CHEMISTRY OF THE NOBLE GASES*
By Professor K. K. GREE~woon , :.\I.Sc., sc.D .. r".lU.C.
University of N ewca.stle 1tpon Tyne
elements were unsuccessful, and for over
60 years they epitomized chemical inertness.
Indeed, their electron configuration, s2p6,
became known as 'the stable octet,' and thisfotmed the basis of the fit·st electronic theoryof valency in 1916. Despite this, many
people felt that it should be possible to induce
the iner t gases to form compounds, and manyof the early experiments directed to this end
have recently been reviewed.l
There were several reasons why chemists
believed that the inert gases might formchemical compounds under the correct conditions. For example, the ionization potential of xenon is actually lower than those ofhydrogen, n itrogen, oxygen, fl uorine andchlorine, all of which readily form covalent
compounds (Xe 280, H 313, N 335, 0 314,
F 402, Cl 300 kcalj mole). Furthermore, theionization potential of radon is very similar
to that of mercury, which read ily form.
cationic species (Rn 247, Hg 240 kcal/ mole).
There were also suggestive experimental
observations. Thus, we are all familiar withthe fact that iodine is virtually insoluble inwater but dissolves freely in the presence ofpotassium iodide to give the brown I 3 anion.
The iodide ion I - has exactly the same
number of electrons as xenon, which m ight
therefore be expected to behave very similady; apparently it does not. The existenceof stable interhalogen compounds, such as
BrF3 , IF5 and IF 7 , also indicates that more
electr·ons than are allowed by the classicaloctet theory can be involved in bonding andthat the isoelectronic species, such as KrF2 ,XeF,1 and XeF6 , might be capable of independent existence.
Finally , t here was the possibility of forming
co-ot·dination complexes by donating a. pairof electrons fmm an inert-gas atom to anotheratom. After all, the inert gases are bristling

The inert gases, or noble gases as they are
now more appropriately called, are a remarkable group of elements. The lightest, helium,
was recognized in t he gases of the sun before
it was isolated on ea.r th as its name (i]A.tos)
implies. The first inert gas was isolated in1895 by Ramsay and Rayleigh; it was named
argon (apy6s, inert) and occurs to the extent
of 0·93% in the earth's atmosphere. The
other gases were all isolated before the turnof the century and were named neon (v€ov,new), krypton (KpVn'TOV, hidden), xenon~€vov, stmnger) and radon (radioactive
emanation). Though they occur much lessabundantly than argon they cannot strictly
be called rare gases; this can be illustrated
by calculating the volumes occupied a.t s.t.p.
by the gases in this lecture theatre (volume~ 106 litres): helium 5·24 1, neon 18·3 I,
argon 9,340 I, krypton 1·141 and xenon 87 ml.
A few chara-cteristic properties a re listed
in Table I , which shows tha.t the inert-gas
atoms have completes and 1' subshells. For
TABLE I
SOX.IE PROPERTIES Ol' THE

At.
No.

Ho

2
No 10
Ar 18
Kr 36
Xe 54Rn 86"OD LE ()ASJ::S

Ionization Boiling
E lectron potential
point
(oC)
config. (kcalj mole)

D.Hvap

- 269
- 246
- 1 6
- 153
-107
- 65

0·022
0·44
1·50
2·31
3·27
4·3

l s2
2s'2p 63s 2 3p6
4s 2 4p6
5s2 5p6
6s2 6p6

567
497
363
323
2 0
247

(kcal/
mole)

this reason there are only weak van der'Vaals interactions between the indi,'idual
atoms and this leads to the very low boilingpoints and heats of va.porization shown. The
ionization potentials (energy required to remove one electron from an atom to give a
cation, ]If+) are very high compa.red witht hose of the alkali metals (90-124 kcalj mole).
In fact, all attempts to prepare well-defined
stoichiometric chemical compounds of these* This article is based on a lecture given to
sixth -form pupils on 20 and 25 November, 1963,
and to the Bedson Club of the University of :><owcastle upon T yne on 6 December, 1963.
176

177

CHEl\'IIS'l'RY OF THE NOBLE GASES

with lone-pairs and their ionization potentials- are often no greater than those of ammonia,
water, ethylen e and other well-known ligands .
However, recent detailed studies of the interaction of xenon and krypton with p owerfulelectron-pair acceptors, such as boron trifluoride and trichloride, indicated that there
is no specific donor- acceptor bond-formationeven at very low temperatures.
This was the position at the beginning of
1962. Then, in ,June of t hat year , N. Bartlett
announced2 t hat when xenon was placed incontact with gaseous platinum hexafluoride
at room temperature an immediate reaction
occurr ed to give an orange-yellow solid:Xe (g) + PtF6 (g) = XePtF6 (s)Xe (g) + 2F2 (g) = X eF 4 (s)The compound can even b e made simply bypassing a mixt ure of xenon and fluorine
diluted wit h nitrogen through a nickel tube
h eated with a bunsen burner ; white cryst alsof xenon t etrafluoride sublime out--an experiment t hat could have been done at an ytime during the past 60 years !
TABLE IIPROPERTIES OF THE XEKON FLUORIDES

X eF 2

XeF4

X eF6

140
4·32
Density (g/ ml)Sol. in liq. HF at 25° (g/ 100 g) 175
H eat of solu tion (kcal j mole)
2·5
Vap . press. (mm H g a t 25°) . .
3
H eat of sublimation (k cal/mole )
..
..
12·3
2·00
Bond length Xe- F (A)

114
4·04
4
6·7
4

46

250
18
30

15·3
1·95

9·0
1·91

:W.p. (00)(3) Stability and bonding theory.
PREPARATION, PHYSICAL PROPERTIES

Xenon hexafluoroplatinate(v) was the first
au t hentic compound of a noble gas. The
idea was taken up by dozens of different
laboratories throughout the world, and within
18 mon t hs a 400-page book on the subject
had been published.3 This gives some idea.
of the speed at which modern advances inch emistry are made.
The first group to start '\vork after t heinitial discovery was one at the Argonne
National L aboratory in Chicago, and three
mon t hs after Bartlett's paper they reported4
that xenon combined directly with fluorine
when the two gases were heated in a nickel
vessel under pressure at 400° :Proper ty

About 20 compounds of xenon have now
been isolated and many more identified inr eaction systems. Compounds of krypton
and radon h ave also b een made, and asystematic account of t his n ew area ofchemistry will now be given. Detailedreferences to the original literature are given
in refs 1, 3, 5 and 6. It will be convenient
to break up the discussion into three parts:(1) Preparation, physical properties andstructures ;(2) Ch emical reactions;--- - - - -- - -AND STRUCTURES

Xenon Difluoride
XeF2 is a white crystalline compound which
is most simply prepared by irradiat ing a
mixture of xenon and excess fluorine for one
day at 25° and 1! atm pressure with lightfrom a high-pressure mercury arc. It is important to freeze out the product at - 78° as
it forms, otherwise some xenon tetrafluoride
is also formed, and these two compounds arevery difficult to separate. Amounts up to
10 g can be prepared in this way in silica
vessels or in nickel vessels with sapphire
windows. The compound can also be prepared by irradiating a xenon- fluorine mixture
with 6°Co y-rays, or by circulating a 1:4 mixture of the gases through a hea ted nickel tube
at 400° and trapping t he product at - 50°.
Some properties of xenon difluoride are summarized in Table II. It can be seen t hat t hemelting p oint, 140°, is far above t h e m.p. ofboth xenon ( - 111 °) and fluorine (- 223°) andthe density is also greater than t hat of solid
xen on (3·1) or fluorine (1·3). Xenon difluoride has a very high solubility in anhydrous hydrogen fluoride, in which it isnon -conducting.
The vapour pressure is sufficient to allow
the compound to sublime readily when
warmed under vacuum and it forms beautiful
crystals in this way. The h eat of vap orization (12·3 kcalj mole) suggests that there isappreciable separation of charge within themolecule, and this coulombic attraction contributes to t he low volatilit.y (cf. 6.Hvap Xe178

EDUCATIO~

3·27, F 2 1·64 kcalj mole). Xenon difluoride
is diamagnetic, indicating that there arc no
unpaired electrons in the structure.
Neutron diffraction and X-ray diffraction
have shown that xenon difluoride crystallizes
in the tetragonal system " ·ith cell constants
a = 4·315 and c = 6·990 ± 0·004 A, and twomolecules in the unit cell, as illustrated inFig. 1. The XeF2 molecule is linear and thisr

4·3 15 2.

1?

I
2·00 2.

IN CHEUISTRY

X e1wn T etrafluoride
XeF4 is best prepared by heating xenon
and fluorine in an atom ratio of 1 : 5 at
13 atm and 400° in a nickel weigh ing-canfor 1 lu·. Up to 50 g can be ha.ndled a.n d
the yield is quantitative. Flow methods
tlu·ough a nickel tube, electrial discharge
methods and y-irradiation haYe a lso beenused. The compound sublimes readily under
reduced pressure a-nd is obtained as beautifully-formed colourless ct'ystals.
The properties of xenon tetl'afiuoride aresummarized in Table II. Compa r·ison withxenon difl.uoride shows that the tetrafluoride
has a lower m.p. and density, and a considerably lower solubility in anhyd roushydrogen fluoride; the heat of vaporization,
however, is 25 per cent higher. Neutron
diffraction and X-ray diffraction results indicate that t he tetrafluoride crystallizes in the
monoclinic system with two molecules in theunit cell, as shown in F ig. 3. The molecule!

1---- 5·03~ - lbFig. 1.

Crystal structure of :XoF ~is also consistent with the infra-red a nd
Raman '\ibrational spectra (both band shapes
and occurrence). The modes of '\ibration are
shown in Fig. 2.
-F--Xe - - F - v ,497 em-•Raman-F--Xe-Fv,

555 em- ' Infra-red

F-~e- F\/Y,

213 em-•F ig. 2.

Infra-redVibrational spectrum of XeF 2

lea
Fig. 3.

Crystal structure of XeF4 ('exploded' v iew)

is square-planar with the F- Xe- F bond angle
90°, and the Xe-F bond distance slightly lessthan that for the difluoride. The molecules
pack in a body-centred array, and the angle
between the planes of the two sets of molecules is 55·2°.
The infra-red and Ra.man spectra of xenon
tetrafluoride are also best interpreted on thebasis of a square-planat· model. Details are
given in Table III. A molecule containing
five atoms should have nine (i.e. 3n- 6) modes

CHEMISTRY OF THE ~OBLE GASES

of vibration, but if it has square-planar
symmetry two pairs of these should be
doubly degenerate, that is, have the sameenergy; v 6 represents one of these pairs
and v 7 the other , so that only seven fundamental modes are expected, of which one,v4 , is inactive in both infra-red and Ramanand can only be observed as an overtone. I t
will also be seen that no fundamental vibrat ion in the infra-red appears in t he Raman
spectrum and vice versa. This 'rule of mutual
exclusion' indicates that the molecule has acentre of symmetry and confirms that it is
not tetrahedral. (A tetrahedral molecule
gives rise to only four fundamental modes,
all of which give Raman lines and two of
which also give lines in the infra-red ; theinfra-red modes are triply degenerate andone of the R aman modes is doubly degenerate,
making a total of nine modes in all, as
required for a pentatomic molecule.)
TABLE IIIVIBRATION SPECTRUM OF XENOX TETRAFLUORIDE(R)

543 cm- 1

v2 (I.R)

291 cm- 1"• (R)

235 cm- 1"• inactive(22 1 cm- 1)"• (R)

502 cm- 1"• (I.:R .)586 cm- 1

v, (I.R)

123 cm- 1

VJ(R) = Raman;

I
I

X
+X+

+ +
X
+
x- +

X
X

)\

(I.R.) = Infra-red

179

Xenon IlexajluoTide
XeF6 is less stable than the difluoride andtetrafluoride. It has been prepared in 10-g
amounts by heating xenon in_a . large excess
of fluorine at pressures up to 200 atm andtemperatures of 300-400°. In a typical experiment a 95 per cent yield of xenon hexafluoride was obtained by heating Xe + 20 F 2
for 16 hr at 300° and 60 atm. Electricdischarge methods have also been used.
Some properties of xenon hexafluoride arc
listed in Table II. The low m.p. and highsolubility in anhydr ous hydrogen fluoride are
noteworthy. Unlike solutions of the difluoride and tetrafluoride, these solutions aregood conductors of electricity. The highvapour pressure and low heat of vaporizationr eflect the more symmetrical structure ofthis molecule. Xenon hexafluoride is thermochroic: the colourless crystals turn paleyellow at 42° and melt to a yellow liquid ;t he vapour is a pale yello\1·-green. All thesechanges are reversible and find explanationin terms of the detailed theory of bonding
and the ene1·gy-level diagram for thecompound.
X-ray and neutron-diffraction studies have
proved difficult because of the reactivity of
the hexafluoride. For example, glass, silicon
and mercury are all rapidly attacked. However, the structure is certainly octahedral,though it is not yet certain whether all bondlengths are equal. The infra-red and R aman
spectra have also only been partly elucidated; there appear to be no coincidences,indicating a centrosymmetrical molecule.
The Xe-F symmetric stretching vibration
occm s at 655 cm- 1 in the Raman spectrum
and the Xe-F antisymmetric stretchingmode is at 612 cm- 1 in the infra-red.
Other Xenon Fluorides
Irradiation of xenon tetrafluorides with
60Co y-rays at - 196° yields xenon monofluoride radicals, XeF, trapped in the crystal
lattice. This species, which is stable only
at low temperatures, has one unpair edelectron; it is paramagnet ic and colours t hehost lattice bright blue.
Crystals of composition XeF3 can be grownfrom the vapour of xenon tetrafluoride containing some difluoride. The compound is

180

EDUCATlOK IN CHEMISTRY

diamagnetic (no unpaired electrons) andcolourless. X-ray crystallography shows thatthe lattice consists of equal numbers of X eF 2
and XeF4 units packed so that the xenon
atoms form a face-centred array. The
crystals at'e monoclinic with cell dimensions
a= 6·64 A, b = 7·33 A, c = 6·40 A, {3 =92° 40'. The density is 4·02 g j ml and all
bond distances and bond angles are normal.
The preparation of xenon octafluoride has
been claimed but not confirmed. Thus, whenxenon was heated with 16 mole-equivalentsof fluorine at 620° and 200 atm and thereaction vessel then chilled to - 78°, yellow
crystals of composition XeF7 .9 0 .3 were said
to be isolated. At room temperature thecompound is apparently an unstable yellow
gas.
.l.

Xenon Oxytetrajtuoride
Xe0F4 is formed when xenon hexafluoride
is kept in glass or si]jca vessels:
2XeF6 + Si02 = 2XeOF, +SiF4

Xenon hexafluoride reacts completely withquartz at 50° in two days. The stoichiometricamonnt of '-rater vapour also hydrolyses thehexafluoride to this product:a centre of symmetry. The spectra of theoxytetrafluoride should also be comparedwith those of xenon tetrafluoride in Table III ;there is a marked similarity except for thetwo extra bands associated with the Xe-0
stretching and wagging modes.
Forceconstant calculations clearly indicate thedouble-bond character of the Xe= O bond.
Similarity with the isoelectronic interhalogen
compound, bromine pentafluoride, is also
notewot·thy.
The molecular integrity of xenon oxytetrafluoride has been further demonstt'ated in atime-of-flight mass spectrometer which also
indicated the presence of the species Xe02F 2 .
In both these compounds xenon is in theoxidation state vr. Another oxyfl.uoride.
Xe0F2 , in which it is in the oxidation staterv, is formed when xenon reacts with oxygen
difluoride in an electric discharge, but little
is known of its properties.
Xe + F 2 0 = Xe0F2

X e1wn '1'1·ioxide:X.e03 was first prepared by an undergraduate at Berkeley University; it was
TABLE IV

XeF6 + H 20 = XeOF4 + 2HF

Xenon oxytetrafluoride is a clear, colourle:sliquid at r oom temperature. It freezes toa white solid at - 40° (or - 28°) and has avapour pressure of 8 mm at 0° and 29 mm
at 23°. The heat of vaporization is approximately 9 kcalj molc.
The infm-red and Raman spectra establi hthat the molecule is a square-based pyramid
0 4 " symmetry) as shown in Fig. 4. The

V!BllA'l'IONAT, SPECTRUM Ol' XE.\'ON
OXY1.'E'l'RAFLUORlDE

)[ode

'11
')2

'Ia

'1,

0

/~-11-7,'
!/
~/
I

Xe

'

F - - - - - - - --- F
Fig. 4.

Stnteture of Xe0F4

detailed assignments are listed in Table IV.
from which it can be seen that there areseveral coincidences in the Raman and infrared spectra, consistent with the absence of

"•
'Is

'17

'Is
'19

I.R.
(gas)

Raman
(liquid)
om- 1

cm- 1

919 (s, p)
566 (vs, p)

1928
(s)
578 (w)

Assignment

X o-0 stretch
X e-F, symmetricstre tc h28() (vw,p)X e-F, sym. out1 288 (s)of-planode formation
231 (w, dp) 1 inactive X e-F 4 sym. in planodeformation
inacti,·e X o-F, out-of530 (s, dp)
phaso str·etch
not obs.
1 inacth·e Xo- F, antisym.
out.of-planed eformation
not obs.
609 (vs) X o-F,1 a ntisym.
stretch364 (mw, dp) 362 (ms) Xe-0 wag
161 (vw, dp) not obs. X e-F4 antisym.
in-planoIdE' formationII

vw, very weak: mw, mediwu-weak; w, weak:
ms. medium-stmng: s, strong; YS, vcr·y sh·on?::p . polarized: dp. d epo l ~trized.

CHFJ:\l£STRY OF THE - OBLE GASES

obtained as white, non-volatile, highlyexplosive crystals by evaporating the aqueousacid hydrolysate of xenon tetrafluoride. During the hydrolysis abont half of the x err
dispropm·tionates into Xe0 a nd x en, andthe other half oxidizes water to oxygen; theoverall reaction can be represented approximately by the equation:
3XeF, + 6H20 = 2Xe + £ 0 2 + Xe03 + 12HF

The trioxide can also be prepared b.v . lowhydrolysis of the hexafluoride :XeF6 + 3H20 = Xe03 + 6HF

The degree of hydration of the tr ioxide is
unknown but oxygen atoms can be exchangedby equilibration with 180-enriched water.
and it seems probable that the reaction is
Xe03 + 3H20 ~ Xe(OH) 6

X-ray structural work on the crystals showsthat xenon trioxide has an orthorhombic unit
cell with a = 6·163 ± 0·008 A. b = 8·115 ±
0·010 A and c = 5·234 ± 0·008 A. There are
four molecules in the unit cell a nd the density
is 4·55 g/ ml. Xe03 is a trigonal pyramidal
molecule as shown in Fig. 5. The bondv, 331 em-'v, 820 em-'Fig . .3.

Vibrational speclnun of X e03

length Xe-0 is 1·76 ± 0·04 A and the bondangle 0 - Xe-0 is approximately 103°. The
molecular symmetry and dimensions are very
similar to those of the isoelectronic iodate
ion, 103 - , for which t he mean bond lengths
and bond angles are 1·82 A and 97<>.
Infra-red assignments are shown in F ig. 5.
There is a marked similarity to t he spectra

181of the ions Cl03- , Br03- and I 0 3- , a nd theprincipal stretching-force constant indicatesconsiderable double-bond character, consistent with the short Xe-0 bond distance.
X enon T etroxide

Xe04 is a volatile yellow solid, unstable atroom temperature. I t is prepared 7 by adding
sodium or barium perxenate (see next section)
to anhydrous sulphuric acid at -5°, thereaction being formally represented by t heequation:
Ba2Xe0 6 + 2H2S0 4 = Xe0 4 + 2BaS0 1 + 2H20

Yields of up to 34 per cent (0·1 g) wereobtained.
The xenon tetroxide molecule is tetrahedral (like osmium tetroxide and theperiodate ion, ro<l-) and has two infra-red
active vibrations as expected: v3 , 877 cm-1 ;v4 , 305·7 cm- 1 . The \apour pressure is
3 mm at - 35° and 25 mm at 0°, whichcorresponds to a heat of sublima tion of
8·0 kcal j mole.

Compounds of K rypton and Radon
Compounds of krypton are less stable thanthose of xenon and this has limited theamount of attention they have received.
Compounds of radon may well p r·ove to be
more stable, but difficulties of handling thishighly radioactive element have also precluded extensive investigation.
Krypton, unlike xenon, does not reactwith platinum hexafluoride at temperatures
up to 50° nor does it react when heated withfluorine. However, the species KrF2 hasbeen identified by matrix-isolation techniques. A mixture of fluorine and krypton
with argon as the matrix was spt·ayed on toa caesium-iodide window at - 253° and thenirradiated with ultra-violet light. Bands
appearing in the infra-red spectrum could be
assigned to the linear molecule KrF2 :v3 , 580 cm- 1 ; v2 , 236 cm- 1 .
Krypton tetrafluoride, K.rF4 , is s!.\id to beformed as transparent, colourless crystals
when an electric discharge is p assed througha gaseous mixture of the elements. Thus,'vhen a mixture of composit ion Kr + 2F2 at
12 mm pressure was passed through an arc

182

EDUCATION IN CHEMISTRY

of approximately 1,000 V/ 30m A between copper electrodes and the products trapped out
at - 188°, a yield of 1·15 g KrF4 was obtainedin 4 hr. The compound has a vapour
pressure of 4 mm at - 30°, 30 mm at 0° and
90 mm at 20°, and a heat of sublimationof 8·8 kca.lj mole. At 20° one-tenth of asample of krypton tetrafluoride decomposedin one hour and decomposition was very
rapid above 60°.
Hydrolysis of krypton tetrafluoride atroom temperature results in the liberationof krypton and oxygen :KrF4 + 2H2 0 = Kr + 0 2 + 4HF
However, slow hydrolysis by ice at -30° to-60° gave 2- 3 per cent of an acid whose
stable barium salt could be isolated at room
temperatme. 8 Hydrolysis of krypton tetrafluoride with a queous barium hydroxide at
0° gave a 7 per cent yield of a compoundconsidered to be barium kryptate, Ba.Kr04 .
The acid may be Kr03 ,xH20, for example
Kr(OH) 6 ; as the stoichiometric formula ofthe acid and the oxidation state of thekrypton have not yet been established it isperhaps appropria.te to call it kryptic acid.
Tracer experiments at the 10- 100 f-tC levelindicate that a radon fluoride, RnF "" can beprepared from mixtures of the gases at 400°
but the stoichiometry of the reaction has
not yet been established. Radon fluoride isstable, distils at 230-250°/10- 6 mm andappears to be an ionic compound. It is
reduced by hydrogen to radon and hydrogen
fluoride at 500°. Attempts to prepare radon
chloride and radon oxide by thermal or
photochemical means have so far been
unsuccessful.
CHEl\IICAL REACTIONS OF XEKOX COl\fPOUKDSThe chemical reactions of the compoundsof xenon have been extensively studied andsome typical examples will now be discussed.
Non-aqueous Chemist1·y
The fluorides of xenon react quantitatively
with hydrogen at about 400° and this has
been used as a method of analysis.

XeF2 + H 2 = Xe + 2HF
XeF4 + 2H2 = Xe + 4HF
XeF6 + 3H2 = Xe + 6HF

Xenon oxytetrafluoride reacts similarly witha large excess of hydrogen at 350° :XeOF4 + 3H2 = Xe + H 20 + 4HF
Reaction with gaseous hydrogen chloride ismuch more rapid and can be carried out
even at - 78° :XeF4 + 4HCI = Xe + 4HF + 2CI 2
It is noteworthy that no indication of axenon chloride was detected in this reaction.
Xenon tetrafluoride is a good fluorinating
agent; it attacks mercury at room temperarure and a solution of the tetrafluoride inanhydrous hydrogen fluoride reacts quantitatively wit h platinum:XeF4 + 2Hg = Xe + 2HgF2XeF4 + Pt = Xe + PtF4
Lilmwise the tetr afluoride fluorinates boron
trichloride to a mixture of boron chloride
fluorides, BCI.,F 3 _x, and carbon tetrachloride
yields CC13 F. Nitrogen dioxide is fluorinatedto nitryl fluoride, whereas sulphur tetrafluoride and selenium tetrafluoride are slowly
oxidized to the corresponding hexafl.uorides:
XeF4 + 4N02 = Xe + 4N02F
XeF4 + 2SF4 = Xe + 2SF6
XeF4 + 2SeF4 = Xe + 2SeF6
Liquid ammonia ignites and r eacts explosively
on contact with xenon tetrafluoride; xenon,
nitrogen and ammonium fluoride are among
the products. The reaction can be represented formally as the initial formation ofhydrogen fluoride, which then reacts withexcess of ammonia to form ammonium
fluoride:
3XeF4 + 4NH3 = 3Xe + 12HF + 2N2
12HF + 12NH3 = 12NH4 F
Xenon tetrafluoride is itself quantitatively
oxidized by dioxygen difluoride at - 78°
according to the equation:
XeF4 + 0 2F 2 = XeF6 + 0 2
This reaction represents a 'chemical' synthesis of a xenon fluoride which involves
fluorination with a compound of fluorine
rather than by the element itself. Otherchemical syntheses have also been achieved.
Thus, when xenon is heated with 2·2 moleequivalents of CF30F at 250 atm and 225°,

183

CHEThfiSTRY Olt' THE NOBLE GASES

xenon difluoride is formed. A similar reaction with FS020 F at 150 atm and 175° also
produces xenon difluoride. The same compound is formed when xenon and carbon
tetrafluoride are passed through an electric
discharge.
It will be recalled that the first authentic
example of a stoichiometric compound of anoble gas was xenon hexafl.uoroplatinate,
XePtF6 . S imilar compounds are formed byother reactive hexafiuorides, e.g. XeRhF 6(deep red), XeRuF6 and XePuF6 . However ,no compounds were formed with the stabler
hexafluorides, IrF6 , UF6 and NpF6 .
When xenon reacts with an excess ofplatinum hexafluoride other compounds areformed with compositions ranging up toXe(PtF6 ) 2 , but their structure has not been
definitely established and they have beenformulated both as ionic and as fluorinebridged covalent compounds: Xe2 +(PtF6 - ) 2
and F 5Pt-F- Xe- F - PtF5 . The ionic structure
is suggested by the infra-red spectrum, whichshows only two absorptions as expected for
the PtF6 - ion. This is also consistent withthe fact that xenon hexafl.uoroplatinate(VI)reacts with rubidium fluoride in iodine pentafl.uoride as solvent t:> form the well-lmown
orange-red salt, RbPtF6 . Likewise, caesium
fluoride reacts with Xe(PtF6 ) 2 in the same
solvent to give CsPtF6 .
Pyrolysis (1 hr at 165°) of a solid having
the composition Xe(PtF6 h.8 gave xenon
tetrafl.uol·ide and a new, brick-red, diamagnetic compound, XePt2F 10 .
Xenon tetrafluoride does not form complexes with sodium fluoride or potassium
fluoride at temperatures up to 200° or withboron trifluoride at temperatures up to 100°.
It therefore shows little tendency to act either
as an electron-pair acceptor or an electronpair donor (ligand). However , xenon tetrafluoride dissolves in antimony pentafl.uoride
to give a gas and a yellow, diamagnetic
complex, XeF2 ,2SbF5 , m.p. 63°. The compound can be obtained directly from xenon
difluoride and distils unchanged under reduced pressure at 120°. A similar strawcoloured complex, XeF2 ,2TaF5 , m.p. 81o , isobtained from xenon tetrafluoride and tantalum pentafluoride. On the basis of theirvolatility and diamagnetism it seems prob-

able that these adducts are bonded bybridging fluor ine atoms, F 5M-F-Xe-F-M_F5 .
Aqueous Solution Ghernistry
The aqueous solution chemistry of xenon
is fairly extensive and somewhat complicated. Species containing Xevr and Xe vm
are well established and several salts havebeen isolated.
Xenon difluoride hydrolyses very slowly
in water but more rapidly in alkaline
solution:
XeF2 + H 2 0 = Xe + }0 2 + 2HF

All the combined xenon is liberated as theelement, indicating that 'hydrolysis' involves
oxidation of water to oxygen and reductionof Xen to Xe0 •Hydrolysis of the teteafluoride gives varying peoducts, depending on the conditons;hydrolysis is again accompanied by oxidationof water to oxygen and also by the disproportionation of xenon(rv) to higher and lower
oxidation states. During the hydrolysis thesurface of the tetrafluoride becomes brightcanary-yellow, and this yellow solid diss0lves
with evolution of gas to give a colourlesssolution. The lifetime of the yellow solid isabout 10 sec in alkaline solution, 8 min inwatee and 20 min in 3M sulphuric acid.
About half of the tetrafluoride is used ·inoxidizing the water and about half disproportionates. The overall reaction can therefore be represented by the equation:
3XeF4 + (6 + x)H 20= 2Xe + ~ 0 2 + Xe03 ,xH20+ l 2HF

E vaporation of the products of acid hydrolysis results in the isolation of xenon trioxide
as noted in the previous section (p. 181). In
the presence of barium carbonate, barium
xenate (Ba3Xe06 ) is formed but this rapidly
goes over to the perxenate, Ba2Xe06 .
In contrast to the variable products
isolated from simple hydrolysis experin1ents,
the reaction with aqueous potassium iodide
can be made quantitative w1der appropriate
condit ions, and this reaction has been used
for analysis of the tetrafluoride:
XeF4 + 4Klaq = Xe + 212 + 4KF

The liberated xenon was measured gasvolumetrically, the iodine determined by the

184

EDUCATION IN CHE""IISTRY

usual thiosulphate titration, and the fluoride
determined gravimetrically as calcium
fluoride.
Hydrolysis of xenon hexafluoride withsmall amounts of water yields the oxytetrafluoride (see previous section, p. 180). Normally, however, the hexafluoride oxidizes
water to oxygen and also disproportionates
to Xe VIII and lower oxidation states ofxenon. XeVIII is fairly stable in alkaline
solution but is unstable in acid, so that
normally Xevr is the highest oxidation
state in products of acid hydrolysis.
Perxenates can be prepared by titrationof a.n aqueous solution of xenon trioxidewith sodium hydroxide (0·5-18M); disproportionation occurs as indicated by thefollowing overall equa.t ion :2Xe03 + 4NaOH + 6H2 0= Xe + 0 2 + Na4 Xe06 ,8H20
Potassium hydroxide gives the yellow, explosive complex, K 4Xe06 ,2Xe03 .
Perxenates are more conveniently preparedby alkaline hydrolysis of xenon hexafluoride.
The white sodium and barium salts are themost stable and from these the less-stableyellow salts of Pb 2 + and U022 + have been
prepared metathetically. The least stable
perxenate yet studied is the black silver salt
which decomposes violently at 150°.
In common with the oxyacid salts of manyother heavy elements, various perxenates canbe obtained, the general formula being
M2nXe04+mYH20. Thus, salts of formulae
Na4Xe0 6 ,2H20, Na4 Xe0 6 ,8H20 , Na5HXe0 7(i.e. n = 2}, y = }) and Na6Xe0 7 ,H20 havebeen isolated. The integrity of several oft hese salts has been established by X-ray
powder photography:
FonnulaStTuctun~ab

Na4Xe06 ,2H20 orthorhombic 6·25
5·77
Na4Xe06 ,8H20 orthorhombic 10·36 10·45
cubic
a 0 =9·391
Na5HXe0 7

c

10·28
11·87

The infra-red spectra of solid perxenates
indicate that the v3 (antisymmetl'ic stretching) mode of the Xe0 6 group absorbs at
650- 680 cm- 1 . This is close to the value ofthe same mode in antimonates, tellurates andparaperiodates and indicates that the Xe- 0
bond in the Xe06 complex is similar to the
Sb- 0 , Te- O and I-0 bonds in the octahedralgroupings Sb06 , Te06 and I06 .
The solubility of the dry sodium salt,
Na4 Xe06 , in water at 25° is 7 gj l and in
0·5M sodium hydroxide 0·2 gj l. It is, however, pethaps a little early to anticipate thegeneral use of perxenate as a gravimetricreagent for sodium!
As already indicated, solutions containing
Xevr are strongly oxidizing.
Iodide,bromide and chloride are all rapidly oxidized
to the element, and iodine and bromine arethen slowly oxidized to iodate and bromate.
Manganese(n) is slowly converted to manganate and then permanganate. F e11 is
rapidly oxidized to Fem and Hg0 to
Hgn. As expected, perxenates, containing
XevrH, are even stronger oxidants. They
oxidize more rapidly and take iodate toperiodate ; water is oxidized slowly in alkaline
solution but rapidly in acid solution. Theseand simila.r observations enable the following
approximate redox tables to be constructed:
Acid

Xe _ _:.. :I·S::..:Y_ _

Base

0 .~
9 Yc___ HXeO,Xe ---"-

S~'ABILITY

X eO,--')'-'·O' -V_ _ H Xe0 6
4- - "0-·9.:..:Y__ HXe061 -AND NATURE OF BO~DIKG

Thermochemist1·y
The difluoride, tetrafluoride and hexafluoride of xenon are thermodynamicallystable and are formed exothermically fromtheir elements. The bonds in these compounds are chemical bonds with an energy
content well within the normal range. It is
known that the mean bond-energy E(Te-F)
in TeF6 is 78 kcaljmole and E(I-F) in IF5
is 63 kcalj mole; extrapolation to the isoelectronic molecule XeF4 might therefore lead
one to expect E(Xe-F) in the range 20-40
kcalj mole, and this is observed.
The standar d heat of formation of crystalline xenon tetrafluoride from gaseous xenon
and fluorine has been determined by studying
the reaction of xenon tetrafluoride ·withaqueous potassium iodide in a rocking bombcalorimeter and analysing the contents after

CHEMISTRY OF THE NOBLE GASES

the reaction.
tions gave:

The mean of four determina-- L':,.H'f(XeF,1 cryst) = 60·1 kcaljmole

As the heat of sublimation of the tetrafluoride
is 15·3 kcalj mole and the heat of dissociationof fluorine is 36·7 kcal j mole, then
- L':,.H[(XeF1 gas) = 44·8 kcalj moleand E(Xe- F) in XeF4 = 29·6 kcalj mole
The relevan t thermochemical cycle is set out
below (heat evolved is positive, heat absorbed
is negative).
Xe(g) + 2F,(g)0 kcal_ _-+--=6-"-0·_.1_.k~c_al_ _- 73·4 kcal

Xe(g) + 4F(g)XeF.(c)-15·3kcal- -+----:-1-;-;18::·2-c:;-;-k-:ca71-----<- XeF,(g)A second method, involving the reductionof xenon tetrafluoride with hydr ogen at 130oin an isothermal calorimeter, gave a preliminary value for the heat of reduction of 202 kcal.
This corresponds to - L':,.H~ (XeF4 gas)= 54·8
kcalj mole and a mean bond-energy E(Xe- F)= 32·0 kcaljmole. This is in reasonable
aareementwith the independent value derived
0
above.
The specific heat and heat capacity ofxenon tetrafluoride have been measured
between - 263° and + 27° and no anomalies
were detected. At 25° the specific heat is0·1367 caljg, and t h e heat capacity,
C :v = 28·334 calf mole. The measurements
lead to the standard entropy, S 298 (XeF~ cryst)= 35·0 ± 1 cal deg-1 mole- 1 , and hence toan entropy of formation, t:,.So = 102·5 caldeg- 1 mole-1 . The free energy of format~on
can then be obtained from the relatwn- L':,.G~ 98 = - L':,.H + T L':,.S = 29·4 kcaljmol~,indicating the considerable thermodynamiCstability of t h e tetrafluoride.
Reduction of xenon hexaflu oride by hydrogen un der similar conditions to those u~ed
for the tetrafluoride leads to the followmg
results:- ,L':,.H(reduction) = 306·5 kcalj mole
- L':,.H1(XeF 6 gas) = 78·7 kcalj moleE(Xe--F) in XeF6 =The fact

31·5 kcaljmolethat xen on hexafluoride is a

185

'normal,' unexceptional compound is furt heremphasized by comparing its mean bondenergy with that of its neighbours in t he
Periodic Table. 5 Thus, if t he mean bondenergies in SbF3 (92 kcal), TeF4 (86 kcal)
and I F 5 (63 kcal) are plotted against atomic
number and extrapolated by one unit to
XeF6 , the 'predicted' bond energy is 31 kcal,
close to t he value obser ved.
The thermodynamic data for gaseous xenon
difluoride at 1 atm and 250° have been calculated from the vibrational spectroscopic data
given in Fig. 2. The free energy (- L':,.G) andentropy (- L':,.S) of formation are :- f\,.G523 (XeF2 gas at 1 atm) = 5·510 kcaljmole- ,L':,.S523 (XeF2 gas at 1 atm) =26·96 cal deg-1 mole- 1
From the relation- L':,.G = - L':,.H + TL':,.S,
the heat of formation at 250° is- L':,.H,23 (XeF2 gas at 1 atm) = 19·6 kcaljmole
The value at 25° will be similar, and fromthis and the heat of dissociation of fluorine
one obtains
E(Xe--F) in XeF 2 = 28·2 kcalj mole
The similarity of the Xe- F bond energy in
XeF2 (28·2 kcal), XeF4 (29·6-32 kcal) andXeF6 (31·5 kcal) is noteworthy, as also isthe apparent slight increase in bond energy
with co-ordination number. This latter trend
may well be within the experimental errorof these diverse determinations (probab ly
about ± 2 kcal) but if substantiated isunusual, since increase in co-ordination
number of an element almost invariably
diminishes the mean bond energy.
Oxy-compounds of xenon are less stablewith respect to formation from the elements
than are the fluorides but this reflects t hegreater stability of the doubly-bonded oxygen
molecule compared with fluorine rather than
the wealmess of the Xe-0 bond. (The bonddissociation energy of oxygen is 118·5 andof fluorine 36·7 kca.lj mole.)
From the heat of explosion of xenon tri oxide as measured in a bomb calorimeter,
the heat of formation was found to be :- ,L':,.H'} (Xe03 cryst) = -96 ± 2 kcalj mole
Exact calculation of the mean bond energy··'E(Xe-0) would require a knowledge of the

186

EDUCATION IN CHEl\ITSTRY

heat of sublimation, as shown by the following
thermochemical cycle:
Xe(g) + i O,(g)

""'

j-

- 96 kcal

XeO,( c)j'

177·8kalXe(g) + 30(g) (

H• .,

.

_ Ll H )XeO,(g)

81 8

""

Hence E(Xe- 0) in Xe03 = 27·3 - t Llfisubl
The heat of sublimation is not measurable
but if a plausible value of about 22 kcal istaken for this quantity thenE(Xe-0) in Xe03 ~ 20 kcal
This is somewhat smaller than t he Xe- F
bond energy but would not, by itself, confer
explosive instability ; the large endothermicheat of formation is due to the fact that theXe-0 bond energy is very much less t han
the 0 =0 bond energy in molecular oxygen.
Extrapolations similar to those alreadymentioned suggest that the heats of formation of xenon chlorides, like that of xenon
trioxide, would be negative. Likewise theheat of formation of krypton tetrafluoride
from its elements should be approximately
zero, but the difluoride is probably slightly
exothermic.
Theories of Bonding
A satisfactory theory of bonding must
explain convincingly not only the stoichiometry and shap e of the noble-gas compounds
but also deal adequately with the observed
physical properties and bond energies . As
the theory becomes more precise, confident
predictions of the stable groupings of atoms
should be forthcoming.
It is clear that van der W'aals interactions
are far too weak to explain the bonding innoble-gas compounds. Such interactions arenormally only l - 5 k caljmole whereas theobserved bond energies are 20- 30 kcalj mole.
Nor does the assumption of ionic bonding
adequately explain t he energies involved.
Even for xenon difluoride a simple calculationshows that the second-stage ionization potential is too large to confer stability on t hestructure Xe2 +(F-) 2 and the similarity inmean bond energies and bond distances forXe] 2 , XeF4 and XeF6 rules out the poss~­bility of structures such as Xe4+(F- )4 andXeG+(F-) 6 . Other lines of evidence aga.i nst
ionic bonding in these compounds are theobserved crystal structures of XeF2 and XeF4 ,the volatility of the compounds and theabsolute values of the bond distances, which
are at least 0·3 A shorter than would be
predicted on the basis of the packing of ions.
The bonding in xenon difluoride can formally be considered in terms of the promotionof a 5p electron on xenon to a 5d level.
However, the energies involved suggest that.
this is unlikely to occur to any appreciable
extent, and calculations indicate that the
5d orbitals are rather too diffuse for goodbonding anyway. It follows that further
promotion to account for the tetra- and h exa fluoride is even less favourable en ergetically .
The valence-bond method then suggests the
mixing in of other structures with theseh ybrids to improve the bond energy, e.g.
F - Xe+F - and F - Xe+-F, but this is notparticularly satisfactory and in any event
merely describes t he ground state of themolecule and says nothing about its electronically excited states and hence its ultraviolet spectrum. 9 It should be noted, however, that this theory correctly predicts a
linear molecule for the difluoride and asquare-pla-n ar molecule for the tetrafluoride.
It predicts a tetragonally distorted octahedron for the hexafluoride, and carefulmeasurements are at present b eing undertaken to determine whether the structure ofthis compound is in fact a regular or an
elongated octahedron.
The most satisfactory treatment of thebonding in xenon compounds is the molecularorbital method. In this th eor y molecularorbitals a.re constructed from atomic orbitals
on the xenon and fluorine atoms. The threecriteria which determine whether a given set
of atomic orbitals can combine to give
molecular orbitals are :(a) t he atomic orbitals should have similarenergies;(b) they should have appreciable spatial
overlap;(c) t hey should have appropriate symmetry.

187

CHEMI STRY OE' THE ~OBLE CASES}lolecular orbitals which have energies below
the mean energy of the atomic orbitals fromwhich they are constructed are called bonding
molecular orbitals. Molecular orbitals whichha\e energies close to those of the constit uentatomic orbitals are called non- bonding molecular orbitals and those having energies above
those of the constituent atomic orbitals arecalled anti-bonding molecular orbitals.
The theory can be illustrated by reference
to xenon difluoride. The outer electronic
configurations of xenon and fluorine arc:
Px

P

Pz

00

kXM><J
Xenon

Fluorine

Let the molecular axis be the z axis. Then
three atomic orbitals must be considered:the 5p. orbital on the xenon atom andthe 2p. orbital on each of the fluorine atoms.
These three atomic orbitals can combine togive three molecular orbitals, one of which
is bonding, one non-bonding and one ant ibonding. The three atomic orbitals originallycontained four electrons in all, i.e. two in the
xenon 5p. and one each in the two fluorine
2p. orbitals. These four electrons are now
placed so that t\m occupy t he bonding
molecula1' orbi tal and two occupy the nonbonding molecular orbital. The situation is
illustrated in the following diagram:
F,

Xe

F,

Descriptiont EX±) EX±) EX£)..... EX±)(fB

Occupancy

anti-bonding

>..

1>0

non-bondingH

bondingttc:w

I EX±) (fB EX±)~~===- · ~.-~-:--.""::..· ~-~

1t is clear that for best overlap the threeatoms should lie in a straight line, as observed .
Notice that t he non-bonding orbital has nocontribution from t he xenon 2pz orbital, since
this has the wrong symmetry for combination,
as indicated by theand - signs. The.
approximate wave-mechanical description ofthese orbitals is:+

P .b = }Y,(Fl)1f'nb =
lJfb

lv'2,P(F1)= !<f(Fl)+lv' 2 Y,(Xe)+ t.p(F2)+OY,(Xe)- .y 2 .P(F2)l+ itj.(F2 )~Y,(Xe )' ~1In these equations, b, nb and ab refer tobonding, non-bonding and antibonding molecular orbitals, lJ', and the atomic orbitals, !{;,
are labelled with an obvious notation. Bysquaring and adding the coefficients in eachveetical and horizontal column it is simple
to check that each atomic orbital has beenused in all just once and that each molecularorbital is normalized to unity.
If the squares of the coefficients are taken
to indicate the extent of electron occupancy,a.nd if it is r emembered that 11\ and P nb eachcontain two electrons, then it becomes2
apparent that the xenon atom has 2()2 ) = 1
p .-electron in its neighbourhood whereas thefree atom had two; hence it has a resultant
net change of 1. Likewise, each fluorine

+

+

2

atom is associated with [2(!)2 2(J2 ) ] = 1%
p.-electrons, so that it has a resultant chargeof - 1/ 2. The theory thus enables us tocalculate approximately the charge distribut ion within t he molecule. A more precise
calculation leads to a charge of 1·4 on the
xenon and -0·7 on each fluorine. As aresult of this partia.l separation of charges"ithin the molecule. the heat of sublimationof xenon difluoride contains a su bstantial
coulombic component in addition to theubiquitous van der 'iVaals attraction, and istherefore considerably greater than the heatof sublimation of xenon itself, as indjca.t ed
in Tables I and II. I ndeed, an estimate of
the heat of sublimation based on an arrayof atoms in the positions indicated in F ig. 1
a nd carrying the partial charges calculated
above gives excellent agreement with theexperimental value.
An important consequence of the molecular-orbital bond theory is that it predicts
enhanced stability the lower t he ionization
potential of the cent ral atom and t he more
electronegative the outer atoms. Compoundsof xenon and radon with fluorine and oxygen
are therefore expected to be the most stable,
whereas krypton fluorides or radon chloride
should have only marginal stability. Compounds of argon and neon are not expected.
The molecular-orbital themy outlined
above is in no way unique to inert-gas compounds or unusual in any of its features. It
had previously been used to d iscuss bonding

+

188

EDUCA'l'ION IN CHEMISTRY

in t he polyhalide ions, such as r3- , 1012- ,1014 - , etc., and in many other situations. It
can be extended to deal with compounds
such as xenon tetrafluoride and xenon hexafluoride simply by forming three-centre
molecular orbitals in the x and y directions
as well as in the z direction. The theory
therefore predicts a square-planar shape for
XeF4 and a regular octahedral shape for
XeF6 . The similarities in bond lengt hs andbond energies for these compounds can also
be understood. P artial charges on xenon
and fluorine are approximately + 2 and - t
in XeF4 ; and +2 and - t in XeF6 .
A final feature of the theory should be
mentioned- its ability to help inter pret
ultra-violet and visible absorption spectra.
Thus when an electron is excited from lf'"'~to lf'a~ it will absorb energy at a frequency
defined by v = 6.E / h where h is P lanck's
constant and 6.E is the difference in energy
between the two levels. l\'Iore extensive
calculations also consider the molecularorbitals of 7T-symmetry, 9 and a detailed
analysis of the various energy levels is atpresen t being carried out for t he complicatedcase of xenon h exafluoride which, as mentioned earlier, is thermochroic.

forced us to thin~ more deeply about t he
fundamen tals of chemical bonding and has
demonstrated in the most direct way possible
how inadequate any 'stable octet' t heory ofvalency must be. The stable octet, which
was the first electron ic theory of valency ,was a great advance when it was originally
expounded 50 years ago. It is distu rb ing
to reflect th a t much elementary teaching ofbond theory still relies on this concept
despite the large areas of chemistry where it
does n ot apply (e.g. ferrous and ferric compounds) and the many advances which havebeen made in our thinking during the last
40 years. An important lesson can also belearned from the fact that if, in 1962, a
student had been asked 'What would happen
if xenon were heated in the p resence offluorine?' and he h ad answered 'a volatile
white crystalline solid of composition XeF4
is formed quantita tively,' he would h ave
been awarded no marks.
REFERENCES

I.

2.
CONCLUSION

Speculations on potential applications ofnoble-gas compounds have been numerous.
Are these compounds a mean s of storing
fluorine in a readily available form; can theybe used as oxidants in rocket-propellant
systems or as specific fluorinating reagentsin chemical reactions ? The use of per xenatesas powerful aqueous oxidants h as alreadybeen mentioned. Specialized uses for shocksensitive oxides su ch as xenon trioxide might
be found. One of the m ost objectionableatomic-fission products is the volatile xen on
isotope, 133Xe; conversion to an in volatilecompound could make t his available inmegacurie amounts. Involatile radon needles
for tumour therapy is a noth er possibility .
Whatever t he future applications of t hese
compoun ds on e thing is certain: the discovery
t hat xenon readily forms a wide ra nge of
stable compoun ds h as proved to be on e oft he most stimulating and significant observ ations in ch emistry for many years . It has

3.

4.

5.

6.

7.

8.
9.

C. L. Chernick, ' Chemical compounds of then oble gases,' R ec. chem . Progr., 1963, 24,
139- 155.
(A review of both t he early
histor y and the recen t chemistry of th eseelements conta ining 123 refer ences.)
N . Bartlett, ' X enon h exafluoroplatinate (v)Xe+[PtF6 ] - , ' Proc. chem . Soc., 1962, 218.
(The fi rst conclusive evid ence for the formation of a chemical compound invol vin g a
noble gas.)
Noble Gas Compounds, ed. H. H. H yman,
Unive rsity of Chicago Press, 1963. (Theproceedings of a conference h eld a t th e
A r·gonne National Laborator y, Chica go in
April, 1963.)
H. H. Claasen , H. Selig and J . G . :.VIalm, J.
Amer. chem. Soc., 1963, 84, 3593. (The firstreport of the direct oxidation of x enon byelemental fluor ine.)
N . Bartlett, 'The chemist ry of the no ble gases,'Endeavour, Jan., 1964, 23, 3- 7. (A non technical a ccount, w·hich includes colou r
photographs of P tF 6 , X eP t F 6 and XeF4 . )G. J. :Moody and J. D. R . Thomas, ' Compoundsof the Noble Gases,' J . roy. I nst. Chem.,
1964, 88, 31-33. (A general rev iew with61 r eferences.)
H. Selig, H. H. Claasen , C. L . Chernick, J . G.
:Maim and J. L. Huston, 'Xen on tet roxide :p reparat ion and some p ropert ies,' Science,
1964, 143, 1322.
A . G. Strong a nd A. V . Grosse, ' Acid of krypton
an d i ts ba ri.um salt,' S cience, 1963, 143,
242- 3.
C. A. Coulson, 'The n ature of the bondin g inxenon fluorides and related molecules,' J .
chem. Soc., 1964, 1442-1454. (An ex cellen t
assessment of the various t heories of bonding.)