1st NORANDA LECTURE

Unusual Oxidation States of the Noble Elements
NEIL BARTLETT, M.C.I.C.,
University of British Columbia,
Vancouver, B.C.

the past six years the fluorine chemistry of thenoble metals has been revised and considerably extended. Weinstock and his co-workers at the Argonne
ational Laboratory have contributed significantly to thechanges (1 ) . They have reported four new hexafluorides(PtF,; <2 l, TcF 6 ( 3 ), RuF G(4 ), RhF 0 <5 )) and have shown (6 )the fluoride previously considered to be osmium octafluoride to be the hexafluoride. In addition, rhenium heptafluoride has been prepared and characterized <7 ).
The known highest oxides· and !}uorides of the second
and third transition series elements are listed in Table 1.
An interesting feature of this comparison is that althoughrhodium hexafluoride exists the trioxide is unknown, but
the highest fluoride of ruthenium is· the hexafluoride despite
the existence of a tetroxide. Osmium is similarly restricted
to a hexafluoride despite the existence of a tetroxide. The
absence of an octafluoride or at least a heptafluoride ofosmium is all the more surprising in view of the existenceof rhenium heptafluoride. The nature of the compounds
produced with varying proportions of oxygen and fluorine,
as shown in Figure 1, is revealing. The attainable formaloxidation state decreases as the proportion of fluorine in
the oxidizing gases increases. In all, the trends point to thehexafluoride being the highest fluoride attainable. We may
suppose that this limit is imposed by a maximum coordination of six for the osmium but this would not exclude
osmium dioxide tetrafluoride, of which no trace has yet
been found. If, on the other hand, the real oxidation statein osmium tetroxide is the same as that in osmium hexafluoride then the electron withdrawing power of theoxygen ligand would be set at 1.5 times that of the fluorine.
This would result in osmium in osmium dioxide tetrafluoride being in the same oxidation state as in osmium
heptafluoride, which is also unknown.
Osmium oxide pentafluoride <8), which as is shown in
Table 2, resembles the related rhenium compound <9 ), is ofespecial .interest in that it is the :first simple osmium compound in which the formal oxidation state is + 7. Osmiumin this oxidation state possesses a d 1 electron configurationand in accord with this the compound i paramagnetic.
Since the oxygen ligand is similar in size to fluorine andince the non-bonding electron should show no stericactivity the molecule is expected to be similar in size andshape to osmium hexafluoride. In fact the oxyfluoride andhexafluoride both crystallize in the orthorhombic system,
the unit cells being similar in size. The molecular shape ofthe oxyfluoride, as determined by X-ray single crystal
methods is shown in Figure 2. The distortions from regularoctahedral shape can be interpreted in terms of distortions
due to the slightly larger ize of the oxygen ligand, superimposed upon a trigonal compression of the molecule. The
latter di tortion may arise from lattice interactions.
The iodine and technetium analogues of osmium oxidepentafluoride presumably can exist. Even ruthenium oxide

W

ITHl

pentafluoride may be preparable although the thermal
stability may be low. A higher fluoride of ruthenium thanthe hexafluoride appears to be unlikely in view of thethermal instability of this fluoride.
The data for the second and third transition series hexafluorides given in Table 3, show the fusion point andboiling point rising w ith increasing atomic number alongeach series as the effective molecular volume decreases.
These trends indicate a significant increase in the intermolecular attraction with atomic number in each series.
Thermal stability falls with atomic number in each seriesalso and there is apparently a correlation between thermal
stability and effective molecular volume, platinum hexafluoride being comparable, in stability and molecularvo lume, with ruthenium hexafluoride, and rhodium hexafluoride with the smallest molecular volume being the least
thermally stable hexafluoride and the most difficult one to
make. Despite some effort in this direction, it has notproved possibJe to prepare palladium hexafluoride ( .5) •In the first reports of platinum <2 ), ruthenium <4 ) andrhod ium (fi) hexafluorides their spontaneous· liberation offluorine at ordinary temperatures was mentioned. Clearly
TABLE 1. -HIGHEST F L UORIDES A rD OXIDES OF THE 2ND AND
3RD TRA SITION SERIES

Fluorides
Oxide

MoF6
MoO a

T c20 1

RuF 6
RuO 4

RhF6
RhO 2

PdF 3
Pd 2O 3AgF 2
AgO

Fluoride
Ox ides

WF6
WO 3

R eF1
R e201

OsF 6
OsO 4

IrF6
IrO 3

PtFG
PtO 3

Au F 3
Au2O 3

TABLE 2. -TcFn

COMPARISO r OF MOF s AND MF 5 COMPOUNDS

R eF 5< 1 l

R eOF s 1bl

OsF5 <' l

OsOF.,< lm .p. (°C)

18 . 734 . 534 . 459 . 8b.p.

33 . 855

47 . 5,..__,90 . 0106 . 6(or t h )<d)-

105 . 7(orth )<dl

103 . 8( orth) < l

0Eff. mo!. vol. (A 3)( a) Cady and Hargreaves, J. Chem. Soc., 1961, 1565.
( b) Aynsley, P eacock and Robinson, Ibid, 1950, 1623.
( c) Bartlett, Jha and Trotter, Pro·c. Chem. Soc., 1962, 277.
( d) Siegel and Northrop, P erson al communication.

----t OsF fj <al
0 + 3F 2
Os + 2 1/2F 2 + 1 /20 2--) OsOF s<bl----t OsO 3F 2 ( m .p . 172 b.p . 200°)(c)Os + F 2 + 3/2O a--) OsO 4 (m.p. 40.1 b.p. 129°)<ct>
Os+ 40 2( a) W einstock and Malm , J. Amer. Chem. Soc., 1958, 80 , 4466.
( b) Baxtlett, Jha and T1·otter, Proc. Chem. Soc., 1962, 277.
( c) Hepwo1ih and Robinson , J. lnorg. Nu.clear CT1em., 1957, 4, 24.
( d) Wartenb erg, Ann. 1924, 440, 97.

Figure I-Oxides, Oxyfluo1·ide and Fluoride of Osmium.

Chemistry in Canada, August, 1963
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33

NORANDA LECTURE

TABLE

any species which will liberate fluorine is a powerful oxidizing agent.Their oxidizing potential, however, remainedlargely unexplored.

=

Orthorhombic unit cell ( 23 ° ± 2) : a
9.540 ± 0.004,
b
8.669 ± 0.004, c
5.019 ± 0.004 A:; V
415.1 .A3 ;Z
4. Space ,g roup Pnma (D;1~, No. 62).

=

=

=

=

3. -

2ND AND 3RD TRANSITION SERIES HEXAFLUORIDESI MoFG(,) T cF s(b ,cI)m.p. (°C )va p. press. 15°C.
~H rusC kca l.moie- 1 )~H vap( kca l. mol~- 1 )E ff. mo!. vol. (A 3 ) Ce>

I'

,

'

I

'

I

•

'

I

'

Os-O
Os-F1
Os-F2
Os-F3
O-F!!
O-Fl
F 2-F18
F2-F/
F a-F0'F1-F2
F i-Fa

I

/

1.74
1.73
1.82
1.75
2.44
2.61
2.52
2.51
2.51
2.52
2.36

87 °

O-Os-F2
O-Os-F3
E2-Os-F3
F 2 -Os-F/
F 3 -Os-F3'
F 1 -Os-F2
F 1 -Os-F3

97
90
87

91
86

ca> Bartlett and Trotter, Unpublished results.

t2g
orbitals

Molecular structure of osmium oxide pentafluoride OsOF6 ca>.

•·~

eg

orbitals

do
MoF6
WF0
NbF6 TaF0 -

z

-O&x

d1
TeF6
ReF8
MoF6WF6 -

d2

9

RuF0
OsF6
TcF6ReF 6-

da
RhF6
IrF0
RuF0 OsF6-

-1
+

~x
+

't-

4
--Or~·
d,I

ds

PtF6
RhF6lrF6-

PtF6-

de

Figure 3-d-electron configurations for 2nd and 3rd trans.
series hexafluorides.

The dissociation of some of the hexafluoride molecules
implies that in these molcules the noble metal atom has asufficiently high electron attracting power to withdraw abonding electron from a fluorine ligand. That this tendency should increase with atomic number in each series
may be seen by reference to the non-bonding valence electron configurations for the various hexafluorides, which are
given in Figure 3. Tungsten and molybdenum use all of
their valence electrons in the hexafluorides but with technetium and rhenium and the later elements, non-binding delectrons remain and are located in the lower lying t 2 gorbitals, in w hich situation they do not effectively screen
the ligands from the nuclear charge, which increases withatomic number . The fluorine ligands in platinum hexafluoride, w ith a d\ 2 g configuration for the platinum, areevidently exposed to a strongly electron attracting central
atom. Since gold hexafluoride would be a d 5 Lzg system thenthe polarizing power of the gold would be greater thanthat of platinum in platinum hexafluoride and in conse-

34

101 . 8

70
38

-

-

-

101 ,099 . 533,418 , 7m.p. (°C )2.0
43 . 8b.p. (°C )17 . 147 . 533 . 853 . 61~Hrus(kca l.moJe- )0 . 42
1 . 1071.76
1.19
6 . 336 . 867 6 , 72
7 . 38~H vap(kcal.mol~- 1 )Eff. moi. vol. (A 3 )Cel 108 . 5 106 , 6 105 . 7 105 .4

61.3
69 . 141.08
7 . 06
104 .6( a) Cady and Hargreaves, J. Chem. Soc., 1961, 1565.
( b) Ch ernick, Claassen and W einstock, J. Amer. Chem. Soc., 1961, 83,3165.
( c) W einstock, Malm and W eaver, Ibid, 1961 , 83, 4,310 .
( d) Selig, Chernick and Malm, J. Ino-rg . Nuclear. Chem. , I 961 , 19, 377.
( e ) Siegel and Northrop, Personal Communication.
TABLE 4. -

,

~

-

54
58

93

' i~
Figure 2 -

37,0
135

RhFs(b)

WF6<a ) R eF 6(a ) OsF 6(a ) IrF 6caJ P t F 6<c)

A'.
. .o.
,, ·,

17 .4
361
0 .92
6 .94
106 , 6

RuFs(b)A NALYTICAL DATA FOR PT0 2 Fs

Found

Pyrohydrolysis: (solid + H 20 va por - followed bysteam - HF
(residue ignited in H 2

F, 32 . 4 %Pt, 57.5

Sodium fusion: Fluoride precipitated as lead chlorofluoride

F, 32. 7

B romine trifluoride: Oxygen evolved measured withgas burettePt0 2F6 requires F , 33.4; 0, 9.4;
Pt, 57.2 %0, 10 . 4

quence the stability should be lower. Nevertheless this does
not mean that gold hexafluoride can not exist. It may be
no more unstable than rhodium hexafluoride.
Prior to the first repon of platinum hexafluoride C2 ), we
had observed that when platinum or platinum salts were
fluorinated in silica apparatus a deep red solid was produced. This solid sublimed in a good vacuum, below 100°.
It was evidently an oxyfluoride, the oxygen being derived
from interaction_of fluorine with the silica: Si0 2 + 2F2 ➔SiF 4 + 0 2 , since fluorination of platinum in nickel apparatus did not give it.
The analysis of this most interesting material cio), which
was largely the concern of my collaborator D. H. Lohmann, proved to be difficult. The solid reacted violently
with water or alkaline solution to give a mixture of products among which ozone, platinum, hydrated platinous
oxide and platinic oxide and the hexafluoroplatinate (IV)ion were identified. It is probable that oxygen difluoride
was also produced in such reactions since fluorine analyses
employing aqueous solution techniques always gave low
results. Only when the hydrolysis was carried out withwater vapor were consistent analytical data obtained. The
data are given in Table 4. They indicated that the redsolid was a compound of empirical formula Pt0 2 F 6 •Of course several formulations are possible for thecompound. Fortuntely the solid gave a simple X-ray
powder diffraction photograph which was indexed on thebasis of a cubic unit cell, a = 10.032.A. The compound wasapparently structurally similar to potassium hexafluoroantimonate (V) the structure of which had. been given
by Bode and Voss (11). Furthermore we found that nitrosy l hexafluoro-osmate (V) and nitrosyl hexafluoro-antimonate (V) were also isomorphous with the compound.
Figure 4 shows the close resemblance of the X-ray powder

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NORANDA LECTURE
photographs. Good agreement was· obtained between theline intensities calculated on the basis of the ionic stru cture, 0 2 + [PtF 6 J- , represented in Figure 5, and thoseobserved. The structure reveals that the PtF 6 octahedra
are compressed along the three fold rotation axis along
which the dioxygenyl ions lie. Each 0 2 + ion has 12
fluorine neighbours, six in a puckered ring and almost
coplanar with the ion, the other six in sets of three, one
above and one below the plane. Apart from the nonspherical nature of the cation, the structure is almost
identical with that proposed for potasium hexafluoroantimonate (V ) .
The other properties of the compound were consistent

KSbFe,

with th e dioxygenyl hexafluoroplatinate( V ) formulation.
In its chemical reactions w hich are show n in Figure 6, it
behaves as a powerful oxidizing agent and as a derivativeof platinum (V ) .
The magnetic behavior of the compound is unusual for
a complex of an element of the third transition series inthat the susce ptibility obeys the Curie Weiss Law, XM ex:
1/ (T + 0) , where 0+ 45 °. This and the magnitude of
the magnetic moment ( ueff 2. 57 B.M) are compatible withthe effect of the dz:! g configuration of the Pt 5 + and theunpaired electron associated with the dioxygenyl ion. Since
for the third transition series, the spin-orbit interaction isg reat, the magnetic moment associated with the d \g-configuration is almost temperature-independent <12 ) , over therange used for the magnetic stud y, and is 1.73 B.M. It isalso to be expected that the magnetic behavior of the
dioxygeny l ion will resemble that of nitric oxide withwhich it is isoelectronic. Magnetic moment values computed for a nitric oxide hexafluoroplatinate (V) ion
combination are similar to those observed for the salt.
Representative values are given in Table 5.

=

T AB L E

5. -S OM E CA LC U LATED AN D OB SER VE D MAGNETIC MOM ENTS F O R 0 2+[PtF5l -T emp . (° K !

Figure 4-X-ray powder photographs of some ABF 6
complexes.

A

113
2 .31

2 .40

2 . 56µ obs. ( B.M. )2 . 122 . 282.34

B

0 = FLUOAIN!.
'9 = PLATINUM
a.. = 1.1:3ib . = 1.74
,C,

A. UNIT CELL OF= 2 _qq

d. . = 2.57-

e = 3 .74

o; [Pt Fl (cub.)B. CONTENTS OF" CUBIC PSEUDO UNIT CELL

Figure 5.

300°µ calc. ( B .M. )0= OXYGEN

TM£ UNITS !II-IOWN IN · e • LIE: ALONG
Tl-IE THREEFOLD AXES INOICA'fED

178

It was clear that if the formulatjon
02 + [PtFt,J - were true, then platinum
hexafluorjde should oxidjze molecularoxygen. A sample of platinum hexafluoride was therefore prepared, andsince it has a vapor pressure of ,_, 80
mm. at room temperature, the interaction with molecular ox ygen was
followed tensimetrically . The deep red(bromine colored) platinum hexafluoride vapor reacted instantly withthe oxygen in a 1: 1 molar ratio togive the familiar red compound.
Interestingly, the solid made by thismethod at low temperatures, is isomorphous and almost isodimensionalwith rhombohedral potassium hexafl.uoroplatinate(V). At higher temperatures transition to the cubic phaseoccurs.
A simple lattice energy calculation
using Kapustinskiis equations (1 3 ) gives
H = ,. . , - 120 kcal. mole- 1 for theprocess: iPtFo] :;> + 0 i11. ➔ o;- [PtFo] ~. Since the firstionization of molecular oxygen is + 281 kcal. mole-1 therequired minimum energy for the electron affinity ofplatinum hexafluoride is - 161 kcal. mole-1 . This value, istwice that of atomic fluorine. Platinum hexafluoride wasevidently the most powerful oxidising agent known
hitherto.

Figure 6-Chemical reaction of 0 / [PtF0 J-.

The way in which the oxidizing power of the hexafluorides increases with atomic number in the third transition series is demonstrated by the nature of the reactions
with nitric oxide which are summarized in Figure 7. Since
the ionizatjon potential of nitric oxide 1 4 is 62 kcal. mole- 1
lower than for oxy gen it can reduce osmium hexafluoride
to the quinquevalent and platinum to the quadrivalentChemistry in Canada, August, 1963

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NORANDA LECTURE
state. Of the third tranS1t1on series hexafluorides, oxygen
will only reduce the platinum fluoride. It is clear howeverthat both rhodium hexafluoride and ruthenium hexafluoride should in view of their thermal instability be closein oxidizing power to platinum hexafluoride and should
also interact with oxygen.
It was quickiy realized that platinum hexafluoride being
capable of oxidizing molecular oxygen would probably also
be capable of oxidizing the heavier noble gas elements. In
particular xenon, the first ionization potential of which <14lis the same as for molecular oxygen and radon with a firstionization potential of 10. 7 e. v. were possibly within range.
Althqugh the lattice energy of a salt is inversely proportional to the sum of the ionic radii, in the case of thehexafluoroplatinates it is· rather insensitive to cationic radiichanges since the complex ion radius is large. A simple
calculation made for the oxidized xenon case showed thatthe lattice energy would be decreased by little more than
10 kcal. mole-1 from that of the dioxygenyl salt. The predicted interaction of xenon and platinum hexafluoride wasconfirmed in a simple and visually dramatic experiment. The
deep red platinum hexafluoride vapor of known pressure
was mixed, by breaking a glass diaphragm, with the same
volume of xenon with pressure greater than that of thehexafluoride. Combination to produce a yellow solid was
immediate at room temperature and the quantity of xenon
remaining was commensurate, within the limits of error,
with a combining ratio of 1: 1 <15 ) .
This experiment was quickly confirmed by the workers
at the Argonne National Laboratory 0 6 ) who went on to
study the interaction of xenon with ruthenium hexafluoride. At that time we were without the facilities for
preparing the second period hexafluorides. The Argonne
workers noted that roughly three molecules· of ruthenium
hexafluoride were consumed per atom of xenon and thata material resembling ruthenium pentafluoride was a product of the reaction. This suggested that this hexafluoride
was, at least in part, acting as a fluorinating agent and thata fluoride or fluorides of xenon must exist. This led to thediscovery of the simple fluorides of xenon.
Meanwhile at the University of British Columbia we
discovered that the xenon-platinum hexafluoride reaction
was not as simple as the .first experiments had indicated.
The reaction was repeated numerous times in a varietyof containers and never with a consistent s-eries of results.
Occasionally , a the results given in Table 6 indicate, a 1: 1
stoichiometry is obtained but usually between one and twomolecules of platinum hexafluoride are consumed per atomof xenon. Unfortunately the 1: 1 adduct does not give a
powder diffraction pattern and attempts to grow single
crystal of the material have met with no success. Since,
however, material of composition XePtF 6 and that approximating to Xe(PtF 6 ) 2 both interact with alkali fluorides in
the non oxidizing medium, iodine pentafluoride: XePtF 6 +RbF ~ RbPtF 6 ; Xe(PtF 6 ) 2 + 2CsF ➔ 2CsPtF 6 , to give
hexafluoroplatinates (V) it is evident that the platinum isin the quinquevalent state no matter what the compositionof the adduct. The most satisfactory explanation for theobservations at this time is that two reactions occur simultaneously the first forming the 1: 1 adduct presumably
Xe + [PtF 6 J- and the second producing platinum pentafluoride and a xenon fluoride : Xe + 2PtF 6 ➔ XeF 2 +2PtF ". The rather scanty spectroscopic evidence, shown in
Table 7, gives some support to these conclusions. Although

the xenon-platinum hexafluoride adduct decomposes at
160° to liberate xenon tetrafluoride, approximately half ofthe xenon remains to form a diamagnetic material of
empirical formula XePt:iF 10 •+ 219 kca l.mole- 1

NO --t NO+ + 1 e\:F s + NO no reaction <a)
ReF s + N O not studied\1

NO + [OsF5J -<b)

OsF s + NO --t

low(NO) 2Irfs(c)IrF6 + NO NO + [ lrFs ] -(b) 20°
temp
+NO
low
PtFs + NOte mp"0'(NO)2PtF6 (b)0 + [ PtFG] -(b) --t+NO(a) G eichmann, Smith, T rond and O gle, foorg . Ch em., 1962, 1 , 661.
( b) Bartlett and Jba.
(c) Robinson and Westla nd, J. Che m . Soc., 1956, 4481.

Figure 7-The reactions of third transition series hexafluorides with nitric oxide.
TABLE 6. -SOME RESULTS FOR THE XENON+ PLATI N M H EXA-F L UOR ID E REACTION

T ensimetric results for gas + gas reaction in a glass bulb
( 1)

(2)

95.0
156 .582 .0
1:1.27

Pressure of PtFs mm.
Initial pressure of X e mm.
Residua l pressure of Xe mm .
Combining ratio Xe: PtFG

70 .0
72 .0

23 .0
1:1. 42

Tensimetric results for PtF s (solid + gas) + Xe (gas) in a glass
bulb

Initial pressure of P t F smm.
Residua l pressure of P tF smm.
Initial pressure of X e mm.
R esidua l pressure of X e mm .
combinin g ratio Xe: P tF G

TABLE

I.R. (cm-

1)

( 1)

(2)

(3)

93 .0
0
108 .017 .0
1:1.02

56 .0
12 .0
27 .5
0
1 :1 . 60

117 .5
0
117 .559 .0
1:2 .07+

KPtF6

O2[PtF 5J-

X e( PtF ~)640 (v.s)
590 (s)6.31 (v.s)

652 (v.s)
550 (s)

3500

3825

0

vis. and U . V. (A)-TABLE 8. - X E ON- RHODIUM HEXAFLUORIDERhF 6 (solid + gas) + X e (gas) in a nickel weighing ca n

(1)

(2)

Wt. of xenon co nsumed g.

0 .0050

0 .0120

Wt. of adduct0.01390 .0337

Combining ratio X e :RhFs

1 :1 . 05

1 :1.10

Rhodium hexafluoride also reacts with xenon and thereacting ratio was, in the two experiments performed, close
to 1: 1 (Table 8). Presumably the adduct is Xe + fRhFf\Jbut again no X-ray patterns were ecured. Since rhodium
hexafluoride was considered, because of its small size andlow thermal stability to be the most powerful oxidizer ofChemistry in Canada, August, 1963

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TABLE 10. -TABLE 9. -AVERA.GE BO JD E TERGY (IN KCAL) FOR SOMEFLUORIDES

SbF 5
90 . 0 (a)T eF s
78 .5 (b )IF1
54 . O (c )

SbF3

TeF 4
88 . 1 (a )IF5
62 . J (d)[XeFs][24] (f)S TRUCTURAL DATA FOR XE ' ON FLUORIDE S

X e non Oifluoride
Crystal Stru cture: Body centred tetragona l ce ll , Z = 2, S.G.
14/ m mm
S. Siegel and E . Gebert, J. Amer. Chem. Soc., 85, 240 ( 1963 ).
2. H. A. Levy and P. A. Agron, J. Amer. Chem. Soc . 85 241( 1963 ).
''Linear molecule
X e-F
2.14 ± 0.14 A (X-ray)

1.

{1.

2. 2. 00 ± 0.01 A (neu tron)
1Q5 (b)

XeFs
31. 5 ce i
XeF4
32. 0 (elXe cg) + ½nF 2cgl - XeF 2
Exothermic whe n A.B.E. > D.H f(F) i.e. 18.3 kca l.

( a) Nichol , Ph.D. Thesis, University of Durham 1958.
( b) R ossini, Wagman, Evans, Lavine and Joffe, Natl. Bur. Standards(U .S .) Circ. 500 ( 1952).
( c) Wiebenger, H avinga and Boswijk, "Ad vances in Inorganic Chemistry
and Radiochemistry", Ed. Emeleus and Sharpe, Vol. 3, Academic
Press Inc., N.Y., 1961, p. 136.
( d) Woolf, J. Chem. Soc ., 1951, 231.
( e) Stein and Plurien, Conference on Noble Gas Compounds, Argonne
at!. Lab., April 1963.
(£) Estimated.

the known hexafluorides it was mixed with gaseous kry pton in an attempt to oxidise this gas. The red vapor of thehexafluoride persisted in the presence of a considerable
excess of the noble gas. Platinum hexafluoride display ed asimilar inability to oxidize krypton.
Within a few weeks of the discovery O 7) of xenon
tetrafluoride by Claassen, Selig and Malm, they and theircollaborators at the Argonne National Laboratory, had also
prepared and characterized a difluoride ( IS, 10 ) and a hexafluoride (20 ). These compounds were soon confirmed in a
number of other laboratories. All thee fluorides, w hich arewhite crystalline solids at room temperature are preparablefrom fluorine and xenon, the nature of the product depending upon the ratio of the reactants and the temperature, the tetrafluoride being thermally more stable thanthe hexafluoride (2 1).
The fluorides excited great interest among theoreticalchemists and some have undoubtedly felt that an unorthodox form of bonding was necessary. The bonding appears,
however, to be little different from that encountered in
the halogen fluorides and interhalogens. Average bondenergies for the fluorides of the elements antimony toxenon are given in Table 9. They include the recently
reported values for the xenon fluorides <22 ) . The smooth
reduction in bond strength from antimony to xenon ispresumably , in part at least, a consequence of the increasein the coordination number of the central atom. (Theaverage bond energies are derived from the heat of formation from the gaseous atoms and not from the atoms intheir valence states. ) It is of interest that extrapolation of
the average bond energies for the highest fluorides of antimony, tellurium and iodine, shown in Figure 8, gives anaverage bond energy for xenon octafluoride of 24 kcal.,indicating a possible existence for this compound.
01; the . basis of simple .electron_ pair repulsion theory
the d1fluonde should be lmear, with the fluorine atoms
at the poles of a trigonal bipyramidal arrangement ofelectron pairs. The tetrafluoride should be a square planarar~angement of fluorine atoms about the xenon atom, there
bemg a non-bonding electron pair both above and below
the I?lane. The he~afluoride, since it possesses a nonbondmg electron pair, should show some deviation fromoctahedral symmetry. The structural evidence has so far
been _in accord with all of these expectations. Table 10
contams some _of th~ structu~al data obtained from X-ray
and neutron diffract10n studies of the di (2 3, 24) and tetra-

38

Xenon Tetrafluorid e
Crysta l Structure: Monoclinic unit cell , z = 2, S.G. P2 i/n
1. S. Siegel and E. Gebert, J. Amer. Chem. Soc., 85, 240 ( 1963 ).
2. J. A. Ibers and W. C. H ami lto n, Science, 139 , 106 ( 1963 ).
3. D. H. Templeton, A. Zalkin, J. D. F orreste r and S. M. Williamson, J. Amer. Chem. Soc., 85 , 242 (1963).
4. J. H. Burns, P. A. Agron and H . A. Levy, Science 139 1209
( 1963 ).
''Molecular symmetry 0 4h:
2 · 1 _92 ± 0 _03 A
X e-F

1

3. 1 . 93 ± 0 . 03 A
1 .951 ± 0 . 002 A
4. 1.954 ± 0 . 002 'A

Molecular species:
Infra red and Ram.a n spectra
H. H. Cl aassen, C. L. Ch ernick and J. G. Malm, J. A mer. Ch em.
Soc., In Press.
Symmetry D 4h· Bond length 1.85 ± 0.2 A (IR )Maximum pu ckering from sq . planar config. 0.5° or 0.02 A
Shift for F from plane.

fluorides <23 , 25 , 26 ) . The neutron diffraction data is themore precise. The infra-red spectra of XeF 2 vapor (27) andXeF 4 vapor (28 ) are consistent with the linear and squareplanar models.
The preparation of the xenon fluorides initiated workon the other noble gases. Radon which would be expected
to form compounds more readily than xenon was quickly
s~o~n <~ 9 ) to for1:1 a fluo_ride but at this time its composit~on 1~ st1_ll uncertam. Stud1es with this noble gas are difficult
smce its isotopes are short lived and are a emiters. Attempts
were mad~ to prepare fluorides of krypton by heating
gaseous mixtures but these experiments met with no success. Kirshenbaum and Grosse <30 ), however, were able toprepare krypton tetrafluoride by passing an electric discharge through krypton, fluorine mixtures at liquid nitrogen temperature. Krypton tetrafluoride which is a transparent colorless solid has a vapor pressure of 115 mm. at
25 °C. and_ i~ m1;1ch less stable than its xenon analogue, thed~composltlon mto t~e elements being rapid at 60 °C. The
difficulty of formation and thermal instability of thisfluoride are consistent with the value of 18 to 19 kcal.
obtained for the average bond energy, by extrapolationfrom the values calculated for the fluorides of arsenic
selenium and bromine. These results are given in Table 11
and ~igure 8. In v:iew of the greater bond energy ofbromme monofluonde (59.4 kcal.) compared with thevalues for the trifluoride and pentafluoride ( 48.3 and 47.7
kcal. respectively) it _is P?ss·ible that the average bondenergy for krypton d1fl.uonde may be appreciably greaterthan 19 kcal. making this a thermally much more stablecompound than the tetrafluoride. It is of interest that
Turner and Pimentel <31 >, by the matrix isolation technique, have isolated a species which they identify as
krypton difluoride. The same authors were unable toprepare an argon fluori~e by similar methods. It appears
unlikely that argon fluondes can exist.
In view of the e~se of formation and thermal stability
of. the xenon flu?ndes there was some possibility thatoxides and chlondes could also exist. For oxides andchlorides of the noble gases to be formed exothermally ,the average bond energy needs to be greater than ,...., 30

Chemistry in Canada, August, 1963

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In view of the instability of the trioxide it is interestingto speculate on the stability of the oxyfluorides. Xenon
oxide tetrafluoride should be an exothermic compound
since the atomic heat of formation using average bondenergies of 32 kcal. and 17 kcal. respectively for Xe-F and
Xe-O is -160 kcal. mole-1 . The combined heat of atomization of four fluorine atoms ( 4 x 18.3 kcal.) and one oxy gen
atom (59.2 kcals.), 132.4 kcal., is less than this. Xenon
dioxide difluoride, by such considerations, should be endothermic. It is of interest that xenon oxide tetrafluoride iseasily made (35 ) by controlled hy drolysis of xenon hexafluoride and is a thermally stable entity (m.p. -28) whereas
xenon dioxide d.ifluoride has been observed only mass
spectrometrically (36).

NORANDA LECTURE
kcal. (since ~Hf(O) = 59.2 kcal. mole- 1 and ~Hf(Cl) =29.01 kcal. mole-I). Consideration of the average bondenergies, estimated from the available thermodynamic data
for the oxides and chlorides of antimony, tellurium andiodine, given in Table 12 and Figure 9, shows howeverthat bond energies for xenon oxides and chlorides areprobably smaller than this. Cons·equemly xenon oxides andchlorides should be endothermic compounds. The failure
to prepare such compounds from the elements is therefore
not surprising. Nevertheless there was some hope that they
could be derived from the fluorides. The Argonne workers
had noted 0 8 ) that the hydrolysis of xenon tetrafluoride
was not complete and that a solid, xenon containing,hydrolysis product formed. A study of this reaction at
the University of British Columbia showed that a fluorine
free solid was produced by the complete hy drolysis ofxenon tetrafluoride. Unforiuntely the detonation of asample of the hydrolysis product brought this investigation
to an abrupt halt. Subsequently this material was shown (32 )to be xenon trioxide, presumably arising by disproportionation of a xenon (IV) species. The trioxide is best made (33 >by the slow and complete hydrolysis of xenon hexafluoride:

Although it may prove possible to prepare xenon
nitride trifluoride ( e.g. by interaction of ammonia withxenon hexafluoride) in all probability , since the heat of
atomization of nitrogen is great (~Hf ( ) = 113 kcal.), it
will be an endothermic compound.

It occurred to us that it might be possible to prepare
the endothermic chlorides of xenon by interaction of thefluorides with boron trichloride. Figure 11 shows somepossible reactions and their estimated enthalpies. Unfortunately, although the reaction has been carried out at-78°C., the last reaction is apparently the only one.

XeF 6 + 3H 2 O ➔ XeO 3 + 6HF. The oxide is nonvolatile and is dangerously explosive, being comparable
in power to T.N.T. There is some indication that it ismore unstable in air, with a relative humidity greater than
20% (33 ). The explosixe nature of the compound is compatible with the high endothermicity ( ~Hf = + 96 kcal.
mole-I) (34l. There is a close resemblance of the crystal
structure of xenon trioxide (35 ), which is represented in
Figure 10, with that of iodic acid. The short oxygen-xenon
bonds (Xe-0, 1.76 A) of the trigonal pyramidal molecule
indicate appreciable double bond character.

There would appear to be little hope of oxides andchlorides of the noble gases, other than those of xenon
and radon, ever being made.

In a prophetic article written some 30 years ago (37)Pauling suggested the existence of XeF 6 and a less stable
XeF8 • He also postulated the existence of xenic acidH 4 XeO 6 and salts ( e.g. Ag 4 XeO 6 ) derivable from this. It
was found, at the Argonne Laboratory that hydrolysis of
TABLE 12. -AVERAGE BO D E ERGIES (INOXIDES AND CHLORIDESKCAL) FOR SOME

Oxides
TeO2
1206
U
54
37
.D.Hf XeO a (c ) = + 95 kcal. mole -Ha>0•
MHf (0 ) - 95
. . Average bond energy < - - - -- - = 14 kcal.

TABLE 11. -AVERAGE BOND ENER GY (IN KCAL) FOR SOMEFLUORIDES

Sb2O3

0

AsFa

SeF4
75 . 5(b )101. 6(c)59 _4(a)BrF[KrF2][,..._,,30]

BrFa
48 .3(a )KrF4[ ,..._,,19]BrFs
47. 7(a)[KrF6]["'18]

6(a) Gunn and Williamson, Conference on
Argonne Natl. Lab., April , 1963.
S11Cl 292 . .5(26 . 1)(27 . 8)(29.7)
Kr cg) + 1 / 2nF 2 ~ Kr F n c g)
Exothermic when A.B.E. > .D.HiFcg) = 18 .3kcal/ 116.04 k al. + Xel

4 C l - - --NOBLE GAS OXIDES

SbF:5

80

60
0

]

BrF:5 ; BrF340
20tEXOTH ERMIC1/Noble gas fluoride

0[KrF2]

XeF ; XeF[.KrF6 ];KrF4

XeF8

6

Average bond energy ( in kcal. )fluorides.

NOBLE GAS CHLORIDES

Sb2 0 3

60
40

4

20

tEXOTHERMIC

fENDOTHERMIC
XeCLiENDOTHERMIC, Noble gas fluorides

Figure 8 -LiHfO- 116 .04

--~

100

(.J

-"'

LiHfO

: -- - --XeCl 4 (g>.D.Hf0 ,...._, + 116. 04 - 40 = + 76 kcai. mole- 1

AsF3

80

oble Gas Compounds,

Chlorides
TeCJ4
54 . 5SbC!a75 . 12c1,,_2- - -- - - - - - .

( a) Wieb enger, Havinga and Boswijk, "Advances in Inorganic Chemistry
and Radiochemistry", Ed., Emeleus and Sharpe, Vol. 3, Academic
Press Inc., New York, 1961, p. 136.
( b) Nichols, Ph.D. Thesis, University of Durham 1958.
( c) Rossini et al. N.B.S. circ. 500.

100

XeOa
[17]

for some

4

Figure 9-Average bond energies (kcal.) for some oxides
and chlorides.

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39

NORANDA LE;CTURE
xenon hexafluoride in sodium hydroxide solution gave apale y ellow solution which gradually loses color, a stable
cry stalline salt being slowly precipitated. Addition ofpotassium iodide to a solution of this salt, followed byacidification, leads to the liberation of eight oxidizing
equivalents per xenon. The formula proposed (:i s) for thesalt isa~nXe0 1+ 11. y i 0 , with 2 being the most probablevalue for n. This is the first established compound ofXe(VIII) and represents the maximum oxidation state forthe element.
X eF ~ + H 2O -t X e + X eO 3 +HF+ 0 ~X eF 6 + 3 H 2O -t X eO 3 + 6HF
Crystal stru cture ( D. H. T emple ton , A. Za lkin, J. D. Forreste r
a nd S . M. Willi a mso n , J. A mer. Chem . S oc. 85 , 817 (1963 ).
Ort horhombic , a = 6.163 ± 0.008, b = 8.115 ± 0.010,
5.234 ± 0.008 A, z = 4, S.G . P212 121c =0± 0.03 A
108 + 2°

1.1s + 0.03 K

101 + 2°
Figure 10-Xenon Trioxide.

Of course it may well prove possible to isolate xenon
tetroxide despite its endothermicity and undoubtedly someeffort will be made towards this end. The tetroxide should
be a regular tetrahedral molecular species and should bevolatile.
Re-fe·rences

3 X eF i<gl + 4 BClacgl-t 3 X eCL <gl + 4 BF.'j< g)
A t .LiHf( kcal ) - 3[4(32 )] - 4[3(92.9 )] - t - 3[4( 10 )] - 4[3( 139.1)]
Li H 1.~a n= ,..._., - 290 krnl
-t 3X cCI 2<gJ + 3C b <gJ+ 4 BF 3<gJ - 3(2(22)]- 3[2( 29 _)],..__, - 4 76 kcal
3 X e(g)+6

Cl 2(gl+ 4 BF \g> - 6[2(2 9 )]= ,..._, - 518 k cal
Figure I I-Possible XeF,1-BCl~ t·eactions.
( 2) a. W einstock, Claassen and Malm , J . Am er. C hem . Soc., 1957 , 79,5832,
b . W einstock, Malm and W eaver, Ibid, 1961, 83, 4310.
( 3) Selig, Cherni ck and Malm, J. Inorg. Nuclear Ch em ., 1961, 19, 377,
( 4) Claass en, Selig, Malm , Chernick and W einstock , J. Arn er. Ch em.
Soc., 1961 , 83, 2390.
( 5 ) Chernick, C1aassen and W einstock, Ib ·i d, 1961, 83, 3 165 .
( 6 ) W einstock and Malm , Ibid , 1958, 80, 4466.
( 7 ) M alm and Selig , J. Inorg. Nuclear Chem ., 1961 , 20, 18 9 .
( 8) Ba1tlett, Jha and Trotter , Proc. Chern. Soc ., 1962 , 277.
( 9) Ayns]ey, Peacock and Robinson , J. Ch ern . Soc., 19,50, 1623 .
10) Bartlett and Lohmann , Ibid, 1962, 5253.
11 ) Bode and Vos , Z. anorg. Chem. , 1951, 264, 114.
12) E arnshaw, Figgis, L ewi aJ'.ld Nyholm , N ature, 1957, 179, 1123 .
13) Kapustinskii, Quart. R evs ., 1956, 10, 28 3.
14) Fi eld and Franklin , " Electron Impact Phenomena", Academic Press,
N ew York, 1957, p . 114.
( 15) Bartlett, Proc. Chem. Soc., 1962, 218.
(16) Claass en , Malm and Se!jg, " X eF 4 et seq" , H. H. H yman , ed . " Nobl e
Gas Compounds", Univ ersity of Chica go Press ( 1963 ) .
( 17) Claass en, Selig and M alm , J. Am er. Chem. Soc ., 1962, 84, 359 3.
( 18) Chernick et al, Sci e nce, 1962, 139, 138 .
( 19) W eeks, Chernick and Matheson , J. Am er. Che m, Soc ., 1962, 84,
4612.
( 20) Malm, Sh eft and Ch ernick, Ibid, 1963, 85, 110.
( 21) Dudley, Gard and Cady, lnorg. Chern., 1963, 2, 228 .
( 22) Stein and Plurien , "Thermochemical Studies of XeF 4, H. H. Hyman ,ed. " Nobl e Gas Compounds", University of Chicago Press ( 1963 ).
( 23) Siegel and G eb ert, J. Amer . Ch em . Soc ., 1963, 85, 240 .
( 24) L evy and Argon, Ibid, 1963, 85, 241.
( 25) T empleton, Zalkin , Forrester and Williamson, Ibid , 1963, 85, 242.
( 26) Burns, Argon and L evy, Science, 1963, 139, 1209.
( 27) Agron, Begun, Levy, M ason , Jones and Smith, Ibid, 1963, 139, 842 .
( 28) Claassen, Ch ernick and Malm, J. Amer. Ch em. Soc ., In Press .
( 29) Fi elds, Stein and Zirin, Ibid , 1962, 84, 4164.
( 30) Grosse, Kirsh enbaum , Streng and Streng, Sci ence, 1963, 139, 1048.
( 31) Turn er and Pimentel, "Preparation of Inert Gas Com p ounds by Matrix
Isolation T echnique", H . H. H yman, ed. "Noble Gas Compounds",
University of Chicago Press ( 1963).
( 32) Williamson and Koch, S.cien ce, 1963, 139, 1046.
( 33) Smith, J. Am er. Chem . Soc., 85, 816.
( 34) Gunn, " Th e H eat of Formation of Xenon Trioxide", H . H. Hyman
Ed . "Noble Gas Compounds", University of Chicago Press ( 1963) .
( 3.5 ) Chernick, Cla assen , Malm and Plurie n, "Xenon Oxytetrafluoride" ,H . H. Hyman, ed. " Noble Gas Compounds", University of Chicago
Press ( 1963).
( 36) Studier and Sloth, "Mass Sp ectrom ebi c Studies of Rare Gas Compounds". Ibid.
(37) P auling, J. Amer . Chem. Soc., 1933 , 55, 1895.
( 38) Malm, Bon e , and Holt, "The Hydrolysi. of Xenon Fluorides andthe Isolation of Stable Xenates", H . H. Hym an, eel. "Noble Gas
Compounds ", University of Chica go Press ( 1963 ).

*( 1 ) W einstock , Rec. Chem. Frog. 1962, 23, 23 .

j ■■ ■■■■•■■■■■■■ a ■■■■■■■■■■■■■■- ■•·•••••••••••• ■■■■■■■■• ■••••••••••••••••••••••• 11 ■■■•••••·•••••■■■■■■■■ II ■■■■■■■■■■■■■■■■■■■■■■ II ■■■■■■■■■■■■■■■■■■■■■■ 11 ■■■■■■■■■■■■■■■■

BooKs IN REVIEWICOMPOUND SEMI CONDUCTORS,
Volume I, Preparation of 111-V Compounds. (Robert K. Willardson, Harvey
L. Goering, Editors; Reinhold Publishing
Corp., 430 Park Ave., ew York 22, N.Y.
xxii + 553 pp. $25.00 U.S. ) - This compilation on compound preparation comprises 50 chapters written by scientists
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for each of the chapters on the basis ofa bibliography compiled from papers
published during the last 10 years. Oneof the authors is S. C. Liang, Cominco,
Trail, B.C.

40

RADIATION EFFECTS ON ORGANIC MATERIALS. (Robert 0. Bolt,
James G .Carroll, editors; Academic Press
Inc., 111 Fifth Ave., ew York 3, N .Y.
xv
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written for Research and Development
people so it goes light on fundamentalsof physics, chemistry and engineering,but emphasizes background and recent
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four general chapters, the following
chapters on materials are included : polymers, plastics, elastomers, coolants, lubricants, adhesives, textiles, coatings andfilms, dielectric fluids, fuels and fluid
shield materials, and coal, wood andexplosives.

+

CHEMICAL CALCULATIONS - An
Introduction to the Use of Mathematics
in Chemistry. Second Edition. (Sidney
W. Benson ; John Wiley & Sons, Inc., 605

3rd Ave., New York 16, N.Y. xii + 254
pp. ) - This new editiori contains a larger
number and variety of problems, withmore weight on 'difficult' ones. The sections dealing with acid-base phenomena
and complexionic equilibria were completely rewritten.
TOPICS IN ORGANIC CHEMISTRY.
(Louis F. and Mary Fieser; Reinhold
Publishing Corp., 430 Park Ave., NewYork 22, N.Y. xii
668 pp. $10.00 U.S. )- Part one surveys : polynuclear hydrocarbons ; aromatic heterocyclic compounds; alkaloids; terpenoids; stero-ids;
vitamins ; chemotherapy; synthetic polymers; dyes. Part Two covers the literatureof organic chemistry from July 1, 1961
to December 13, 1962, and as such supplement the authors' previous text "Advanced Organic Chemistry" .

+( See, also p age .5 7 )Chemistry in Cttnada, August, 1963

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