31 Mo and W 32) have been preparedand characterised . The spectra of theions are related to those of the xenon
analogues but additional peaks in the
Raman for [Kr2F3f have been interpreted
in terms of an unsymmetrical (rather
than symmetrical) cation. The oxidising
power of [KrFt and [ Kr 2F3f is strong er
than that of KrF 2, capable of oxidising IF5
to [IF6f, O 2 to [02f, Xe to [XeF 5f, BrF5 to[BrF6f and Au to [AuF 6f and [AuF 5].
Noble gas chemistry is now well established. Still inextricably linked withfluorine chemistry, it has provided ameans to accomplish novel fluorinations,of introducing us to new classes ofcompounds and new high oxidation state
species. In the next few years it promises
to take us into the rea lms of app liedchemistry, with noble gas fluorides beingr

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