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Fluorides of Copper Silver Gold and Palladium.

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Fluorides of Copper, Silver, Gold, and Palladium
By Bernd G. Muller*
Dedicated to Professor Rudolf Hoppe on the occasion of his 65th birthday
Fluorine, by far the most reactive of the non-metals, is capable of forming a large number
of compounds with nearly all other elements (exceptions (so far): He, Ne, and Ar), even
under comparatively “mild” reaction conditions. These compounds usually differ markedly
from those of the heavier halogens in composition, structure, and chemical and physical
properties. Thus, for example, it is generally quite easy to prepare fluorides containing elements in high oxidation states (often their maximum), as in AgF,, CsAgF4, PdF3, CsAuF,
etc., whereas the corresponding chlorides, bromides or iodides are in many cases (still) unknown. Conversely, the synthesis of fluorides containing these elements in middle or low
oxidation states often meets with considerable difficulty, even where it is possible at all, as,
e.g., in the case of CuF, AuF, PtF,, SeF,. Finally, there are also some examples of compounds MX,, which with X = F are stable, but with X = C1 are unstable or decompose easily
(e.g. CoFJCoCI,, VF4/VC14, PbF4/PbC14, AsF,/AsCI,). Consequently, fluorine compounds are of great general interest.
1. Introduction
The nineteen-fifties heralded a new phase of development in fluorine chemistry. It began, o n the one hand,
with the preparation of complex transition metal fluorides
(and oxides) in high and in some cases previously unknown oxidation states (K2NiF6,K3CuF6, CsAgF, etc.),[’-’]
first by W. Klemm, then later by R . Hoppe and co-workers,
and, on the other, with the use of BrF, as a solvent and
fluorinating agent (e.g. by H. J . Emelkus et al.). It culminated in the (independent) discovery of the first noble gas
fluorides by N . Bartlett (“XeFYF6”)L41
and R . Hoppe et al.
(XeF,)[’I in 1962. In the following years it became possible,
using more sophisticated methods of fluorination (e.g.
high pressure fluorination, or the use of KrF, as a fluorinating agent) to synthesize further compounds of unusual
constitution, some likewise in previously unknown oxidation states, for example Cl: PtF;
KrF* AuF; ,[’I and
C S ~ C U F ~(see
) ~ ]also Refs. 19, lo]).
However, despite the remarkable progress that has been
made regarding preparative methods, our knowledge of
this area of inorganic fluorine chemistry, where single
crystal X-ray diffraction measurements are essential for
the unambiguous characterization of the constitution and
structure of the compounds, is still only fragmentary. U p
to now such investigations have mainly been restricted to
thermally stable, involatile and, in most cases, ionic fluorides which are capable of being grown in the form of single crystals from the melt, with air and moisture excluded
if necessary, and using flasks o r crucibles of appropriate
materials (the Bridgman/Czochralski technique and its
variants).[”] Only seldom is it possible to prepare anhydrous monocrystalline fluorides without decomposition:
[*I
PrivwDoz. Dr. 8. G. Miiller
Institut fur Anorganische und Analytische Chemie der Universitat
Heinrich-Buff-Ring 58, D-6300Giessen (FRG)
Angew. Chem. Inl. Ed. Engl. 26 (1987) 1081-1097
1. from concentrated H F solutions (e.g. cs,PtF6,
K2MnF6)I2,13] or under hydrothermal condition^;"^'
2. by crystallization from anhydrous H F as solvent (e.g.
XeF,.SbF,,“’]
H30+TiF,,[’61 AgF2.2SbF,”71)-a
method whose application is limited by the sparing solubility of many (ionic) fluorides in HF;
3. by growing crystals from the melt,[lX1
either by transport
or by gel
4. by sublimation without decomposition, under an inert
atmosphere (e.g. Ar), under fluorine, o r in a high vacuum (e.g. RhF,,”91 A u F ~ , ~ ”(XeF,):+PdF:-L2’1).
]
Very few examples of the preparation of single crystals
of fluorides which are easily decomposed thermally (especially in the absence of F2) have thus far been reported in
the literature.
As already known, the structures of many metal fluorides can be derived from those with the most densely
packed anions, as a result of their essentially ionic bond
character, with the cations distributed between the available sites (usually octahedral, less frequently tetrahedral) in
a usually ordered manner corresponding to the molecular
formula.
Following this simple building principle, many classes
of compounds can be assigned to a small number of basic
structural types. Thus, for example, in the VF3 structure
vanadium can be replaced by any of the similarly charged
ions (A13+, Ga3+, Ti3+, Cr3+, etc.), or by M”/M1”
( pM”M1VF6with M“ =Zn, Ni, Co, .. ., M’“ =Ti, Sn, Pd,
Pt ...), o r by M‘/MV ( a M I M V F , with M1=Li, N a ...,
MV= Sb, Pt ...)[,’I without altering any basic structural features.
The same holds true for complex fluorides, a n example
being the RbNiCrF,-type structure derived from the pyr o ~ h l o r e s . ‘Here
~ ~ ~substitution of other metals in place of
Rb, Ni and C r gives a remarkably large number of new
quaternary fluorides with the same structure and arrangement of cations. With fluorides, therefore, one is often in
0 VCH Verlagsgesellschafi mbH, 0-6940 Weinheim. 1987
0570-0833/87/1111-108l$ 02.50/0
1081
the fortunate situation of being able to dispense with single
crystal measurements and to limit oneself to interpreting
the generally less informative powder diffraction data, providing an exact single crystal structure determination is already available for at least one example.
There are, nevertheless, certain exceptions which d o not
allow such simplified analogies and isostructural considerations. These are the cases where one or more of the cations involved in building u p the structure require a less
symmetrical arrangement of the ligands, owing to the existence of eg states with odd-numbered occupancy (JahnTeller effect). Those which are relevant for fluorides include Cr’+ and Mn3+ (3d4 “high spin” configuration:
t&e,!), Cu2+ and Ag’+ (3d9/4d9 configuration: &e;), and
Ni3+,C u 4 + , and Pd3+ (3d7/4d7 “low spin” configuration:
t2,ei).
Unusual structures are sometimes also found in the case
of fluorine complexes of C u 3 + , Pd2+, Ag3+, and Au’+
(nd’ configuration, where n = 3 , 4 , 5 ) , as a result of the
rarely occurring coordination number C.N. [ M ” + ] = 4 in
relation to fluorine (with F- ions in a square planar configuration around the central ion).
Because of these special structural characteristics, such
fluorides have already been the subject of many of the
early studies, but there are some further reasons why the
fluorides of the coinage metals seem a particularly attractive subject for investigation:
1. The elements Cu, Ag, and Au lie on the border between
transition metals and Main Group metals (“meta-metals”), but are “true” transition metals in the narrower
sense insofar as, in contrast to the elements Zn, Cd, and
Hg, d-electrons can still be involved in their bonding.
2. They are the only elements (apart from, in the widest
sense, a few lanthanoids, e.g., Ce, Pr ...) which occur in
oxidation states higher than that corresponding to their
position in the Periodic System of the Elements, i.e.
their outer electron shell configuration.
3. The individual physical and chemical properties of the
elements differ more widely, and are less susceptible to
extrapolation than in any other group of the Periodic
System-one need only consider their chemistry in
aqueous solution.
4. The fluorides are of particular importance insofar as
they are the only compounds in which certain oxidation
states can exist.
An interpretation of the, in some cases remarkable, physical properties (magnetism, spectra, conductivity etc.) of
these compounds requires a detailed knowledge of the
structures; consequently, careful single-crystal measurements are essential.
However, there are usually considerable difficulties in
preparing pure microcrystalline samples of the fluorides of
interest in this connection, and the production of suitable
single crystals poses even greater problems. The following
report describes some methods for synthesizing thermally
less stable fluorides of the coinage metals and palladium,
and outlines the associated problems (e.g. side reactions
and decomposition of products) that are commonly met
with; it also gives an overview, mainly in the form of ta1082
bles, of our present state of knowledge of this area of
chemistry.
2. Copper Compounds
Copper occurs in combination with fluorine in the oxidation states I , +2,
3, and +4, with the + 2 state as
clearly the predominating one. Fluorine complexes of
mono-, tri- and tetravalent copper are listed in Table I . Because of the very large number of known copper(i1) compounds, Tables 2 and 3 (p. 1084) list only a few representative examples chosen more or less arbitrarily, but limited
to those whose structures have been determined on single
crystals.
It is surprising that, apart from CUASF~/’~I
fluorides of
Cu’ (including C u F itself) are so far totally unknown, despite numerous attempts at producing them. This contrasts
with copper(1) oxides, e.g. CuzO, ~ [ C U , O , ] , [ ~or~ ]chlorides such as CuCl and KCuCI,. Copper differs markedly
in this respect from the homologous element silver, in
which it is well known that the + 1 oxidation state predominates, and not only as regards its chemistry in aqueous solutions. Copper is also unique amongst the coinage
metals (even though based on only a few examples) in its
ability to form compounds having the + 4 oxidation state,
namely M:[CuF6], where M ’ =(K), Rb, Cs.[’.29.30.3’1
These
are of particular interest from the structural chemistry
standpoint in having (according to original powder diffraction data) a slightly distorted orthorhombic (and from
more recent single-crystal data tetragonal) structure of the
same type as K2PtC16.Thus, single crystal studies are necessary in this case to determine the Cu4+ coordination
sphere more exactly (d7 “low spinlhigh spin” configuration), and to give a deeper understanding of other features,
e.g. magnetic properties. Understandably, however, their
preparation is by no means easy. This is related to the fact
that pure samples of, for example, Cs2CuF6, can only be
obtained by (in solid reaction terms) rapid conversion of,
e.g., Cs2CuF, using F2. If one attempts prolonged annealing (in order to grow crystals), decomposition occurs by
uncontrollable and incompletely understood mechanisms
which are affected by pressure, temperature, and the material of the crucible. Some of the possible products from
side-reactions or decomposition are indicated in Scheme 1.
One is therefore limited to either carrying out experiments over longer periods at relatively low temperatures,
or to repeated annealing for short periods (at higher temperatures) on the same sample. This also applies for all
other compounds which can only be prepared under similar conditions. Using the second (rather laborious) method, it has recently been possible to prepare single crystals
even of CsZCuF6,and to determine the structure in detail.
The basic problems involved in synthesizing single crystals of thermally sensitive, highly reactive fluoro-complexes of the transition metals (especially those containing
alkali metals or alkaline earth metals) are already evident
from the above discussion, namely the large number of
possible side-products or decomposition products, and of
reactions between one or more components (especially
CsF) and the walls of the vessel during lengthy experi-
+
+
Angew. Chem In1 Ed. Engl. 26 (1987) 1081-1097
Table I . Copper fluorides with oxidation states
Compound
Ref.
Color
Cu As F,
(241
[25, 261
NaKuF,,
+ I,
+3, +4.
Angles
["I
C.N. [Cu'")
dJCu"'-F]
[A1 [a1
Magnetic
properties
a=55.7
(6)
-
diamagnetic
6
1.845(2~)
1.917(2x)
1.919(2x)
(6)
-
p=2.66 B.M.
(contaminated with
NaF/CuF2)
p=3.03 B.M.
a = 8.894
6
1.830
p=2.83 B.M.
a = 5.58
C = 12.04
a = 7.067
b = 7.277
C = 10.32
a = 6.282
4
1.730
diamagnetic
Cu'+ 4 + 2
1.980(4x)
2.130(2x)
1.900
1.772(2~)
l.750(4 x )
-
Structure
Lattice
constants
colorless
LiSbF,
a = 5.49
green
(R3
cryolite
(P21/n)
[A]
a = 5.454
b = 5.667
C = 7.843
a = 8.51
KKuF,,
[261
pale green
Cs2KCuF,, [c. d]
[26, 271
pale green
Cs[CuF4]
1281
orange
CsCuCuF,
1261
black
[dl CsrCuFh
18, 291
orange-red
K2PtCI,, Variant
(14/mmm)
RbOF,,
[30, 311
orange-red
K2RCI, Variant
(Immm?)
KiFeFh [b]
(Fm3m)
elpasolite
(Fm3m)
KrBrF,
(14/mmm)
CsNiNiF,
(Imma)
8=90.2
Cu'+ 6
4i2
8.885
a = 8.835
b = 5.978
C = 6.007
p s 1 . 7 6 B.M.
C=
-
-
-
[a] Interatomic distances found from powder data, where these are given, are by their nature inaccurate (except in structures without F parameters, eg. AgF), and
are often based on analogy considerations, or (occasionally) on MAPLE calculations. The same remarks apply to the rest of the tables. [b] The structures which
have been postulated for these and other compounds of the form M:M"'F6 ( M ' = K , Rb, Cs; M"'=Ga, Fe, Co ...) are in need of correction, as shown by a
number of studies (see, for example, [26, 32-34]). [c] lsostructural type: MiM'CuF,, light green to beige ( M ' = K , Rb, C s ; M'=Li, Na, K, Rb) [261. [d] Single
crystal data available for this compound.
Cs2CuF4. CsNiF6. Cs,NiF,,
T
Cs2CuF4
5
4OOOC.
p(F&
n
t
% 1-2
J
d
200-400 bar
Cs2CuF6
5
CsNiF,, Cs,CuF6.
4OOOC. t
22
CsCuF,...
d
Cs2CuF4, CsCuCuF6, CsCuF4. Cs3CuF6,..
I
Cs2CuF4, Cs3ALF6, CsAIF4. CsCuF4, CsCuAlF6. CsCuCuF6 . . .
Scheme I . Side-reactions in attempts to prepare Cs2CuF6 single crystals
ments. Further, a reduction in F2 partial pressure (as a result of fluorine being lost through reaction with the autoclave material) can cause thermal decomposition of those
compounds which are particularly sensitive.
Also, owing to the fact that practically all the fluorides
of interest here are insoluble (without causing hydrolysis
or decomposition) in the usual solvents, and that other
more appropriate methods of purification such as distillation or sublimation are not available, it is impossible in
principle to remove impurities from samples after preparing them.
The fluorine complexes of trivalent copper are generally
much more stable, and have consequently been more thoroughly studied. A notable example is the successful preparation of single crystals of Cs2KCuF, in closed copper crucibles (under argon!). In this case the basic elpasolite
structure of the complex evidently permits an "optimal"
arrangement of the individual ions (stable matrix with a
high lattice energy) which gives exceptional thermal stability.[*'] Other examples too support this conclusion.
Angelr.. Clwm. In,.
Ed. Engl. 26 (1987) 1081-1097
In addition to these, copper(1ri) fluorides are known,
namely CS[CUF,][~*~
and CSCUCUF,,[~~]
but these can only
be made under more extreme conditions (F2 pressures
above 400 bar), and are correspondingly lacking in stability. The existence of Cs[CuF,] (isostructural with K[BrF,]) is
"reassuring", since in this respect copper follows the behavior of the two other coinage metals, which in their + 3
oxidation states with fluorine show predominantly (for
Ag3+) or exclusively (for Au3+) a coordination number of
4, with a square planar arrangement of F- ions around
M3+.
CsCuCuF, is interesting as regards its structural chemistry, being a very sensitive compound of mixed valency
C u 2 + / C u 3 + . According to powder diffraction data it is
isostructural with C S N ~ N ~ F , ,but
' ~ ~ the
' results d o not allow one to determine either the positions of the Cu'+ (3d",
Jahn-Teller effect) and Cu3+ ions or the F parameters
(proximity to the strongly scattering Cs ions) and thus to
calculate reliable values for the distances d[Cu"+-F]. Nor
is it easy in such cases to propose plausible structural mod+
1083
els, as the structures of C S P ~ P ~ F , R
, ~~~C’U~ P ~ F , ,and
’~~]
C S A ~ F ~ F ,show,
‘ ~ ~ ]and single crystal data are therefore essential.
In general, a considerable amount of knowledge exists
concerning the structure and constitution of copper( 11)
fluorides, which, as a result of their 3d9 configuration, differ markedly from compounds of similar formula (so far as
these are known) formed by the adjacent elements nickel
and zinc. The reason for this is that CuF, and fluorine
complexes of C u 2 + (compared with Cr”, Mn3+, or Ag2+
compounds) can usually be prepared without difficulty,
even as single crystals, and there exists a correspondingly
large amount of exact (single crystal) structural data-see
Tables 2 and 3 . Ternary and quaternary systems containing
Table 2. Copper fluorides with oxidation state
~~~~
~
Compound
~
~401
CuPtF,,
+ 2 (cationic). Single-crystal data measured for all compounds
~
Ref.
CuF2
[381
Color
Structure
colorless
orange
Lattice
constants
m i l e variant
(P2dn)
LiSbF, variant
(Pf)
ReO, (variant)
(Fm3m, disorded)
unique type
(pi)
a-CuZr F,
141, 421
colorless
FCu[AuF,]
[431
yellow
RbCuPdFS
1381
orange
CdPdPdF, variant
(&ma)
KCuAIFo
[Ma1
colorless
CsAgFeF, type
(Pnma)
KCuCrF5
[Mbl
green
unique type
(P2JC)
Table 3. Copper fluorides with oxidation state
Compound
Ref.
C u 2 + are usually obtained by annealing suitably chosen
mixtures in copper bombs or Au/Pt tubes (welded under
Ar). The only exceptions to this are compounds in which
one o r more of the cations involved (in polynary complexes) is thermally sensitive, that is decomposes with
change of oxidation state upon heating, or can be readily
split off in the form of a binary volatile fluoride. This is the
case in the CuMF, series (M=Ti, Sn, Mn, Cr, Pt ...), for
example, and as a result it has only recently been possible
to determine the structure of CuPtF, single crystals.*38’The
synthesis of CuPtF, must be carried out in a closed autoclave, because of the volatility of PtF5 and PtF,; however,
for the same reason too high a pressure of F2 must be
avoided. Once made, however, CuPtF, has-apart from its
Color
a = 3.307
b= 4.546
c = 4.599
a = 4.952
b= 4.982
C = 9.624
a = 5.613
a = 5.536
b= 4.979
C = 3.707
a = 6.296
b= 7.199
C = 10.76
a = 6.731
b= 7.040
c = 9.743
a = 7.266
b= 9.945
C = 6.769
[A1
Angles
C.N. [Cu2+]
d6Cu2+-q
[A1
Magnetic
properties
1.910(2x)
I .929(2 x )
2.305(2x)
1.930(2x )
2.007(2 x )
2.122(2 x )
1.985
-
p=2.06 B.M.
2+2+2
1.854(2x)
2.021(2 x )
2.221(2 x )
antiferromagn
g(251 K)=
1.32 B.M.
4+2
1.928(2x)
I .959(2 x )
2.190(2x)
-
2+4
1.873
1.881
2.12 l(4 x )
1.872 2.007
1.878 2.239
1.999 2.261
-
1’1
4+2
B=
96.57
a = 89.9
2+2+2
B = 104.2
y = 120.4
(6)
a=107.6
98.77
y = 98.63
b=
a=
2+2+2
92.79
/L=
1.84 B.M.
-
+ 2 (anionic). Single-crystal data measured for all compounds.
Structure
Lattice
constants
[A]
Angles
[“I
C.N. [Cu”]
d [ C u 2 + - R[A]
KCUF,
[45]
colorless
perowskite variant
(14/mcm)
a = 5.863
C = 7.847
2+2+2
1.888(2x)
1.96212 x )
2.258(2 x )
K,CUZF,
1461
colorless
unique type
(14/mmm)
a = 4.156
~=20.52
1+2+1+2
I .907
1.960(2x)
CS~CUF~
[47]
pale blue
K2NiF4
(14/mmm)
a = 4.403
C = 14.03
CSICUAFIC
1291
bluish cast/
colorless
unique type
(pi)
a = 8.505
a = 113.9
fl= 1 11.8
pale blue
unique type
(P2,/c)
b= 7.740
C=
6.268
a=11.73
b= 6.266
C = 15.68
a = 5.643
c = 10.67
CS~CU~FIP
1291
SrCuF,
1481
blue-green
Ba2CuF,
[49]
colorless
n-Ra2Cu5F,,
[SO]
colorless
KBrF4
(14/mcm)
unique type
(Bbma)
unique type
(C2/C)
a = 5.937
b= 5.837
c = 15.85
a = 18.17
b= 6.652
C = 10.328
1.933(2 x )
I .961(2 x )
2.474(2 x )
y = 103.1
B=
99.19
C U ( ~4 )+ 2
Cu(2) 2 + 2 + 2
Cu(l) 2 + 2 + 2
Cu(2) 3 + 2 + 1
Cu(3) I 1 + I + I + I
4
+
412
fl= 117.1
Cu(l) 2 + 2 + 2
Cu(2) 4 + 1 + 1
Cu(3) 2 1 - 2 1 2
I084
1.927
2.225(2x)
+I
1.902(2x) 1.944(2x)
2.251(2 x )
1.865-2.226
1.853-2. I7 I
1.850-2.405
1.845-2.353
1.858
1.862(2 x )
1.867(2 x )
2.320(2 x )
1.875-2.372
1.91 1(4 x )
2.275
2.469
1.865(2p)
2.056(2x) 2.148(2x)
Angew. Chem. Int. Ed. Engl. 26 (1987) 1081-1097
susceptibility to hydrolysis-surprisingly good thermal staThe copper fluorides, considered as a whole in cornparibility, and single crystals can be grown in welded platinum
son with those of the other coinage metals, remain the
most thoroughly characterized, especially those of Cu2 .
tubes (under Ar) at about 750°C (!).
The surprising results that are occasionally found (espeGenerally speaking, the coordination sphere of C u 2 + is
cially concerning the stoichiometry and environment of
of most interest. Usually one expects this to occupy an
C u 2 + , as for example in Cs7Cu6FI9or FCu[AuF,]) on.the
elongated octahedral environment with a coordination
one hand, and some yet unsolved problems (e.g. regarding
number C.N.[CuZ+]=4+2, but in the limiting case
the existence of copper(1) fluorides, or the exact structures
C.N.ICu2+]=4 (square planar). Here, though, it is to some
of Cs2CuF6 and CsCuCuF6) on the other, show that it is
extent a ''philosophical'' question as to whether one decertainly worthwhile to take a fresh look at problems rescribes the environment as (still) octahedral or (already)
garded as already solved, or at systems believed to be fully
square planar-the transition is a gradual one. Numerous
understood.
examples are known-Tables 2 and 3 list just a few, with
the corresponding interatomic distances. In addition, however, quite a few fluorides have been reported, mainly in
recent work, in which C u 2 + has been found to be in a n
"abnormal" coordination state, or in which different coor3. Silver Compounds
dination states occur alongside each other in the same
compound. Abnormally coordinated are, e.g., KCUAIF,[~"'
Silver, like copper, shows valencies of 1, +2, + 3 and
and
KCuCrF6[44b1 with
C.N. [Cu2+]= 2 4,
(formally) + 4 in its fluorides. In contrast to copper, howand
ever, compounds of Ag+ are by far the most stable, while
FCU[AUF,][~~'
and
with C.N.[CuZt]=2+2+2;
oxidation states are all very unstamixed coordination is found in C S ~ C U with
~ F ~ ~those
~ ~in ~the~ remaining
~
ble thermally (in the absence of fluorine), and are accordC . N . [ C u Z + ] = 3 + 2 + 1 ; 2 + 2 + 2 ; I + 1 + 1 + 1 + 1 + 1 (this
ingly difficult to prepare as pure samples or single crystals.
assignment is, of course, arbitrary-but some evidence is
obtainable from ECoN calculation^^^'^^^^). Other phases
In addition to these "chemical" differences one also
often finds, in comparable systems with the same oxidaexist in the CsF/CuF2 system besides Cs7Cu6FI9;in this
tion states, differences in constitution and structure, and
case Cu2+ resembles its neighbor element nickel, which
associated differences in the coordination geometry at the
similarly forms numerous compounds with C s F (cf., e.g.,
central ion.
Ref. [53]), rather than the homologous element silver, for
Table 4 lists a few selected silver(]) compounds, together
which only CSA~F,''~'and Cs2AgFJ'"'' are known up to
with all the known fluorides containing silver in the + 3
now.
+
+
+
Table 4. Silver fluorides with the oxidation states (+0.5),
Compound
Ref.
Color
bronze
+3, + 4 ( + 3 / + 5 ) .
Structure
anti-Cdl,
(R3m)
colorless
NaCl
(Fm3m)
colorless
CSCl
(fi3m)
inverse NiAs-type
(P6,mc)
perowskite variant
monoclinic
colorless
colorless
Ag,CuFeF, [b]
+ I,
Lattice
constants
dbAg"'-F]
[A1
Magnetic
properties
6
2.45 1
-
6
2.467
diamagnetic
a = 2.945
8
2.550
(diamagnetic)
a = 3.264
C = 6.226
a = 4.06
b = 3.83
c = 4.06
a = 7.243
b= 10.47
C = 7.769
6
2.436
(diamagnetic)
blue
unique type
(C2/m)
a = 9.351
b= 6.991
c = 7.801
red-brown
unique type
-
yellow
KBrF4
(14/mmm)
elpasolite
(Fm3rn)
yellow
C.N. IAg"+]
a = 2.916
5.691
a= 4.395
weberite
(Imrn2)
orange
Angles
1'1
C=
colorless
purplered
orange
[A]
a = 5.90
-
p=
91.2
4+4
-
2.421(2 x )
2.540(2 x )
2.788(2 x )
2.907(2 x )
2.355(4 x )
2.798(4 x )
paramagnetic
g = 2 . 8 3 B.M.
4
1.90
paramagnetic
p=1.15B.M.
diamagnetic
6
2.13
4+4
@=115.7
c=11.15
a = 9.175
paramagnetic
p = 2 . 6 B.M.
paramagnetic
p = 1.39 B.M.
KiPtCI6
(Fm3m)
a = 8.907
K2RCI6
(Fm3m)
unique type
(I4/mrnm)
a = 9.01
paramagnetic
-
diamagnetic
[a] Single-crystal data measured for this compound. [b] Isostructural are Ag,M"M"'F,, colorless (MI'= Mg, Co, Ni ...; M"'=AI, Ga, Sc, In ...) [60]. [c] Isostructural are M'AgF,, yellow (MI= Na, Rb) 1641.
Angew. Chem. In/. Ed. Engl. 26 (1987) 1081-1097
1085
and + 4 (or +3/+5) oxidation states, insofar as these
have already been described in detail.
Ag2F is anomalous and is unique in its stoichiometry,
being the only subfluoride so far known. Ag,F can be prepared, for example, by reaction o f AgF with finely divided
silver, o r by electrolysis of AgF in H F ; it is obtained as
bronze-colored platelet-like
and is an electrical
conductor. The structure is of the same type as that of antiCdl,, and consists of alternating Ag double layers (metallic bonds) and intercalated fluorine ( F - ) layers (ionic
bonding). Ag,F is thus one of the very few examples of
cluster compounds (in the broadest sense) containing F-.
It is reminiscent-with a little imaginative licence-of
the suboxides of Rb and C S , [ ~although
~]
in these cases
Rb,O$+ or Cs,,O:+ clusters are present (two Rb6 or three
Cs, face-sharing octahedra with an 0,- ion at the center
of each octahedron).
Fluorides of Ag+ are (mainly) ionic. As might be expected on comparing atomic radii (for C.N.=6,
and r[Na+]= 1.02 A[691), they are often
r[Ag+]= 1.15
found to resemble the corresponding N a + compounds
(this is true, for example, for the binary fluorides AgF and
NaF). A detailed description of the Ag' fluorides in this
respect is therefore unnecessary. Furthermore, their preparation does not usually call for any special precautions;
they can be obtained (similarly to the copper(i1) fluorides,
which are entirely comparable in this respect) by heating
intimately mixed samples of the corresponding binary
fluorides under Ar in sealed Ag bombs or welded tubes.
Only the preparation of dry AgF presents some difficulties; direct fluorination of finely divided silver at room
temperature (using a 1 : 10 mixture of F2 with N,) has
proved to be a suitable method. The presence of some
AgF, here causes no problems; on the contrary, by evolving fluorine it inhibits the partial decomposition to elemental silver (probably caused by a reaction with the container walls) at high temperatures.
A
It can further be seen from Table 4 that fluorine complexes of silver in the + 3 and + 4 (+3/+5) oxidation
states are rare; so far, only a few have been prepared as
pure samples and studied in detail. Such complexes include, aside from C S , K A ~ F , , ' ~the
~ ~ species M'AgF,
( M ' = N a , K, Rb),@,' for all of which there exist X-ray
powder data.[,,] Also magnetic data are available for
and spectroscopic data for M'AgF, (MI= K,
Rbi4'I). Other compounds which have been reported'"' are
CsAgF, and "BaAgFS" (or "Ba[AgF,],", which is analogous to B~[AUF,];~"~),but no further details of these are
yet available.
The purple-red paramagnetic complex Cs2KAgF6 (with
an elpasolite structure!) occupies a special position, being
up to now the only confirmed example of a fluoroargentate(ir1) with Ag3+ being 6-coordinative with F-. The rest
of the silver(ir1) fluorides, which again are not many in
number, contain silver in a square planar environment
(C.N.[Ag3+]=4, with the possible exception o f AgF,); this
certainly also holds true for CsAgF,, "BaAgF,"
(Ba[AgF,],), and the recently reported XeF,.AgF3
( = XeF:AgF,),Ih7I
all of which are yellow and diamagnetic. with still unknown structures.[*]
1086
Little is known, at least with regard to structure, about
the red-brown compound AgF,, which is by far the most
sensitive and most difficult of the trivalent silver compounds to isolate; it is prepared by the oxidation of AgF,
with KrF2 in H F (20"C, 2 d), but the product may still contain some impurities. Its color and magnetic properties indicate an octahedral Ag3 environment, comparable with
the isoelectronic Pd2+ in PdF, (C.N.[Pd2+]=61721)rather
than (as in the case of the mainly planar fluoroargentates(1ri)) with Au3+ in AuF, (C.N. [ A U ~ + ] = ~ ' ~ The
'~).
mixed valency formula Ag2+[Ag4+F6]'- has also been
proposed (somewhat analogous to PdF, = Pd"PdfVF6), although this to some extent conflicts with the findings for
Cs2AgF6 (here the replacement of Ag4+ by Ag3+/AgS+
should explain some experimental results).
For C S , A ~ F , ' ~ ~(obtained
]
as an orange-colored, lightsensitive, and very unstable substance by high pressure
fluorination of, for example, Cs,AgF,), the formula
Cs2Ag/,':Agi5F6 appears entirely plausible (the isoelectronic PdF, too can be envisaged as a mixed valency compound Pd"Pd'VF6J'31 and Ag3+ can also be replaced by
G a 3 + , as in CsZGab':Agi5F61661)),
but the actual structure,
as also in the case of AgF,, can be clarified only by single
crystal studies. It is surprising in this context that it has not
yet been possible to obtain any pure Ag" compounds (such
as ''CsAgF,"), although C S A U F , ' ~and
~ ~ the (isoelectronic)
C S , P ~ F ~ can
[ ~ ' ~be prepared without difficulty at normal
pressure (pF2
5 1 bar).
It is easy enough to emphasize on paper the importance
of synthesizing single crystals of these very unstable substances, but putting this into practice may still be only a
distant prospect, at least with the techniques currently
available. It appears that for these compounds (and also
for PTF,,[~~]
Cs,KNdF7 and
and others) a
frontier has now been reached, with further progress hindered by the difficulties of synthesizing single crystals using existing methods. The decomposition (with release o f
F,), which already commences at room temperature or
slightly above leads one to suspect that the substances described here are already thermodynamically unstable toward decomposition to "low valency" fluorides and elemental fluorine in the temperature ranges specified.
Although AgF2 was first described more than fifty years
ago, and its magnetic properties were also investigated
soon afterward^,'^^.^'] there was for a long time uncertainty
about its true nature. At first it was envisaged, by analogy
with A g o (= Ag'Ag"'02[8'1) as a mixed valency compound
(= Ag'[Ag'''F,]), and it was only later shown, by more precise magnetic measurements (combined with a thorough
understanding of collective magnetic properties), and by a
neutron diffraction structure determination (on powder
samples),'**' that silver in AgF2 is actually in the 2 oxidation state.
Surprisingly, fluorine complexes of Ag" -in contrast to
those of Ag3+1631-were not synthesized until comparatively recently. Then, however, their intense colors (at least
+
+
[*] In the meantime the structure of XeFiAgF, has been determined by a
single-crystal structure analysis; as suspected Ag'+ has C . N . = 4 and resides in a square planar environment ( B . Zemva. private communication).
Angew. Chem. lnl. Ed. Engl. 26 (1987) 1081-1097
around the central ion (As") than in other M 2 + ions of
comparable size.
2. In the series Pd2+, Ag2+, Cd2+ one finds, as a result of
the increase in the radius T M 2 + (or in the ratio r M 2 + /
rF-), a transition from the rutile structure with
C.N.[M2+]=6 (PdF,) to the fluorite structure with
C.N.[MZ+]= 8 (CdF2). CdPdF4,['01 which is (formally)
"isoelectronic" with AgF2, and the isostructural high
pressure modifications of PdFz191.9z1
and of AgF2l9I1already crystallize in a CaFz-type structure with
C.N.[M2+]=6 (+2), with a highly distorted octahedral
arrangement of F- around M"; for AgF2 itself and for
the corresponding fluoride complexes one therefore
might expect to find (at the transition) a n essentially
singular structural behavior.
3. Neither binary nor simply constructed complexes (i.e.
those in which there is no additional stabilization from
organic ligands such as py, bpy etc.) of Ag2+ are yet
known, and, as is shown by the high redox potential of
Ag2+ (Eo(Ag+/Ag2+)= 1.98 V in aqueous solution!),
they are scarcely to be expected. The highly unstable
for simple inorganic fluorides, and in comparison, e.g.
with copper(I1) fluorides), their widely varying magnetic
properties, and their individual structural properties with
sometimes unusual stoichiometry, stimulated the preparation of further silver(r1) compounds in rapid succession
(Table 5).
In a few cases it was possible to derive the crystal structures from powder diffraction data (based on isostructural
types), while "plausibility" considerations and MAPLE
calculation^[^'^ led to "credible" interatomic distances
d[Ag" -F], and gave the first indications concerning the
coordination behavior of divalent silver. As a rule, however, more exact information on structure could not be obtained, and spectroscopic data@31or magnetic measurem e n t ~ [ ' ~provided
]
little further clarification.
Divalent silver is particularly interesting, at least from
the structural chemistry standpoint, for several reasons:
1. Ag2+, with its 4d9 outer electron configuration, is like
C u 2 + a cation in which, owing to the Jahn-Teller effect,
the ligands need to be less symmetrically arranged
Table 5. Silver fluorides with oxidation state +2.
Compound
Ref.
Color
Structure
Lattice
constants
C.N. [Ag"]
d!Ag2+-Fj
IAI
Magnetic
properties
AgF2
[54, 821
blue-black
unique type
(Pbca)
a = 5.548
b= 5.834
C = 5.094
4
2.07
(2.59)
antiferromagn
p = 1.07 B.M.
brown
perowskite variant
(Pbnm)
a = 6.186
b = 6.207
C=
8.300
2t4
2.08(2 x )
2.20(4 x )
p=O.81 B.M.
[A]
Angles
["I
-
antiferromagn
KZNiF4
(14/mmm)
a = 4.581
14.192
4
C=
violet
KBrF&
(14/mmm)
a = 6.038
c=11.46
4
2.05
blue-green
unique type
(Pnma)
(I=
7.338
b = 7.564
C = 10.554
2+4
2.049 2.277(2 x )
2.055 2.300(2*)
p = 5.95 B.M.
ultramarine blue
CsAgFeF,
(Pnma)
U=
7.380
2+4
2.063 2.293(2 x 2.081 2.297(2 x )
p = 1.30
blue-green
CsAgFeF,
(Pnma)
a = 7.19
b = 7.39
C = 10.32
2+4
2.045 2.31 l(2 x )
2.051 2.316(2x)
(antiferromagn.)
green
RbNiCrF,
(Fm3m)
a = 10.79
6
2.0s
antiferromagn
p=1.16 B.M.
light-blue
(CuF'tF,)
(PI)
-
-
-
paramagnetic
p = 1.99 B.M.
unique type
a = 9.248
b = 6.677
C=
9.064
Adl) 4+2
Ag(2) 4 + 2 + 2
2.066(4
. x ). 2.354(2
. x .)
1.998
2.147(2p)
2.092
2.608(2 x )
2.789(2 x )
antiferromam
/?= 91.14
paramagnetic
p=1.87 B.M.
violet
b= 7.241
C=
blueviolet
(C2/c)
paramagnetic
p = 1.92 B.M.
antiferromagn.
antiferromagn
B.M.
10.352
blue
unique type
(pi)
a = 7.809
b= 5.700
C=
5.832
a = 106.1
8=111.5
y = 96.62
2+2+2
2.048(2 x )
2. M ( 2 x )
2.339(2 x )
light blue
unique type
(pi)
a = 9.044
b= 5.596
C=
5.518
a = 118.76
4+2
/?= 91.53
2.067(4 x )
2.367(2 x )
a = 5.224
b = 5.446
C = 8.779
/?= 78.02
__
weakly ferromagn.
p = 1.49 B.M.
p
= 1.95
B.M.
y= 102.42
(I=
75.78
5(+2)
2.09(4 x )
2.43(2 x )
-
y = 65.29
[a] Structurally related are: M'AgF,, brown ( M ' = R b , Cs) 1541. [b] MiAgF,, violet ( M ' = K , Rb) structure not yet known 155). [c] Isostructural type: M"AgF,,
violet (M"=Ca, Sr, Cd, Hg) 1841. [dl Single-crystal data measured for this compound. [el Isostructural type: M'AgM'I'F,, blue to green (M'= K, Rh, Cs; MIi'= K,
light blue to black (M'"=Ti, Pb,
Fe, C a ) 1391. [fI lsostructural type: M1AgM"'F6, green ( M ' = C s ; MI"= In, TI) (851. [g] Probably of isostructural type: AgMLVF6,
Pd, t'F ...) 1861. [h] lsostructural type: Ag,Zr2F,4, blue-violet (single crystal data), related are: M"Ag,MbvFI,, red-violet (M"=Mg, Ni, Zn: MIV=Zr, Hf) and
AgMYMi'F,,, emerald-green (M"=Ca, Cd, Hg; M"'=Zr, HO [87]. [I] lsostructural type: Ag[NbF6I2, light blue (881 and possibly AgF2.2SbFs, green [89].
Angew. Chem.
In!. Ed. Engl. 26 (1987) 1081-1097
1087
oxides Ag3041X91
and Ag203,'931which have recently been
prepared as single crystals by anodic oxidation at low
temperatures (about OOC), d o show, however, that oxoargentate(rr), o r oxoargentate(ii1) complexes are conceivable. Consequently, if information about the spectroscopic, magnetic o r structural chemical properties of
Ag" is required, one has in principle to rely on studies
of the corresponding fluorides.
The synthesis of microcrystalline (powder) samples of
AgF, and complex fluorides of Agz+ is now relatively
easy, by direct reaction of suitable starting mixtures with
elemental fluorine ("Klemm/Hoppe fluorination^'"^^').
Growing single crystals suitable for X-ray structure analysis, on the other hand, sometimes involves considerable
difficulties. The reason for this is simple: fluorides of divalent silver decompose with loss of fluorine (in the absence of F2) at temperatures between 150 and 200°C depending on the particular compound, and, if crystals are
required, prolonged annealing (6-10 weeks or longer) at
relatively high temperatures (above 580°C, say) under a
partial pressure of F2(depending on the compound) is necessary. There is, of course, much corrosion of the vessel
(with associated impurities and side-reactions) if common
materials are used, and it is therefore not surprising that
structure determinations from single crystal data have only
been successfully achieved in recent times.
The first such success was achieved with silver(r1) compounds of the composition M1Ag"M"'F6 (M'=Cs, Rb;
MI1'= AI, Fe).L391
Owing to the high reactivity of M F (wall
reactions) and the possibility of silver being oxidized further with the formation of M[AgF4], careful attention to
maintaining correct reaction conditions (F2 partial pressure, crucible material, temperature) was essential.
2 CsCl
+
Ag20
(CsAgFeF6). CoAgF,,
+
It was found that Mg boats "passivated" with fluorine
(formation of MgF2) were especially suitable as sample receptacles, since MgF, forms neither mixed crystals nor any
kind of ternary compounds with AgF,, and is also very stable towards other fluorides, especially those of the alkali
metals (cf. Scheme 2).
In the first structure determinations carried out on
CsAgFeF,, Ag2+ was found to have a compressed octahedral environment (an unusual situation for a "Jahn-Teller
ion"; cf. Table 5). In order to confirm this unexpected result, the investigations were repeated (with similar results
in each case) on CsAgAIF, and RbAgAI,-.Fe,F, (Fe3+
being partially replaced in the latter case by A13+ from the
corundum boats used!). However, since in parallel with
this the same structure was also found in the case of
NH4Fe1'Fe"lF6'95'(same space group, same occupancy of
the corresponding lattice positions, comparable interatomic distances), with no involvement of a cation with d4
or d9 configuration in this case, the compression of the octahedron around Ag'+ could well be mainly a structural
effect. It appears that in both the elpasolites and the pyrochlores (and variants thereof), the C u 2 + ion (cf. Table 3 )
or Ag2+ ion adapts to the usually less favorable environment with C.N.[M2+]=2+4, so as to optimize the overall
arrangement of the differently charged ions.
This rather vague notion is, however, substantiated to
some extent by the existence and structure of CsPd"Pd"F5
and RbCu"Pd"F,. In both these compounds the distribution of all the atoms at the lattice positions corresponds
exactly with that in CsAgFeF, and NH4FeFeF6, except
that one F - position remains unoccupied to preserve the
electric charge balance. This leads to half the number of
Pd2+ ions having a square planar configuration with
Fez03
(CsAgFeF6), CsALF,,
I
(-CIF,
FZ/Ar
z1:5
CsAgF,,
ClF3)
CsFeF4, AgF,.
FeF3. ALF,.
0,
/
AgF. AgF,,
CsFeF4,
CaFeF5. AgFeF,...
CaF>
T
F,/Ar
5
2
M g boot
T N 580-620 OC
550%
I: 10
,
/
CsAgFeF6 (single crystals)
(CsMgFeF,.
A$.,
CsAg, .xMg,FeF6)
t z 6-10 weeks
F2-passivoted
(CsAgFeF6). CsAgF,.
(CsMgF3), AgF,.
CsFeF,,
\1
FeF3 . . .
CsAgFe, .$L,F6,
CsALF,.
AgF. FeF3...
CsFeF,,
AgF,...
Scheme 2. Conditions needed for preparing CsAgFeF, single crystals, with side reactions.
1088
Angew. Chem. Int. Ed. Engl. 26/1987) 1081-1097
C.N.[Pd2+]=4 in relation to F-, instead of being in a
compressed
octahedral environment
(like Ag2+,
Fe2+)).[’7.381 Figures l a and I b show the similarity of the
two structures.
at the outset, remain essentially unchanged on substituting
many different cations. It is therefore not surprising that
this structural principle continues to hold true for other
compounds of completely different stoichiometry. This is
so, not only for AgPdZrzFl, (cf. Tables 5 and 8) and
NaM”ZrzF,, (M” = Ag, Pd), formally derived from
Ag(l)Ag(2)2ZrZF,4 by replacing AgZ+(l) by N a + and
“eliminating” AgF,), but also for CaMnF,,[97J and for
NaMn3F,o,’981as can be seen from Figures 2-5.
All these crystal structures are based on coordination
polyhedra sharing edges or corners in the “equatorial
plane” (i.e. octahedral, pentagonal and hexagonal bipyramids) and all have nearly planar layers, built up in a
“base o n peak” arrangement (i.e. they are linked through
F- ions in a trans position).
It can be seen from this description that loosely packed
layers containing only F- ions (e.g. with a shortest interatomic distance d[F--F-]=3.48
in AgPdZrZFT1
(Fig. 2))
alternate with layers which, in addition to F - , contain all
the cations (shortest fluorine-fluorine distance in the
above compound: d[F-F] =2.27 A!). The individual structures differ only in the order of stacking, according to the
various stoichiometries.
As a final example of the synthesis of single crystal samples containing Ag’+, we consider Ag[TaF6lz. Owing to the
ease of decomposition into the binary components (in a
stream of fluorine gas at temperatures above l5O0C) and
the volatility of TaF,, one must in this case use a closed
system (autoclave). The production of single crystals is
briefly outlined in Scheme 4.
It is interesting here to note that single crystals of
Ag[TaF& are always found in the cooler, upper part of the
autoclave. Since AgF, (m.p. >62O”C, under F,) certainly
does not sublime under these conditions ( T 54OO0C), TaF,
must participate (with the formation of volatile complexes
of unknown constitution that are stable only at high pressures) in the transport through the gaseous phase.
The layer structure is very simple: three AgF, octahedra
(with approximate D4hsymmetry-see also Table 5 ) are attached in the cis position above a very obviously distorted
TaFb octahedron (cf. Fig. 6). The remaining three F - of
the TaF6 octahedron are terminal, i.e. no linkage via further cations takes place between each of the layers. Thus
each AgF, octahedron has six different TaF, octahedra,
each TaF, octahedron, however, only three different AgF,
octahedra as ligands (for further details cf. Ref. [90]).
A
Fig. I . a) Crystal structure of CsAgFeF, with chains of [FeF,] and [AgF,]
octahedra; b) crystal structure of CsPdPdFs with chains of IPdF,] octahedra
and planar (PdF,] units.
The synthesis of Ag3Hf2F,4single crystals is somewhat
more difficult, as we here have a phase which is comparatively rich in silver. Accordingly, the compound readily decomposes and annealing must therefore be performed under nearly pure fluorine (cf. Scheme 3). Any small variations in the fluorine concentration above the sample cause
partial decomposition with the formation of Ag+, and the
sample melts and “creeps” over the edges of the boat.[”]
The existence of a compound with the composition
Ag3HfzF14is quite surprising, since in the case of C u 2 + , o r
the appreciably larger ions Sr2+, Pb2+ and Ba2+, under
similar conditions one obtains only M”MfVF6(M“ =Cu,
Sr, Pb, Ba; M’“=Zr, Hf).141.961Its structure too is very unusual, as Agz+ is found to have markedly different coordination numbers at adjacent sites (namely 4 + 2 and
4 + 2 2). Consequently, it is possible to selectively replace[x7Jthe first of these two types of Ag2+ ions by C u z + ,
Ni2+, Znz+ o r Mg2+ and the other type by Ca2+,C d Z +o r
Hg2+ (although not by SrZ+ o r BaZ+, which produce
bright blue compounds of unknown structure).
This compound, whose structure at first sight appears so
complicated, provides a further example of the stability of
certain types of ionic structures, which, as was mentioned
+
Ag2HfF6, AgF,.
HfF, ...
F2/Ar
‘r
5
1 : 10
580-600°c.
t z 1-2 h
2 10 : 1
l4 T z5 58O-60O0C.
?,
5:1
t
z5 6-10 w e e k s
I
CaF2 boat
F2/Ar
Mg boat. F2/Ar
F
Ag Hf
’
Ag3Hf,F14
(single crystals)
(Mgfig, .,Ag2Hf2F14)
T N 58O-60O0C
t z5 8-9 weeks
AgAg2. ,CaHf2F14, CaAgF4, Ag3Hf2F14, Ca[AgF4]2..
Scheme 3. Conditions for the synthesis of Ag3Hf2F,, single crystals, with side reactions.
Angew Chem Inr. Ed. Engl. 26 (1987) 1081-1097
I089
A
A
5
C
B
A
Fig. 2. Layer structure and layer sequence in AgPdZr,F,, (NaAgZr,F, ,).
Black squares: Pd2+:black hexagons: Ag+: black pentagons: Zr4+: circles:
F-.
A
Fig. 3. Layer structure and layer sequence in Ag,Hf2FI,. Black squares:
Ag2+(1);black hexagons: Ag2’(2); black pentagons: HC’ : circles: F
A
D
D
C
C
8
A
Fig. 4. Layer structure and layer sequence in CaMnFs. White squares: Mn’+ ;
white pentagons: Ca2+(i);black pentagons: Ca2+(2);circles: F-.
1090
Fig. 5. Layer structure and layer sequence in NaMn,FIo. Black hexagons:
N a + ; white squares: Mn’+(l): black squares Mn’+(2); circles: F-.
Angew. Chem. I n t . Ed. Engl. 26 (1987) 1081-1097
I
Ag[TaF6I2
(powder)
p(F2)
‘
N 10-100
T
s 250-350
t
N
OC
bar
Aq,O
+ 2 Ta205
2-10 d
AqTaF,
+ TaF5
(+1/2
p ( F 2 ) N 2-3 kbar
r
n 380-4OO0C
t ::4-6 w e e k s
F2)
As we can see from the examples given here, silver(I1)
compounds are actually quite unpredictable in their structural chemical behavior. The conditions needed to prepare
them-at least for single crystals-vary considerably and
cannot be simply extended from one to another, so that
one must treat each case individually and find the optimal
conditions, sometimes through a tedious succession of experiments calling for much patience.
Fig. 6. The layer 5:rucIure of Ag[TdF&, consisiing of corner-sharing [AgF,]
and [TaF,] octahedra.
4. Gold Compounds
In contrast to the rather “enigmatic” behavior of silver
fluorides, and to a lesser extent that of copper fluorides,
the chemistry and especially the structural chemistry of
gold fluorides are remarkably lacking in surprises. This is
largely due to the fact that neither gold(I), gold(I1) nor
(clearly defined) gold(1v) fluorides have thus far been described. Only fluoroaurates(Ii1) and a few fluoroaurate(v)
complexes are known. They are all yellow, the coordination number relative to fluorine being always
C.N.[Au3+]= 4 for the trivalent gold compounds (square
planar configuration), and C.N. [Au”] = 6 for the pentavalent gold compounds (octahedral environment).
Further, all the compounds are diamagnetic, excepting
where one of the other cations is paramagnetic. Their preparation too is generally comparatively easy, and is carried
out under similar conditions for the different compounds.
Angew. Chem. Int. Ed. Engl. 26 (1987) 1081-1097
’
Aq[TaF&
(single crystals)
Scheme 4. Conditions for the synthesis of
Ag(TaF& single crystals, with side-reactions.
In the case of the ternary gold(1Ir) fluorides, the volatility of AuF, in the stream of fluorine (T>350”C) causes
something of a problem; nevertheless it is possible1”] to
obtain AuF3 as orange-red needle-shaped crystals up to
several mm in length. Single crystals of the ternary fluorides listed in Table 6 can be obtained (similarly to the
C u 2 + and Ag2+ compounds) by annealing fine mixtures of
the starting components in welded Au tubes (under dry
Ar).
It is interesting from a crystallographic point of view
that so many different structure types are found for the
same empirical formula M”[AuF4],. It appears that here
the structure is determined (in a subtly graded way) by the
sizes (i.e. ionic radii) of the individual cations. Thus for
M“[AuF4], with different metals one finds different types
of structures as follows: M ” = Mg, Ni or Zn (monoclinic),
M ” = Pd (orthorhombic), M ” = C a or Sr (monoclinic),
M “ =Cd, Hg or Pb (tetragonal-P), M i ’= Ba (tetragonal-I),
with coordination numbers ranging from C.N. [MZ’]=6
(e.g. Mg2+) to C.N.[M2+]= 12 (BaZ+).14’]Copper, instead
of forming “Cu[AuF4],”, forms the compound FCu[AuF4],
in which only one of the F - ions in CuF, is formally replaced by [AuF4]-. The structure of this compound (see
Fig. 7) is surprisingly simple: linear [-F-Cu-F-] chains lie
along the [OOl] direction, around which, at the same
“height” as Cuz+, four isolated square planar [AuF,] units
are so arranged that one F- ion of each such unit completes the (highly distorted) octahedron around the
Cuz+.1431In contrast the structures of FM:’i[AuF4]s-in
which we again see that the F- ions of MF3 are only partly
formally replaced by [AuF,]- -are extraordinarily complicated. Remarkable is the occurrence of [M:lrFl7] units
(two corner-sharing trebly capped trigonal prisms), which
are linked by (isolated) [AuF4]- units to form a threedimensional structure (cf. Fig. 8).
Comparatively little is known about gold(v) fluorides;
the first known gold(v) compound was described only
quite recently.”021 This is rather surprising, since CsAuF,
can be obtained by the direct conversion of CsAuC1, with
elemental fluorine at about 300°C ; to prepare homogeneous samples, however, increased fluorine pressures are recommended.*’O’lIt appears that the surprisingly easy formation and the (relatively) good stability of CsAuFd (toward
decomposition into CsAuF, F,) were so unexpected that
the results of early experiments led, at first, to wrong inter-
+
1091
Table 6. Gold fluorides with oxidation states + 3 and +5.
Compound
Ref.
Color
Structure
redorange
unique type
(P6 I 22)
KBrF,
(I4/rnmm)
yellow
yellow
yellow
yellow
yellow
a = 9.788
C=
7.600
a = 5.536
b = 4.979
C=
3.707
b=
1.91(2x)
2.04(2 x )
1.915
diamagnetc
4
1.892(2 x )
1.924 1.928
diamagnetic
4
1.898(2x)
1.931 1.937
diamagnetic
4
1.907(2 x )
1.930(2 x )
antiferromagn
p = 1.32 B.M.
4
1.889 1.905
1.890 1.926
paramagnetic
p = 7.47 B.M.
4
1.883 1.975
1.885 2.008
paramagnetic
p=11.45 B.M.
-
-
-
-
diamagnetic
diamagnetic
6
1.85
1.86(4x)
I .90
diamagnetic
-
-
paramagnetic
p = 1.66 B.M.
-
diamagnetic
109.3
a = 107.6
B=
98.77
y= 89.63
a = 8.955
b = 7.960
B=
90.83
diamagnetic
7.960
-
(RhFs?)
LiSbF6
Magnetic
properties
4
a = 8.310
c=25.897
C=
orange
yellow
d[Au"+-F]
IA1
4
a = 6.182
C = 11.91 1
unique type
unique type
(~4~2~2)
unique type
(P2h)
C.N. [Au"+]
16.26
a = 5.847
b = 5.518
C = 10.83
unique type
(pi)
Angles
["I
a = 5.149
C=
unique type
(P2 I / C )
(14)
yellow
Lattice
constants [A]
a = 5.24
a = 96.5
(R5)
Xe2F;AuF;
[a]
yellowgreen
Xe2F;RuF;
(fima)
lemonyellow
LiSbFb
(R3)
yellow
unique type
a = 9.115
b = 8.542
15.726
a = 5.001
C=
O'AuF;
IF: AuF;
a = 9.573
a = 99.5
[a] Single-crystal data measured for this compound. [b] Isostructural are M'[AuF,], yellow (MI= Na-Cs) [104]. [c] Isostructural are M"[AuF,],, yellow (MI'= Mg,
Ni); (M"=Cd, Hg, Pb, Pd, Ag) have unknown structure 1701. [dl Isostructural FM:"[AuF,jS, yellow (M"'=ln, TI, Bi, La-Tb) [99]. [el Isostructural are
FM:"[AuF&, yellow (M"'=Dy-Lu) [loo]. [fl lsostructural are MAuF6, yellow ( M = N O + , K) [loll.
pretations ("CsAuF,. 1.5 HF"); the same also applies in
the case of "M[AuF,].xHF" ( M = K, Rb: x = 1.5-2.0)."051
Only sparse information exists as yet on gold pentafluoride, which is, as expected, the most unstable of the gold(v)
compounds. It is obtained in the form of a dark orangered, diamagnetic X-ray amorphous solid upon careful
thermal decomposition of 0:AuF;
(or KrF+AuF,). In
the solid state (m.p. = 75 "C) tetrameric units are probably
present (as in RhF, and other pentafl~orides"'~),and elec-
Fig. 7 Crystal btructure of FCu[AuF,], with corner-sharing chains of [CuF,]
octahedra and planar [AuF,I units.
Single crystal studies have so far only been carried out
on [Xe2Fl + [ A U F ~ ] - , [whose
~ ~ ~ I preparation under comparatively "mild" conditions ( T = 400°C, pF,= 70 bar,
t = 4 8 h) is possibly made easier by the volatility (and high
mobility) of XeF,, or perhaps of both components (AuF,,
XeF,).
In AuF,, as in all other examples which have been
studied in detail, Au" is in an octahedral environment; the
interatomic distances d[Au5+-F-]= 1.85-1.90 are of the
expected order of magnitude. Other fluorine complexes
of pentavalent gold, e.g. MI'[AUF,]~ (M"=Ca, Sr) and
"FM~"[AuF,]," (M"' = La, Pr ...), can be prepared by
> 3 kbar) of M1'[AuF4IZ
high-pressure fluorination (pF2
or FM:"[AuF,], respectively, but u p to now they have only
been obtained as microcrystalline powders, owing to increasing instability.['001
A
1092
Q
Fig. 8. The [MFF,,]/[AuF.,] units in FM2[AuF& (M = In, La-Tb) and their interlinkage. Filled circles. Au"'; open circles, F - ; hexagons, Mn"'.
Angew. Chem. Inl. Ed. Engl. 26 (1987) 1081-1097
tron diffraction studies (at 493 K) have provided evidence
for the existence of [AuF5I2and [AuF5I3 species."061
Viewed as a whole, the fluoroaurate(v) complexes are
comparatively difficult to prepare, with increased fluorine
pressures being generally necessary, but most of them (e.g.
MAuF,, where M = K, Rb, Cs) are appreciably more stable
than, for example, Cs2AgF6 or Cs2CuF6.
In the interest of completeness it should also be mentioned here that in Raman spectroscopic measurements on
a mixture of NO+AuF, and NO+AuF, also the yellow
compound [NO ']2[A~F6]2- was identified. However, in
view of the comparatively small amount of information
available, further more detailed studies are needed (especially magnetic measurements) to confirm the existence of
this only known example of a gold(1v) compound.11071
The
same applies to the recently reported compound "AuF,"
(prepared by the reaction of AuF, with F2 in vacuum (!)),
for which again there exists only spectroscopic evidence.'"']
5. Palladium Compounds
The existence of fluorides of di-, tri- and tetravalent palladium is firmly established (cf. Table 7). Results on
0: PdF; containing Pd" (identified by Raman spectroscopy) need to be definitely substantiated, since under comparable reaction conditions (pF2
2 4 kbar, T= 300-4OO0C,
t = 2 h-3 d) no evidence has so far been found for the formation of "CsPdF,".
Further, the X-ray powder diffraction patterns of
0: PdF,, containing many lines, are considerably different from those of the structurally related dioxygen(1 +)
compounds 0: MF, (M = Rh, pt, Au ...) having the same
general f o r m ~ I a . " ~ ~ l
There is little difficulty in synthesizing alkali metal and
alkaline earth metal fluoropalladates(1v) (by conversion of
the appropriate starting mixtures with dilute fluorine at
temperatures of 300-500"C)175.''I
and the compounds
MI'PdF, (with M"=Ni, Mg, Zn ...),lllll including
Pd[PdF,] it~elf."~'The brick-red, diamagnetic, very unstable compound PdF, is, in contrast, obtainable only by
high-pressure fluorination of PdF3.1T12]Its structure (determined by neutron diffraction on powder samples) is unusual, and consists-in contrast to SnFJ1131-of [PdF,] octahedra linked at their corners with terminal F- ions in the
cis p~sition.'"~'The interatomic distances d[M4+-F-] are
comparable to those in SnF,.
U p to now, single crystal studies have only been carried
out on [XeF:]2[PdF,]Z-, prepared under conditions identical to those for single crystals of Xe2F&AuF,.i'171 Like
the isoelectronic Au5+ ion, Pd4+ here has an octahedral
environment of fluorine ions; the interatomic distances
have been found to be in the range d[Pd4+-F-]= 1.861.92 A, as expected and previously estimated from known
Pd'" complexes.1751
The constitution of PdF,-similarly
as in the case of
AgF2-was at first in doubt, but one can now regard the
mixed valency formula Pd"[Pd'"F6] as proven, despite the
lack of any single-crystal data. O n the other hand, "true"
palladium(1rr) fluorides are extraordinarily rare, having
been reported for the first time only recently.11'5.1161
Their
preparation is sometimes not easy, and in certain cases one
needs to use high pressures (inert gas, e.g. Ar at about 4
kbar). The compounds M:M"PdF, (elpasolites!), however, can be prepared by heating the binary fluorides under
argon at atmospheric
In view of the fact that C o 2 +(with a n identical arrangement of valence electrons) always has the t&e; ("high
spin") configuration in fluorides, and that magnetic measurements on Ni3+ compounds (e.g. on CS,KN~F,[""~)indicate a thermal equilibrium between "high spin" and
"low spin" configurations (with the possibility of a Jahn-
Table 7. Palladium fluorides with the oxidation states + 3 and + 4
Compound
Color
Structure
"PdF3"
= Pd"Pd'"F,
K2NaPdF, [a]
blackviolet
light green
(RT)
NaPdF,
gray
KBrF, variant
(monoclinic)
a=11.84
b = 5.36
C=
5.86
brick
red
unique type
(Fdd2)
a = 9.339
b = 9.240
yellow
BaGeF,
(RTm)
LiSbF,
(R3
K2PtC16
(Fm3m)
unique type
(Pca2,)
a = 7.213
Pd Fa
Ref.
LiSbF6
elpasolite variant
(F4/mmm)
Lattice
constants
d!Pd"+-Fl
[A1
Magnetic
properties
6
2.04
4i2
1.95(4x)
2.14(2 x )
paramagnetic
p=2.88 B.M.
(paramagnetic)
(4)
-
(paramagnetic)
6
1.91(2X)
1.94(2 x )
2.00(2 x )
diamagnetic
6
1.82
diamagnetic
yellow
yellow
Angle
C.N. [Pd"+]
a = 4.98
c=13.48
6
1.85
diamagnetic
a = 9.00
6
I .89
diamagnetic
a = 9.346
b = 12.78
C=
9.391
6
1.860-1.902
diamagnetic
a = 5.01
14.12
a = 8.30
C=
8.72
C=
C=
yellow
[A]
C=
fl=
114
5.828
6.989
~
~
~
~
~
~
~
~
~
~
[a] Isostructural o r structurally related are M:M'PdF6, light green to beige ( M ' = K , Rb, Cs; M'=Li, Na, K, Rb) [1151. [b] Isostructural are M"PdF,, yellow
(MI'= Ba, Sr) [ I IS]. [c] Isostructural are M"PdF,, yellow ocher (M"=Ni, Mg, Zn ...) [112]. Id] Isostructural o r structurally related are MlPdF,, yellow ( M ' = K ,
Rb, Cs) [ I 101. [el Single-crystal data measured for this compound.
Angew. Chem Int. Ed Engl. 26 (1987) 1081-1097
I093
~
Fluorides in which Pd2+ has regular octahedral coordination can be violet colored (PdF2[72.i20.i2'1
1, b rownviolet (RbPdF3L901)(both antiferromagnetic), or blue
(NaPdZr,F, AgPdZr2F,,, paramagnetic["]). In addition,
we have dark blue CdPdF,, black HgPdF4,["" and HPPdFJ91,921
(antiferromagnetic) with a highly distorted octahedral arrangement of F- around Pd2+.
As a rule, the divalent palladium compounds listed in
Table 8 are quite easy to prepare. If required, they can be
obtained as single crystals by annealing mixtures of the appropriate binary fluorides under inert gas (Ar) in welded
Pd tubes. The preparation of pure samples of PdF2, however, presents something of a problem; PdF3 (obtained by
fluorination of, e.g., metallic Pd) is allowed to react with
powdered palladium (which has previously had all traces
of oxygen removed by heating in high vacuum), under
inert gas (Ar) in welded Pd tubes at temperatures between
600 and 950°C. By this method, using T> 950°C and t = 46 weeks, it has been possible to grow excellent millimetersized, amethyst-colored single crystals of PdF2.L1201
Of the many palladium(r1) fluorides which are now
known, and which, in some cases, have been characterized
by single crystal measurements, aside from those already
mentioned, the compounds CsPdPdF,, MiPdF, (M' = K,
Rb, Cs), and CdPdF4 are especially noteworthy. In
Rb3PdF,, which is an "alkali-rich'' phase, unusual
amongst fluorometalate(l1) complexes, two crystallogra-
Teller distortion involving Ni3+), there is naturally great
interest in the comparable palladium(i1I) fluorides. ESR
and magnetic measurements indicate the existence of the
"low spin" case (tzgeA). If the arrangement of F- ions
around the Pd3+ ion could be determined, with exact values for the distances d[Pd3+-F-], this would give information concerning the nature and magnitude of the JahnTeller distortion, which would be of considerable importance.["'I
Divalent palladium behaves in complex fluorides as an
inorganic "chameleon", i.e. it adapts its coordination number, and accordingly the structure, color, and magnetic
properties in the compounds concerned, to the situation
imposed by the type and number of the accompanying cations.
In the alkali metal and alkaline earth metal fluoropalladate(1r) complexes one generally finds a square planar
environment (C.N.[Pd2+]=4), e.g. in M$PdF5 (MI= K,
Rb, Cs)
and in M"[PdF4] (MI'= Ba (orange), Sr
(red), Ca, Pb
all of which are diamagnetic. In
the compounds M'Pd"Pd"F, (MI = K, Rb, Cs), which are
orange-brown and antiferromagnetic, PdZ+occupies a half
octahedral and half square planar environment of Fions; the PdZ+ ions in octahedral positions can be selectively replaced by M2+, [M" = Mg2+ or Z n z + (yellow,
diamagnetic) and M2' =NiZ+ or Cu2+ (yellow and
orange-brown respectively, antiferromagneti~)].'~'.~~~
Table 8. Palladium fluorides with oxidation state +2.
Ref.
Compound
[72,120,121]
PdF2 [al
Color
amethyst
I1221
PO1
CsPdInF, [el
11231
CsPdPdFS [a,
fl
1371
1371
NaPdZr,F,, [a]
PdPtF, [a]
[A]
Angles
C.N. [Pd2+]
101
diPd2'-F]
[A1
Magnetic
properties
a = 4.961
C = 3.391
a = 5.327
4+2
2.1514 x )
2.17(2 x )
antiferromagn.
p = 1.90 B.M.
6
2.175
a = 5.403
6
2.205
a = 5.521
4
I .99 1
antiferromagn.
p=2.10 B.M.
antiferromagn.
p=2.58 B.M.
diamagnetic
6
2.15
antiferromagn.
p = 1.32 B.M.
-
antiferromagn.
p = 1.79 B.M.
brownviolet
brownviolet
orangebrown
unique type
(lmrna)
a = 6.53
b= 7862
C = 10.79
4
yellow
unique type
(Pd/mbrn)
a = 7.462
4
1.926
diamagnetic
4
1.940
diamagnetic
4
1.944(2 x )
1.975(2 x )
2. l54(4 x )
2.165(2 x )
paramagnetic
p=2.88 B.M.
6
2.160(2 x )
2.172(4 x )
paramagnetic
p=2.83 B.M.
6
2. I 76
paramagnetic
p = 2.88 B.M.
dark blue
CaPdF, [a, c]
RbPdF, [d]
rutile
(P4/mnm)
Lattice
constants
CdPdFa
(pa31
CaF, variant
Va3)
KBrF,
( I 4/mmm)
perowskite vari ant
(Pm3m)
RbNiCrFh
(Fd3m)
HP-PdF2
CdPdF, [a, b]
Structure
I371
yellow
Rb3PdF5
(P4/mbm)
I381
orangebrown
unique type
(fima)
1611
blue
NaAgZr2F,,
(pi)
1611
blue
unique type
(C2/m)
c = 10.570
a = 4.298
a = 10.89
I .95 1
antiferromagn.
p = 1.56 B.M.
6.457
a = 7.519
C = 6.455
a = 6.269
b= 7.199
C = 10.16
C=
a = 7.910
b = 5.146
c = 5.145
a= 9.351
b= 6.991
C = 7.801
a = 5.038
c=14.31
(antiferromagn.)
a = 106.7
p=112.2
y = 91.9
6
D= 115.7
[a] Single-crystal data measured for this compound. [b] Isostructural are HgPdF,, black 1901. [c] lsostructural are M"PdF,, purple to orange (M"=Sr, Ba, Pb) (1201.
[d] Structurally related are M'PdF,, brown-violet (MI = K, TI) 1901. [el Isostructural are CsPdM"'F6 (MI" = Fe, Sc .. .) [92]. [fl lsostructural are M'M"PdF5, yellow to
orange-brown ( M ' = K , Rb, Cs, TI; M"=Pd, Mg, Ni) [37]. [g] Isostructural are MiPdF, and M'M\'PdFs, yellow ( M ' = K , Rb, Cs, TI) 1311. [h] Isostructural are
M'CuPdF5 ( M ' = K , CS) [loo].
1094
Angew. Chem. Int. Ed. Engl. 26 (1987) 1081-1097
phically different types of Rb’ and F - ions are present.
Those rubidium ions which occupy sufficiently large holes
in the lattice can be specifically replaced by cesium (to
give CsRbzPdF,), with no appreciable change in the lattice
constants, but attempts to carry out a similar partial exchange of CI - for F- have so far been unsuccessful. Most
surprisingly, however, the same types of compounds and
structures have recently been found in complex platinum
Figure 9 shows a
hydrides, e.g. Rb3PtHs or
pictorial impression of the structure of the palladium compounds, consisting of channels with square and triangular
cross sections.
Fig. 9. Crystal structure of RbiPdFs and CsRbzPdFS with tetragonal and trigonal channels.
Finally, CdPdF, (and also HgPdF,) is of interest from a
structural chemical standpoint, since it undergoes the transition from a KBrF,-type structure, through a variant of
the CaFz type (CdPdF,), to the pure CaF, type (CdFZ), as
has already been suggested in the case of CdAgFJs4] (with
c / a = 1.99). In view of this, it seems plausible that high
pressure modifications of PdFz and AgF2, metastable at
normal pressure and isostructural with CdPdF,, can exist.
6. Concluding Remarks
Despite the remarkable progress that has been made in
this area of inorganic solid state chemistry the present state
of our knowledge regarding the fluorides of copper, silver,
gold and platinum, which has here been described only selectively, still leaves much to be desired. In the first place
there is a shortage of available methods for preparing pure
homogeneous samples (partly due to a lack of suitable
methods of purification), and secondly the growing of single crystals (especially those of fluorides with poor thermal
stability) is often hindered by severe problems in finding
Angew. Chem. Inr. Ed. Engl. 26 (1987) 1081-1097
suitable materials, and by the occurrence of obscure or unpredictable side reactions.
It is possible in some cases that the presence of one or
more volatile substances (even when one can only speculate about their existence as intermediate products or species of unknown composition) plays an important or even
a crucial role. Two final examples will again serve to illustrate this.
1. When one tries to make single crystals of Ba[AuF6], (FZ
pressure > 3 kbar) one obtains, in addition to crystals of
that compound (of moderate quality), ruby-red single
crystals of BaNiF, (formed by Ni from the monel-metal
crucible!). This compound can normally be prepared
(under comparable conditions, but without Au) only
with difficulty, and always in microcrystalline
form.11z5.1261
A possible cause of this surprising occurrence of BaNiF6 crystals is the formation of AuF5 (as a
“mobile phase” side-product) during the preparation.
However, in the absence of firm evidence this remains
mere speculation.
2. Single crystals of MnF, are only obtained (under an Fz
pressure of > 3 kbar) if traces of Oz are present (leading
to formation of the volatile, structurally related compound 0: Mn,F,). In oxygen-free experiments under
otherwise identical conditions, the samples produced
are X-ray amorphous.[’27.
”*I
Finally, we are still far from having a deeper understanding of these structures, even in such simple systems
which are formed predominantly by ionic bonding.
Thus, it remains unclear (and incomprehensible) why, in
the system MF,/ZrF, with both small (MI’= Ni, Co, Fe . ..
through to Cu) and large ( M ” = C a , Cd, Eu ...) cations,
compounds of the type M”ZrF, are formed (usually variants of the ReO, structure), whereas with Ag’+, Ag,ZrZFl4
is formed (with a noticeably more complicated structure),
and-even less comprehensible-why no compounds at all
are formed with Pd2+.
Simple models or concepts based on purely ionic bonding are inadequate to explain these results. Consequently,
even in fluorides of metals in low oxidation states, there
must be additional factors (apart from simple geometrical
influences such as those of ionic radii or local distortions
of the immediate environment due to the Jahn-Teller effect). For example, there could be a definite space requirement for occupied d orbitals (with a preferred orientation),
as has been suggested for NbF4,11291
or there may be significant contributions from covalent bonding (considering the
volatility of AuF, and HgFz!).
This is quite at variance with the view expressed by Wilhelm KIemm (on the occasion of a tribute in 1940 to the
work of 0. Ruff in fluorine chemistry): “ R u f f s work on
the preparation of fluorides was so complete and thorough
that little more remains to be done. In future research in
this field the emphasis must increasingly be on determining their physical properties”.11301
On the contrary, now as
then we need new or more refined and sophisticated methods for preparing purer, well-characterized samples, including single crystals (as indeed KIemm himself shortly
afterwards demonstrated when he introduced new fluorination techniques). The substances so prepared will then
I095
lay down the basis for measuring their real physical properties and interpreting them correctly.
I wish to thank Herrn Dip1.-Chem. D. Kissel, Herrn Dr.
M. Serafin, Herrn Dr. K . H . Wandner and, in particular,
Frau W . Schmidt f o r their extremely valuable help in pteparing this review.
Received: June 25, 1987 [A 641 IE]
German version: Angew. Chem. 99 (1987) 1120
Translated by Dr. J. K . Becconsall. Northwich (UK)
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[ I I] J. Grannec, L. Locano, P. Hagenmuller in [lo], Chap. 2.
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[I31 P. Bukovec, R. Hoppe, J. Fluorine Chem. 23 (1983) 579.
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[IS] D. E. McKee, N. Bartlett, Inorg Chem. 12 (1973) 1722.
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