CHEMISTRY 3810
Problem Set #7
Warning: cited literature articles have been added to the end of this assigment; you may not wish to print off all the pages.
Topic: Chemistry of the Group 17 and 18 Elements. You are responsible for Ch. 12 of Shriver-Atkins
1.
Preferably without consulting reference material, write out the halogens and noble gases as they appear in the periodic table,
and indicate the trends in (a) physical state (s, l, or g) at room temperature and pressure, (b) electronegativity, (c) hardness of
the halide ion, and (d) colour.
2.
Describe how the halogens are recovered from their naturally occurring halides and rationalize the approach in terms of
standard potentials. Give balanced chemical equations and conditions where appropriate.
3.
Sketch the chlor-alkali cell. Show the half-cell reactions and indicate the direction of diffusion of the ions. Give the chemical
equation for the unwanted reaction that would occur if OH– migrated through the membrane and into the anode compartment.
4.
Sketch the form of the vacant sigma* orbital of a dihalogen molecule and describe its role in the Lewis acidity of the dihalogens.
5.
Nitrogen trifluoride, NF3, boils at -129°C and is devoid of Lewis basicity. By contrast the lower molar mass compound NH3
boils at -33°C and is well known as a Lewis base. (a) Describe the origins of this very large difference in volatility. (b)
Describe the probable origins of the difference in basicity.
6.
Based on the analogy between halogens and pseudohalogens: (a) Write the balanced equation for the probable reaction of
cyanogen (CN) 2 with aqueous sodium hydroxide. (b) Write the equation for the probable reaction of excess thiocyanate with
the oxidizing agent Mn0 2(s) in acidic aqueous solution. (c) Write a plausible structure for trimethylsilyl cyanide.
7.
(a) Use the VSEPR model to predict the probable shapes of IF6+ and IF7. (b) Give a plausible chemical equation for the
preparation of [IF6][SbF6].
8.
Predict whether each of the following solutes is likely to make liquid BrF3 a stronger Lewis acid or a stronger Lewis base: (a)
SbF5, (b) SF6, (c) CsF.
9.
Predict whether each of the following compounds is likely to be dangerously explosive in contact with BrF3, and explain your
answer: (a) SbF5, (b) CH30H, (c) F2, (d) S2Cl2.
10. The formation of Br 3– from a tetraalkylammonium bromide and Br 2 is only slightly exoergic. Write an equation (or NR for no
reaction) for the interaction of [NR4][Br3] with excess I2 in CH2Cl2 solution, and give your reasoning.
11. Explain why CsI3 (s) is stable with respect to the elements but NaI3 (s) is not.
12. Write plausible Lewis structures for (a) ClO2 and (b) I205 and predict their shapes and point groups.
13. (a) Give the formulas and the probable relative acidities of perbromic acid and periodic acid. (b) Which is the more stable?
14. (a) Describe the expected trend in the standard potential of an oxoanion in a solution with decreasing pH (b) Demonstrate this
phenomenon by calculating the reduction potential of ClO4– at pH = 7 and comparing it with the tabulated value at pH = 0.
15. With regard to the general influence of pH on the standard potentials of oxoanions, explain why the disproportionation of an
oxoanion is often promoted by low pH.
16. Which oxidizing agent reacts more readily in dilute aqueous solution, perchloric acid or periodic acid? Give a mechanistic
explanation for the difference.
17. (a) For which of the following anions is disproportionation thermodynamically favourable in acidic solution: ClO–, CIO2– ,
ClO3– and ClO4–? (if you do not know the properties of these ions, determine them from a table of standard potentials.) (b) For
which of the favourable cases is the reaction rate very low at room temperature?
18. Which of the following compounds present an explosion hazard? (a) NH4ClO4, (b) Mg(ClO4)2, (c) NaCI0 4, (d)
[Fe(OH2)6][ClO4]2. Explain your reasoning.
19. Explain why helium is present in low concentration in the atmosphere even though it is the second most abundant element in
the universe.
CHEMISTRY 3810
Problem Set #7
20. Which of the noble gases would you choose as (a) the lowest temperature liquid refrigerant, (b) an electric discharge light
source requiring a safe gas with the lowest ionization energy, (c) the least expensive inert atmosphere?
22. Give the formula and describe the structure of a noble gas species that is isostructural with (a) ICl 4–, (b) IBr 2–, (c) BrO3–, (d)
CIF.
23. (a) Give Lewis formulas and formal charges for ClO–, Br 2, and XeF+. (b) Are these species isolobal? (c) Describe the chemical
similarities as judged by their reactions with nucleophiles. (d) Rationalize the trends in electrophilicity and basicity.
21. By means of balanced chemical equations and a statement of conditions, describe a suitable synthesis of (a) xenon difluoride,
(b) xenon hexafluoride, (c) xenon trioxide.
24. (a) Give a Lewis structure for XeF7–. (b) Speculate on its possible structures by using the VSEPR model and analogy with other
xenon
25. Given the bond lengths and angles in I5+, describe the bonding in terms of two -center and three-center s bonds and account for
the structure in terms of the VSEPR model.
26. Until the work of K.O. Christe (Inorg. Chem. 25, 3721 (1986)), F2 could be prepared only electrochemically. Give chemical
equations for Christe's preparation and summarize the reasoning behind it.
27. Predict the line-patterns in the 19F NMR spectrum of the following halogen fluorides. Assume that the axial position(s) is(are)
deshielded more, and that there is no chemical exchange: a) ClF3
b) BrF5
c) IF7
28. Provide an MO interpretation of the bonding in ClF3. Treat the F atoms as simple spheres (pseudo s orbitals), but consider all
the valence s and p orbitals of the central atom, as usual.
29. The Se 2I42+ cation in [Se 2I4][Sb2F10]2 has the unusual shape shown below. Like many halogen compounds, it is extremely
electron rich. This means that there are many electrons per nucleus. One bonding model for SeI2+ involves the p bonds
indicated at the right. Use these p orbitals as a basis set to construct an MO description of the dimer in Se 2I42+. It is not
necessary to provide symmetry labels for the dimer orbitals, but provide an energy-level diagram and orbital topologies. Is your
model consistent with the very weak bonding between the two halves of the dimer indicated by the observed bond distances?
(Reference: T. Klapotke and J. Passmore, Acc. Chem. Res. (1989) 22, 234-240).
30. Provide an MO description of the bonding in XeF4. Use the example done in class for BrF5 as a guide.
31. The structure of XeF6 has been the subject of continuing discussion and debate among chemists. It is not an easy molecule to
study! Describe the idealized VSEPR and actual gas phase structures of XeF6. What structure does XeF6 adopt in solution
(NMR evidence) and in the solid state? Reference: K. Seppelt, Acc. Chem. Res. (1979) 12, 211-216.
CHEMISTRY 3810
Problem Set #7
32. Interpret the NMR spectra of the C5F5N-Xe-F+ cation. (a) shows the 19F signal for the Xe-F fluorine atom only, while (b)
shows the 129Xe spectrum. C5F5N is a pentafluoropyridine molecule, attached to Xe via the N atom. The nitrogen is natural
abundance.
33. Interpret the 125Xe NMR spectrum of OXe(OTeF5)4 shown below. The structure resembles that of OXeF4, with the –OTeF5
groups replacing the fluorine ligands. –OTeF5 has the structure indicated below. Each of the larger lines in spectrum is split
into quintets by 4 Hz. The separation between the quintets is 55 Hz, while the separation between the centres of the three
clusters is 640 Hz.
34. Is the poorly resolved 125Xe NMR spectrum shown below consistent with the proposed structure Of C6F5X+?
Volume 25
Number 21
Inorganic Chemistry
October 8, 1986
0 Copyright 1986 by the American Chemical Society
Communications
Chemical Synthesis of Elemental Fluorine
Sir:
The chemical synthesis] of elemental fluorine has been pursued
for at least 173 years2 by many notable chemists, including Davy,2
F r e m ~M
, ~o i ~ s a nand
, ~ Ruff.5 All their attempts have failed, and
the only known practical synthesis of F2 is Moissan’s electrochemical process, which was discovered exactly 100 years ago.6
Although in principle the thermal decomposition of any fluoride
is bound to yield fluorine, the required reaction temperatures and
conditions are so extreme that rapid reaction of the evolved fluorine
with the hot reactor walls preempts the isolation of significant
amounts of fluorine. Thus, even in the well-publicized case of
K3PbF7,7*sonly trace amounts of fluorine were i ~ o l a t e d . ~ , ~
These failures, combined with the fact that fluorine is the most
electronegative element and generally exhibits the highest single
bond energies in its combinations with other elements,1° have led
to the widely
belief that it is impossible to generate
fluorine by purely chemical means.
The purpose of this communication is to report the first purely
chemical synthesis of elemental fluorine in significant yield and
concentration. This synthesis is based on the fact that thermodynamically unstable high-oxidation-state transition-metal
fluorides can be stabilized by anion formation. Thus, unstable
NiF,, CuF,, or MnF, can be stabilized in the form of their corresponding MF62-anions. Furthermore, it is well-known that a
weaker Lewis acid, such as MF,, can be displaced from its salts
by a stronger Lewis acid, such as SbF,.
K2MF6 2SbF5
2KSbF6 [MF,]
(1)
+
+
+
If the liberated MF4 is thermodynamically unstable, it will
spontaneously decompose to a lower fluoride, such as MF, or MF,,
with simultaneous evolution of elemental fluorine.
[MFd
-
MF, +
Since a reversal of (2) is thermodynamically not favored, fluorine
can be generated even at relatively high pressures.
Consequently, the chemical generation of elemental fluorine
might be accomplished by a very simple displacement reaction,
provided a suitable complex fluoro anion is selected which can
be prepared without the use of elemental fluorine and is derived
from a thermodynamically unstable parent molecule. The salt
selected for this study was K2MnF6. It has been known16 since
1899 and is best prepared from aqueous H F s01ution.l~
2KMn04 + 2KF
+
+ lOHF + 3 H 2 0 2
50% aq
2K2MnF6/
In the context of this communication, the term “chemical synthesis of
elemental fluorine” implies the generation of F, by purely chemical
means and excludes either techniques such as electrolysis, photolysis,
discharge, etc. or the use of elemental fluorine for the synthesis of any
of the starting materials. The regeneration of fluorine from materials
prepared from fluorine obviously is just a method for chemically storing
but not for chemically generating fluorine.
Davy, H. Phil. Trans. R. Soc. London 1813, 103, 263.
Fremy, M. E. Ann. Chim. Phys. 1856, 47, 44.
Moissan, H. C. R. Hebd. Seances Acad. Sci. 1886, 102, 1543; 1886,
103,202,256,850 1884,99,655,874; 1885,100,272, 1348; 1885,101,
1490; 1886,102,763, 1245; 1886,103, 1257; 1889,109,862,637; Ann.
Chim. Phys. 1887, 12, 472; 1891, 24, 224; Bull. So?. Chim. Fr. 1891,
5 , 880.
Ruff, 0.Z. Angew. Chem. 1907.20, 1217; Z . Anorg. Chem. 1916,98,
27.
Moissan, H. C. R. Hebd. Seances Acad. Sci. 1886, 102, 1543.
Brauner, B. Z. Anorg. Chem. 1884, 7, 1 .
Clark, G. L. J . A m . Chem. SOC.1919, 41, 1477.
Argo, W. L.; Mathers, F. C.; Humiston, B.; Anderson, C. 0. Trans. Am.
Electrochem. SOC.1919, 35, 335.
See for example: Cotton, F. A,; Wilkinson, G. Advanced Inorganic
Chemistry; Interscience: New York, 1972; p 1 1 3 .
Schmitz-Dumont, 0.;Opgenhoff, P. Z. Anorg. Allg. Chem. 1952, 268,
57.
Ryss, I. G. In The Chemistry of Fluorine and Its Inorganic Compounds;
USAEC Translation 3927; AEC: Oak Ridge, T N , 1956; p 3 1 .
Cady, G.H. In Fluorine Chemistry; Simons, J. H., Ed.; Academic:
New York, 1950; pp 293-294.
O’Donnell, T. A. In Comprehensiue Inorganic Chemistry; Bailar, J. C.,
Emeleus, H. J., Nyholm, R., Trotman-Dickenson, A . F., Eds.; Pergamon: Oxford, U.K., 1973; pp 1010-1013.
Naumann, D. In Fluor und Fluorverbindungen; Spezielle Anorganische
Chemie in Einzeldarstellungen, Vol. 2; Schneider, A,, Ed.; Steinkopff
Darmstadt, W. Germany, 1980; p 3.
0020-1669/86/1325-3721$01.50/0
(2)
‘/2F2
HF
8H20
+ 3 0 2 (3)
The literature yield of 30% was increased to 73% and can
Probably be improved further by refining the washing procedure
(use of acetone instead of HF).” The other starting material,
SbF5, can be prepared]9 in high yield from SbCI, and HF.
SbCI, + 5 H F
SbFs 5HC1
(4)
+
+
Since both starting materials, K2MnF6and SbF5,can be readily
prepared without the use of F2 from H F solutions, the reaction
K,MnF6
+ 2SbF5
-
2KSbF6
+ MnF3 + ll2F2
(5)
represents a truly chemical synthesis of elemental fluorine.
The displacement reaction between K2MnF6 and SbF5 was
carried out in a passivated Teflon-stainless-steel reactor at 150
“ C for 1 h. The gas, volatile at -196 O C , was measured by PVT
and shown by its reaction with mercury and its characteristic odor
to be fluorine. The yield of fluorine based on (5) was found to
be reproducible and in excess of 40% but most likely can be
improved upon significantly by refinement of the experimental
conditions. Fluorine pressures of more than 1 atm were generated
in this manner.
In summary, the purely chemical generation of elemental
fluorine can be achieved in high yield and concentration by a very
simple displacement reaction between starting materials that can
(16) Weinland, R. F.; Lauenstein, 0.Z. Anorg. Allg. Chem. 1899, 20, 40.
( 1 7 ) Bode, H.; Jenssen, H.; Bandte, F. Angew. Chem. 1953, 65, 304.
(18) Chaudhuri, M. K.; Das, J. C.; Dasgupta,
_ . H . S. J . Inorz. Nucl. Chem.
1981, 43, 85.
(19) Ruff, 0. Ber. Dtsch. Chem. Ges. 1906, 39, 4310.
0 1986 American Chemical Society
Inorg. Chem. 19186, 25, 3122-3124
3122
be prepared in high yields from H F solutions and have been known
for 80 years or longer. As in the cases of noble gas20or N F 4*]
chemistry, the successful chemical synthesis of elemental fluorine
demonstrates that one should never cease to critically challenge
accepted dogmas.
Acknowledgment. The author is grateful to R . D. Wilson for
his assistance with some of the experiments, to Drs. C. J. Schack,
W. W. Wilson, and L. R. Grant for help, and to the U S . Army
Research Office and Office of Naval Research for financial
support.
(20) Bartlett, N. Proc. Chem. SOC.,London 1962, 218.
(21) Christe, K. 0.;Guertin, J. P.; Pavlath, A. E. Inorg. Nucl. Chem. Left.
1966, 2, 83.
Rocketdyne
A Division of Rockwell International
Canoga Park, California 91303
Karl 0. Christe
Received August 20, 1986
Electrochemical Generation of Iron(1V)-Oxo Porphyrins
and Iron(1V)-Oxo Porphyrin a Cation Radicals
Sir:
With weakly basic ligands, such as chloride or perchlorate, the
first electrochemical l e oxidation of an iron(II1)porphyrin is
porphyrin-centered, resulting in the formation of an iron(II1)~,~
porphyrin a cation radical.] We were the first to s h ~ wthat
when the ligands are the strongly basic oxy anions, HO- and
CH30-, the first electrochemical l e oxidation is iron-centered,
providing an iron(1V) porphyrin. In our investigations (dry
CH2CI2solvent), potentials were determined by cyclic voltammetry, coulometry was determined by controlled-potential oxidation at the potentials corresponding to the CV peak positions,
and the nature of the products was established by low-temperature
spectroelectrochemistry and comparison of the spectra to those
of known porphyrin species. In addition, the identification of
electrochemically generated iron( IV) porphyrin species was verified
by chemical conversion to known species at the same oxidation
level. To obviate p-oxo dimer formation all investigations employed (meso-tetrakis( 2,6-disubstituted pheny1)porphinato)iron(111) hydroxide and methoxide salts. In this manner, we showed
that (tetrakis(2,4,6-trimethylphenyl)porphinato)iron(III) hydroxide ((TMP)Fe"'OH) on 1e oxidation (+ 1.O 1 V)4 provides
an iron(1V) porphyrin ((TMP)FeiVO),Sand the second l e oxidation (+1.13 V) gives an iron(1V) porphyrin *-cation radical
(('+TMP)FelVO).
In a recent communication in this journal, Groves and Gilbert
reexamined the electrochemistry of (TMP)Fe"'OH in wet CH,CI,
saturated with Na2C03.6 Their results substantiated our original
discovery that the first and third oxidation peaks observed with
(TMP)Fe"'OH are for the formation of (TMP)Fe"O and (2+TMP)FeIVO,respectively, but disputed the value of our second
( 1 ) (a) Phillipi, M. S.; Goff, H. M . J . Am. Chem. SOC.1982, 104,
6026-6034. (b) Phillippi, M. A.; Shimomura, E. T.; Goff, H. M. Inorg.
Chem. 1981, 20, 1322-1325. (c) Cans, P.; Buisson, G.; Duee, E.;
Marchon, J.-C.; Erler, B. S.;Scholz, W. F.; Reed, C. A. J . Am. Chem.
SOC.1986, 108, 1223-1234.
(2) Lee, W. A.; Calderwood, T. S.; Bruice, T. C. Proc. Natl. Acad. Sci.
U . S . A . 1985, 82, 4301-4305.
(3) Calderwood, T. S.;Lee, W. A.; Bruice, T. C. J . Am. Chem. SOC.1985,
107, 8272-8273.
(4) All potentials given in this paper are vs. a saturated caromel electrode
(SCE). The potentials given by Groves and Gilbert are corrected to
S C E by using the conversion factor in their paper.
(5) The nature of the oxo ligand is not known. Though the iron(1V)-oxo
porphyrin is written as (TMP)Fe"O, the oxo ligand may be -OH or
perhaps a second oxo ligand may be present since it is virtually impossible to free a polar organic solvent of all traces of water. In our
previous publications (ref 2 and 3), we used the notation (TMP)Fe"OH.
(6) Groves, J. T.; Gilbert, J. A. Inorg. Chem. 1986, 25, 123-125.
0020-1669/86/1325-3722$01.50/0
oxidation potential for the formation of ('+TMP)Fe'"O. The two
l e oxidation potentials that they report are + I .01 and + I .40 V.
They attributed our results to the presence of chloride ion impurity
and to the absence of N a 2 C 0 3and water. We show here that
chloride ion is not present in our system; we provide additional
data in support of our assignment of potentials for l e oxidation
of iron(1V) porphyrins to the corresponding iron(1V) porphyrin
K cation radicals; and we show that the potential (+1.01 V)6
assigned by Groves and Gilbert for the l e oxidation of (TMP)Fe'I'OH to (TMP)Fe"O is in actuality due to two l e oxidations.
The following observations establish the absence of all chloride
ion. In our experiments pure (TMP)Fe"'OH was used and the
solvent and electrolyte system was devoid of chloride ion. Reactions were carried out at -71 OC where CH2CI2solvent does
not undergo oxidation. In Figure 1 there is shown repetitive visible
spectral scans of the first two sequential l e oxidations of
(TMP)Fe"'OH (conditions, positions of isosbestic points, peak
heights, etc. provided in the caption). That the spectral changes
of parts A and B of Figure 1 are associated with l e oxidations
follows from their generation by controlled-potential coulometry.
From the isosbestic points there is seen to be no competitive change
of ligand nor accumulation of intermediate. Inspection of Figure
1A reveals the absence of the spectral characteristics (absorbance
at 380 and 510 nm) of (TMP)Fe"'CI. Indeed, spectroelectrochemistry at -71 "C with (TMP)Fe"'CI shows that l e oxidation
at 1.18 V is accompanied by a decrease in the Soret absorbance
at 420 nm and an increase in absorbance at 398 nm with an
isosbestic point at 528 nm. Additional evidence for the absence
of chloride ion in our experiments is shown by the observation
that the presence of trace concentrations (10-5-10-4 M) of [ ( n C4H9)4N+][CI-] in a solution of (TMP)Fe"'OH results in a CV
where the first oxidation is no longer reversible. Such is not the
case with the CV's we have r e p ~ r t e d . ~It. ~is known' that the
reduction of iron(II1) porphyrin to iron(I1) porphyrin is strongly
influenced by the axial ligand. Employing (TMP)Fe"'OH, we
find that the le-reduction potential is at -1.05 V while the potential
for l e reduction of (TMP)Fe"'CI occurs at -0.75 V (dry CH2CI,,
25 "C). The samples of (TMP)Fe"'OH employed in the electrochemical and spectroelectrochemical studies showed no evidence
of a peak potential at -0.75 V. In our hands the electrochemical
oxidations of (TMP)Fe"'OH have been found to be both chemically and electrochemically reversible.
Groves and Gilbert (working in a solvent composed of CH2CI2
wet with water and saturated with Na,C03) reported that the
oxidation of (TMP)Fe"O to ('+TMP)Fe'"O is irreversible and
occurs at a higher potential (+1.40 V ) than the potential ( + I . I 4
V) for the (TMP)Fe"'CI to ('+TMP)Fe"'(CI), oxidation. A CV
similar to theirs, with the exception of the absence of the irreversible peak at +1.40 V, is obtained for (TMP)Fe"'OH in CH2C12
that has been wet by being passed through air-equilibrated alumina
(Figure 2A). By simple visual observation of the CV, it might
appear as though the first oxidation wave (1.01 V) represents a
single l e process, as they assumed. Controlled-potential coulometry (at 1.06 V) at -71 "C showed that there are two electrons
( n = 2.1 f 0.2) associated with this wave. For the low-temperature
controlled-potential coulometry, the system was first calibrated
by using (TMP)Fe"'CI. The coulometric oxidation was monitored
by change of current measurement with time and also by running
a CV when n is equal to 1. Therefore, the first oxidation wave
must represent two le oxidations that are so close in potential that
they cannot be distinguished by CV.' A possible explanation for
(7) (a) Scheidt, W. R.; Reed, C. A. Chem. Rev. 1981, 81, 543-555. (b)
Jones, S. E.; Srivatsa, G. S.; Sawyer, D. T.; Traylor, T. '3.; Mincey, T.
C. Inorg. Chem. 1983, 22, 3903-3910.
(8) There is no reason that two l e oxidations so close in potential should
resemble a 2e oxidation with CV. Bard and Faulkner (Bard, A. J.;
Faulkner, L. R. Electrochemical Methods; Wiley: New York. 1980;
pp 233-235) point out that when the potential difference between two
l e processes is less than 100 mV the individual waves are merged into
a broad wave. Also, they show that if the difference is 35.6 m V (theoretical), which occurs when there is no interaction between the redox
groups on the substrate, then the observed wave has all the characteristics of a l e transfer.
0 1986 American Chemical Society
Acc. Chem. Res. 1989,22, 234-240
234
Sulfur and Selenium Iodine Compounds: From Nonexistence to
Significance
T.KLAPOTKEand J. PASSMORE*
Department of Chemistry, University of New Brunswick, Fredericton, New Brunswick, Canada E3B 6E2
Received May 6, 1988 (Revised Manuscript Received March 21, 1989)
It has been our goal to prepare quantitatively in
one-step reactions simple compounds that are novel in
terms of their stereochemistry and bonding, are first
examples of new classes of compounds, and open up
new areas of chemistry. Such achievements are often
discoveries rather than planned syntheses, and the
sulfur and selenium iodine cations (Table I) described
here were prepared as a result of an unsuccessful search
for S13AsF6.'
Binary sulfur iodides are unstable under ambient
condition^,^-^ and selenium iodides5 are unknown.
Before our work, there were no examples of stable
species at room temperature containing covalent S-I
ar Se-I
except Seh2-.6 We have now prepared,
usually quantitatively, a number of salts of the sulfur
and selenium iodine (and bromine) cations shown in
Table I, all of which contain covalent S-I and Se-I
bonds. In addition, these cations provide examples of
stable derivatives of S7,Se6,thermodynamically stable
npr-npr bonds ( n I3), and r*-+ bonds. Many of the
cations maximize intercationic halogen-chalcogen contacts and thus have cluster-like characteristics, e.g., the
cube-like SQI,~+(Figure 10) and the distorted righttriangular prismatic S214,+(Figure 4). The bonding
encountered in these cations has been helpful in understanding the puzzling geometries of more complex
a fuller
related species, e.g., S:+, Sei+, S4N4,and S202-,
account of which is given in ref '7.
Instability of Neutral Sulfur and Selenium
Binary Iodides
Solid S212has been characterized at -90 0C,2and SI,
at 9 K,34 but no structural evidence has been presented
for the corresponding binary selenium iodide^.^ The
instability of the S-I and Se-I bonds can be attributed
to their very low ionic resonance stabilization energies
as the electronegativity of iodine is about the same as
that of sulfur and selenium. Thus AH (eq 1) and AH
(eq 2) are -18.0 and -1.3 kJ mol-l, respectively.8 They
2 -s-I(g)
-s-s- (g) + Iz(g)
(1)
2 -Se-I(g)
-Se-Se- ( g ) + 12(g)
(2)
are even more unstable in the solid state due (in part)
-
+
Thomas Klapotke was born in Gijttingen in 1961. He studied chemistry at
t h e Technical University of Berlin under the supervision of Prof. Kopf and
obtained his Ph.D. in 1986. In 1987 he won a Feodor-Lynen-Scholarship of
the Alexander von Humboldt Foundation and spent a year as a visiting
scholar with Prof. Jack Passmore at the University of New Brunswick in Canada. Since 1988 h e has been at the Technical University of Berlin working
on fluoro organometallic chemistry.
Jack Passmore was born in Barnstaple, Devon, England, and received his
B.Sc. and D.Sc. from the University of Bristol, England, and his Ph.D. degree
(with Dr. Neil Bartlett) from the University of British Columbia in Vancouver.
He did postdoctoral work during the year 1968-1969 at MacMaster University, HamiRon, Ontario, Canada (with Dr. Ronald J. Giliespie). He then joined
the faculty at the University of New Brunswick in 1969, where h e is presently
Professor of Chemistry.
Table I
CharacterizedoBinary Sulfur and Selenium Halogen
Cations
F
c1
Br
I
S7B?
Se
s;I'
Br2S+SSBr
[(S71)21]3+
SeF3+ SeC13+
SeBr3+
Se13+
Se7+SeSeC1 SezBr6+
Se212+
Br2Se+SeSeBr 12Se+SeSeSe+Izb
(Se I+
See$n
Structure of cations determined by X-ray crystallography.
Identified in solution by 77Se NMR.
to the large sublimation energy of solid I, (62.3 kJ
mol-l). For example, CF3SI is detected as a gas, but
readily disproportionates in the solid state above -100
"C according to eq 3.9a CH3SI behaves similarly and
is also only stable in the solid state at very low temperatures, decomposing to CH3SSCH3and
2CF,SI(s)
-+
CF,SSCF,(s)
+ I~(s)
(3)
The structure of Ph3CSI, which is stable10cin the solid
state at -78 "C and in solution in the dark, has been
determined.lobEvidence for RCOSI (R = aryl) has been
presented,loabut the material has not been structurally
characterized. No neutral compound containing a room
(1) Passmore, J.; Taylor, P.; Whidden, T. K.; White, P. S. J. Chem.
SOC.,Chem. Commun. 1976,689.
(2) (a) Dasent, W. E. Non-Existent Compounds;Marcel Dekker: New
York, 1965; p 162. (b) Daneky, J. P. In Sulfur in Organic and Inorganic
Chemistry; Marcel Dekker: New York, 1971; Vol. 1, p 327. (c) Peach,
M. E. Int. J. Sulfur Chem. 1973,8(1), 151.
(3) (a) Vahl, G.; Minkwitz, R. 2.Anorg. Allg. Chem. 1978,443,217and
references therein. (b) Manzel, K.; Minkwitz, R. 2.Anorg. Allg. Chem.
1978,441, 165.
(4) Feuerhahn, M.; Vahl, G. Inorg. Nucl. Chem. Lett. 1980, 16, 5.
(5) Behrendt, U.; Gerwarth, U. W.; J e e r , S.; Kreuzbichzer, 1.; Seppelt,
K. In Gmelin Handbook of Inorganic Chemistry, 8th ed.; Springer-Verlag: Berlin, Heidelberg, New York, Tokyo, 1984; Supplement Vol. B2,
No. 10.
(6)Greenwood, N. N.; Earnshaw, A. Chemistry of the Elements;
Pergamon Press: Oxford, New York, Toronto, Sydney, Paris, Frankfurt,
1986 and references therein.
(7) Burford, N.; Passmore, J.; Sanders, J. C. P. From Atoms to Polymers: Isoelectronic Analogies. In Molecular Structure and Energetics;
Leibman, J. F., Greensberg, A., Eds.; VCH Vol. 8, in press and references
therein.
(8) (a) Johnson, J. P.; Murchie, M. P.; Passmore, J.; Tajik, M.; White,
P. S.; Wong, C.-M. Can. J. Chem. 1987, 65, 2744. (b) Steudel, R. In
Chemistry of the Non-Metals; Walter de Gruyter: Berlin, New York,
1977.
(9) (a) Minkwitz, R.; Lekies, R. Z. Anorg. Allg. Chem. 1987,544, 192.
(b) Shum, L. G. S.; Benson, S. W. Int. J.Chem. Kinet. 1983,15(5), 433.
Benson, S. W. Chem. Reo. 1978, 78, 23.
(10) (a) Kato, S.; Hattori, E.; Mizutta, M.; Ishida, M. Angew. Chem.,
Int. Ed. Engl. 1982,21,150.(b)Minkwitz, R.; Preut, H.; Sawatzki, J. 2.
Naturforsch. 1988,438,399. (c) Guaraldi, G.; Ciuffariu, E. J. Org. Chem.
1970,35,2006. (d) Maloof, M.; Smith, S.; Soodak, M. Mech. React. Sulfur
Compd. 1969, 4, 61. Frankel-Conrat, H. J. Biol. Chem. 1955,217, 373.
0001-4a42/~9/0122-0234$01.50/0 0 1989 American Chemical Society
Acc. Chem. Res., Vol. 22, No. 7, 1989 235
Sulfur and Selenium Iodine Compounds
Q
F'6'
-.-eI.-----7 --Q
iLII
0..
...
'
FYI
1111
-
d
@
FU1
Figure 2. S71+cation in S71SbFs.
Se
- 1
0-F
Figure 1. Se13+cation in SeIaMF6(M = As,Sb). Weak contacts
are indicated as here and in other figures.
....
S71MF6(M = As, Sb), vastly different from those calculated for SI3MF6. Subsequently S71MF6were prepared quantitatively, in liquid SO2 or AsF, solution16
according to eq 6-8.
4
' /8s8
temperature stable covalent S-I bond has so far been
characterized: although the human thyroid is thought
to contain such a compound.lW It is likely part of a
protein, isolated from other S-I bonds, and thus kinetically stabilized. The first synthesis and structure
of a stable neutral iodo selenide RSeI (R = 2,4,6tB~3C6H2)
has recently been reported;'l is it likely kinetically stabilized by the bulky substituent (cf. kinetically stable RP=PR and R2Si=SiR2; R bulky
groups).l2
Sulfur and Selenium Iodine Cations
Preparation and Characterization of Se13MF6(M
= As, Sb). Although neutral binary selenium iodides
are unstable, salts of Se162-6have been known for some
time. More recently we prepared SeI3MF6 according
to eq 4 and 5 as well as various other routes.13
SOZ(1)
+ 312+ 3AsF5
SOZ(1)
6Se + 91z + 10SbF5
2Se
-
2Se13AsF6+AsF3 (4)
6Se13SbF6+ (SbF3)3.SbF5 ( 5 )
The heat of reaction 4 was estimated to be -100 kJ
The Se-I bonds in Se13+ are probably only
slightly stronger than those in Se12. The crystal lattice
energy stabilizes Se13AsF6substantially.sa Presumably
SeI,2- salts are also stabilized by their crystal lattice
energies. The X-ray crystal structures of SeI3MF6
confirmed the identity of the pyramidal Se13+cationsa
(Figure 1). The average selenium-iodine bond distance
is 2.510 A and is similar to the sum of the covalent radii
of Se and I (2.493 A). SBr3MF614(N.B. SBr4 is not
known) and Te13MF68a*15
have also been prepared.
Preparation of S71MF6, (S71)4S4 ( A S F ~ ) and
~,
(S71)21(SbF6)3.2AsF3.The successful preparations of
salts of Se13+led naturally to the attempted synthesis
of the analogous SI3+compounds.' Initial reactions
were carried out with an excess of sulfur relative to the
stoichiometric amounts indicated in eq 4 and 5. The
chemical analyses of these products corresponded to
(11) du Mont, W.-W.; Kubiniok, S.; Peters, K.; von Schnering, H.-G.
Angew. Chem., Int. Ed. Engl. 1987,26, 780.
(12) (a) Cowley, A. H. Ace. Chem. Res. 1984, 17, 386. (b) West, R.
Angew. Chem., Int. Ed. Engl. 1987,26, 1201. (c) Yoshifuji, M.; Shima,
I.; Inamoto, N.; Hirotsu, K.; Higuchi, T. J. Am. Chem. SOC.
1981,103,
4581.
(13) Passmore, J.; Taylor, P. J.Chem. SOC.,Dalton Trans. 1976,804.
(14) (a) Passmore, J.; Richardson, E. K.; Taylor, P. Inorg. Chem. 1978,
17,1681. (b)Murchie, M. P.; Passmore, J. Inorg. Synth. 1986,24,76. (c)
Brooks, W. V. F.; Passmore, J.; Richardson, E. K. Can. J. Chem. 1979,
57, 3230. (d) Passmore, J.; Richardson, E. K.; Whidden, T. K.; White,
P. S. Can. J. Chem. 1980,58, 851.
(15) Passmore, J.; Sutherland, G.; White, P. S. Can. J. Chem. 1981,
59, 2876.
3AsF5
12
+ 312 + 10SbF5
42/ss8
S,(ASF&
(X
= ca. 19)
-+
2S7IAsF6
AsF3
(6)
6S7ISbF6 + (SbFs)&bF,
+ I2
-
(7)
2s71AS.F~+ [ ( X - 14)/8]s8 (8)
Attempts were made to prepare S81+according to eq
9 and 10, but S71AsF6was formed in both reactions; I2
and s8 (for eq 9), and KAsF, and sa (for eq lo), were
also quantitatively produced. Presumably S81+is
I3ASF6 + nS8
Ss(ASF6)z + KI
+
+
S7IASF6 + 12 + ( n - 7/8)s8 (9)
S7IAsF6 + KASF6 + '/es8
(10)
formed initially but it disproportionates to the more
thermodynamically stable (estimated at ca. 26 kJ mol-')
S71MF6and sulfur.16 An important factor in the stability of S71+,relative to S81+,is the lower ionization
energy of SI (836.4 kJ mol-') relative to that of s8 (872.4
kJ m0l-I).l7 The ionization energy of S5 (830 kJ mol-')
is less than that of s6 (868 kJ mol-') or S4 (1000 kJ
and consistently the radical cation S5+has been
detected in solution but not S4+or S6+.l8 Other S7+
derivatives to have been characterized are S71+,1J6
S7Br+,19(S71)213+,20
and S7+-S6-S7+(S192+).21 This
suggests that the odd unipositively charged rings SI+
and S5+(with or without substituents) are most stable
than the even-membered rings S8+or s6+.This is in
contrast to the situation for neutral rings, where evenmembered rings are the most stable, with the stability
sequence s8 > s6 > S7 >>> S5 (not isolated).22 We
have been unable to synthesize S81+(and S8Br+)1v16y19
and our attempts to make
or S61+(and S6Br+),19,20
(and S5Br+)led to the isolation of (S71)4S4(A~F6)6,20,23
(S71)2I(SbF6)3.2AsF3,20and (S7Br)4S4(A~F6)
6. l9 These
compounds were prepared quantitatively according to
eq 11 and 12 from sulfur, iodine, and the corresponding
pentafluoride.
(16) Passmore, J.; Sutherland, G.; Taylor, P.; Whidden, T. K.; White,
P. S. Inorg. Chem. 1981,20, 3839.
(17) Wagman, D. D.; Evans, W. H.; Parker, V. B.; S c h u " , R. H.;
Halow, I.; Bailey, S. M.; Chruney, K. L.; Nuttall, R. L. J.Phys. Chem.
Ref. Data 1982, Suppl 2, 11.
(18) Low, H. S.; Beaudet, R. A. J. Am. Chem. SOC.
1976, 98, 3849.
(19) Passmore, J.; Sutherland, G.; Whidden, T. K.; White, P. S.; Wong,
C.-M. Can. J. Chem. 1985,63,1209.
(20) (a) Passmore, J.; Sutherland, G.; White, P. S. J. Chem. SOC.,
Chem. Commun. 1979,901. (b)Passmore, J.; Sutherland, G.; White, P.
S. Inorg. Chem. 1982,21, 2717.
(21) Burns, R. C.; Gillespie, R. J.; Sawyer, J. F. Inorg. Chem. 1980,19,
1423.
(22) (a) Steudel, R.; Mausle, H.-J. 2.Anorg. A&. Chem. 1981,478,156
and 177. (b)Steudel, R.; Strauss, R.; Koch, L. Angew. Chem., Int. Ed.
Engl. 1985, 24, 59.
(23) Passmore, J.; Sutherland, G.; White, P. S. J. Chem. SOC.,Chem.
Commun. 1980, 330.
236 Acc. Chem. Res., Vol. 22, No. 7, 1989
4s8
+ 212 + 9AsF5
-
+ 312 + lOSbF5
'78s~
SOlW
Klapotke and Passmore
(S71)4S4(ASF&rj + 3AsF3
(11)
AeFs(1)
2( (S7I)2I)(SbF6)3*2AsF,+ (SbF3)3SbF5 (12)
We were surprised to find that Sz+ had been formed
in reaction 11, as a large excess of AsF6 in SOz or AsF,
oxidizes sulfur only to Sg2+. In fact S42+was only prepared by heating s8 and liquid SbF6 at 120 "C for
several days!24-26 It was postulated that iodine facilitates the oxidation to S4'+ and subsequently s8 was
by AsF5 in liquid
quantitatively oxidized to S4(A~F6)2
SOz in the presence of trace halogen (Iz,Br2, C12)within
minutes of the reaction mixture warming up to room
temperat~re.~~-~~
S t r U C t U r e S O f S7IMF6, (S71)4S4(ASF6)6, and (s7I)zI(SbF6)3.2AsF3 The structures of all four salts were
determined by X-ray crystallography. The S71+cations
in both S71MF61p16
salts and in (S71)4S4(AsF6)620
are
essentially identical; S71+(Figure 2) consists of a seven-membered homoatomic sulfur ring in a slightly
distorted chain configuration similar to those in y- and
6-S7,27
S70,28and Slg2+,21
but with an exocyclic iodine.
The geometries of S70 and S71+are similar, with similar
bond-length a l t e r n a t i ~ n s . These
~ ~ ~ ~alternations
~~~~~~
may be viewed as arising from the alternations present
in S7and the presence of a positively charged tricoordinate sulfur atom (connected to the iodine).'J6 The
extent of the lengthening and shortening is greatest near
the source of the perturbation: S(l)-S(7) is very long
(2.389 (4) A, bond order of 0.37)31and S(7)-S(6) very
short (1.900 (5) A, bond order of 1.76) (Figure 2). In
valence-bond terms, the structure can be viewed as
consisting of structure A and a number of other resonance structures that delocalize the charge into the ring,
the most important of which is B.
S
S
$+I
y
+
1
s-s
/
s-s
A
s-s
s=s
4
B
The S+-I distances (2.30-2.37 A)16 in all S71+salts
(including (S71)4S4(AsF6)6)
correspond to a bond order
of 1 (sum of covalent radii: 2.37 A) and are all shorter
than the s-1 bond length (2.406 (4) A) in (C6H6)3CSI.10"
These are the only examples of structural determinations of covalent S-I bonds. (CH )2SISbF6,(CH3)&SbC16,32aand (CH3)(CF3)SIMF632
fi have recently been
(24) Murchie, M. P. Ph.D. Thesis,University of New Brunswick, 1986.
(25) Ruff, 0.;Graf, H.; Heller, W.; Knock, Eer. Dtsch. Chem. Ges.
1906, 39, 4310.
(26) Gillespie, R. J.; Passmore, J. Adu. Znorg. Chem. Rudiochem. 1976,
17, 49 and references therein.
(27) (a) Steudel, R.; Reinhardt, R.; Schuster, F. Angew. Chem., Int.
Ed. Engl. 1977,16, 715. (b) Steudel, R.; Steidel, J.; Pickardt, J.; Schuster,
F. 2. Nuturforsch. 1980,35B, 1378.
(28) Steudel, R.; Reinhardt, R. Angew. Chem., Znt. Ed. Engl. 1977,16,
716.
(29) Steudel, R. Angew. Chem., Znt. Ed. Engl. 1975,14, 655. Steudel,
R.; Schuster, F. J.Mol. Struct. 1978,49, 143.
(30) Gillespie, R. J. Chem. SOC.Rev. 1979, 8, 315.
(31) Campana, C. F.; Lo, F. Y.-K.;Dahl, L. F. Inorg. Chem. 1979,18,
3060.
(32) (a) Minkwitz, R.; Preutzel, H. 2. Anorg. A&. Chem. 1987,548,
97. (b) Minkwitz, R.; Werner, A. 2. Nuturforsch. 1988,43E, 403.
zU
Figure 3. (S71)23+cation in (S71)zI(SbF6)3.2AsF,.
Figure 4. The S2I4'+cation in S414(AsF6)z.
Figure 5. The Szand 21z dimers joined via x*-x* interactions
to give SzI,2+.
reported to be stable up to -20 OC or -35 OC, respectively, and were characterized by Raman and NMR
spectroscopy.
(S71)4S4(AsF6)6
contains discrete S71+and Sz+ cations
and AsF, anions.m The+S
:
cation has a square-planar
geometry similar to those in S4(AsF6)z.0.6S0~3
and
(S7Br)4S4(&F6)2.19
The structure of the (S7&I3+cation
consists of two equivalent S71+units that have geometries similar to those observed in S71MF616and
(S71)4S4(~F6)620*23
(Figure 3), joined by a bridging iodine
atom. The structure is approximately described in
and S71+,and
terms of two resonance structures S7IZ2+
thus the bridging sulfur-iodine bond (2.674 (7) A) has
a formal bond order of 0.5. The I(l)-S(6) intercationic
distance (3.394 (3) A) in S71+itself and the corre(3.381 (9) A)
sponding I(l)-S(3) interaction in (s7I)J3+
are significantly less than the corresponding sum of the
van der Waals radii of 4.0 A. In addition, the bridging
iodine atom also has a very weak contact with each of
the S71+units (1(2)-S(6), 3.777 (8) A). Thus both S71+
and (s71)z13+
have cluster-like characteristics.
Preparation and Structure of S214(AsF6)2.In an
attempt to prepare other salts of novel sulfur-iodine
cations (e.g., s21(AsF6),cf. SzFAsF6), we reacted s4(AsF,), with an excess of iodine. One product was
characterized as SZI4(AsF6),
and was subsequently
synthesized quantitatively in liquid sulfur dioxide according to eq 13.33 Systematic attempts to prepare
f/4S8 + 212 + 3AsF5
SzI,(AsFrj)z + AsF3 (13)
-
S13AsF6were unsuccessful. All reactions lead to S&(AsF,),(s) and Iz(s) although S13AsF6formation was
(33) Passmore, J.; Sutherland, G.; Whidden, T. K.; White, P. S. J.
Chem. SOC.,Chem. Commun. 1980, 289.
Sulfur and Selenium Iodine Compounds
0
.. .- . .. . - . .... .
=
Se
0 - I
Figure 6. The Se2142+cation in Se21,(Sb2Fll)2.
Acc. Chem. Res., Vol. 22, No. 7, 1989 237
0S:,*-o+o
-{+-0IG\p
*Io~--oI.
Figure 8. Two Se12+radical cations joined via weak n*-a*
interactions to give Se2142+.
SezIz+cation has an eclipsed structure (Figure 6) similar to that of S204z-42with two SeI, units 'oined by a
weak selenium-selenium bond (2.841 (2) ) and very
weak iodine-iodine interactions (3.756 (2), 3.661 (2) A)
(Figures 6 and 8)41(i.e., the structure of Sez142+only
n, Bonding
n; Nonbonding
l'ly Antibondlng
superficially
resembles that of Sz142+).
Occupancy 2 e
2c
le
The Sez142+cation may be regarded as two SeIz+
Figure 7. The Se12+radical cation n MOs derived from pz selradical cations, joined, in part, by overlap of the odd
enium and iodine AOs.
electron in each of the P* SeIz+ molecular orbitals
(Figure 7) resulting in some bonding between all six
shown to be thermodynamically feasible; the disproatoms and a formal selenium-iodine bond order of
portionation to solid Sz14(AsF6)zand Iz(s) was shown
1.25i4l (Figure 8). Consistently, the selenium-iodine
to be even more thermodynamically f a v ~ r a b l e . ~ ~ - ~ ~
bond distances in Sez142+(2.436 (2)-2.450 (2) A) are
The structure of Sz14(AsF6)z
consists of discrete SzI?+
significantly shorter than those in Sei,' (2.510 (2)-2.513
and AsF6- ions with weak anion-cation contact^.^^^^^
and Sez12+are different from one
(2) A).& Thus Sz142+
The SzI2+cation has a unique distorted right triangular
another, and from their isoelectronic counterparts PZI4=
prismatic structure (C, symmetry) (Figure 4)33and,
and A s ~ Iwhich
~ , ~ have
~ classical all Q eclipsed geomeunexpectedly, does not have the structure as the isotries. However, they both are cluster-like, and both
electronic Pz14.34
contain npr-npa (n > 2) and P*-P* bonds. That the
The S-S (1.828 (1) A,cf. Sa,Sz(X3Cg)1.8894 A3,) and
structure of Sez12+is different from Sz12+supports our
1-1 bond lengths (2.597 (2) A,cf. Iz(g),2.662 As and 12+, contention that the geometry and bonding in the sulfur
2.557 (4) A37)in Sz142+
indicate bond orders of 2.33 and
cation are dependent on the equality of the IEs of Sz
1.33, respectively. The S-S bond distance in this cation
and Iz.
is the shortest reported. The S-I distances of 2.858 (6)
Preparation O f ( S e 6 I ) , ( A S F 6 ) , and S e & ( A S F 6 ) p
A and 3.195 (6) A are comparable to sulfur-iodine
Several
allotopes of sulfur and their derivatives have
distances in sulfur-iodine charge-transfer complexeszc
been
isolated
and characterized, including S, (n = 6-13,
and are longer than that in [(HzN)zCS]zI+(2.629
18,20,
and
SnO (n = 6-10),28
and SlzOzin
which has a formal S-I bond order of 0.5, but they are
S1202-2
SbC15,46
and
in
various
cations,
for
example,
less than the sum of the van der Waals radii (4.00 A).39
S71+,16
S7Br+,19
(s71)z13+,20
and
(s8)zAg+.47
Selenium
The Sz142+
cation may be regarded as consisting of
forms48aonly the unstable rings Seg,48bSe6,48cand
S2@+and 218%' units, weakly bonded together via two
Se7,48d3e349
in addition to polymeric gray selenium. Demutually perpendicular sets of P*-P* orbitals (Figure
rivatives
of
selenium rings had not been reported prior
5) by electrons in P* orbitals. Thus, P bonding in the
to
our
work.
(S71)4(S4)(A~F6)620~23
has the greatest
cation is maximized. The equidistributions of charge
thermal
stability,
of
the
salts
containing
sulfur-iodine
over all three dimer units (S2.66+,2I,O."+) and the recations,
and
therefore
we
attempted
to
prepare (Se7sulting bonding situation may arise from the near
I)4Se4(AsF6)6.50
Selenium
and
iodine
were
reacted with
equality of the ionization energies of Sz (9.36 eV) and
11
but
by using
AsF,
in
liquid
AsF3
as
indicated
in
eq
Iz (9.3995 eV).40 Therefore, Sz142+
is an example, par
selenium
instead
of
sulfur.
However,
the
reaction
excellence, of a thermodynamically stable species that
15,
and
subsequently
Se61proceeded
according
to
eq
contains a npr-npr (n I 3) bond.7
Preparation and Structure of S e z 1 4 ( S b z F 1 1 ) 2 . It
was postulated that the structure and bonding in Sz12+
(41) Nandana, W. A. S.; Passmore, J.; White, P. S.; Wong, C.-M. J.
were a consequence of the near equality of the ionizaChem. SOC.,Chem. Commun. 1982, 1098.
tion energies (IEs) of Sz and Iz. To test this hypothesis,
(42) Dunitz, J. D. Acta Crystallogr. 1956,9, 579. Kiers, C. Th.; Vos,
A. Acta Crystallogr. 1978, B34, 1499.
we attempted the synthesis of Sez12+(IE of Sez = 8.33
(43) Baudler, M.; Stassen, H.-J. 2.Anorg. Allg. Chem. 1966,343,244.
eV),40 and subsequently Se214(Sb2F11)2
was prepared
(44) Steudel, R. Top. Curr. Chem. 1982,102,149. Steidel, J.; Steudel,
according to eq 14 and its structure determined.4l The
R. J. Chem. Soc., Chem. Commun. 1982, 1312.
d
(34) Leung, Y. C.; Waser, J. J . Phys. Chem. 1956, 60, 539.
(35) Fink, E. H.; Kruse, H.; Rameay, D. A. J . Mol. Spectrosc. 1986,
119. 337.
(36) Karle, I. L. J. Chem. Phys. 1955, 23, 1739.
(37) Davies, C. G.; Gillespie, R. J.; Ireland, P. R.; Sowa, J. M. Can. J.
Chem. 1974,52, 2048.
(38) Lin, G. H.-Y.;Hope, H. Acta Crystallogr. 1972, B28, 643.
(39) Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry,4th
ed.; Interscience: 1982.
(40) Rosenstock, H. M.; Droxl, K.; Steiner, B. W.; Herron, J. T. J.
Phys. Chem. R e f . Data 1977, Suppl. 1 , 6 .
~~
(45) Steudel, R. Comments Znorg. Chem. 1982,1, 313.
(46) Steudel, R.; Steidel, J.; Pickardt, J. Angew. Chem., Znt. Ed. Engl.
1980, 19, 325.
(47) Roesky, H. W.; Thomas, M.; Schimkowiak, J.; Jones, P. G.; Pinkert, W.; Sheldrick, G. M. J. Chem. SOC.,Chem. Commun. 1982, 895.
(48) (a) Cherin, P.; Unger, P. Inorg. Chem. 1967,6,1589. (b)Cherin,
P.; Unger, P. Acta Crystallogr. 1972, B28, 313. Marsh, R. E.; Pauling,
L.; McCullough, J. D. Acta Crystallogr. 1953,6, 71. Foss, 0.;Janickis,
V. J. Chem. SOC.,Dalton Trans. 1980, 624. (c) Miyamoto, Y. Jpn. J.
Appl. Phys. 1980, 19, 1813. (d) Steudel, R.; Strauss, E. M. 2.Naturforsch. 1981,36B, 146. (e) Steudel, R.; Papavassiliou, M.; S t r a w , E.-M.;
Laitinen, R. Angew. Chem., Znt. Ed. Engl. 1986, 25, 99.
(49) Steudel, R.; Papavassiliou, M.; Krampe, W. Polyhedron 1987,
7(7), 581.
(50) Nandana, W. A. S.; Passmore, J.; White, P. S. J. Chem. Soc.,
Chem. Commun. 1983, 526.
238 Acc. Chem. Res., Vol. 22, No. 7, 1989
a
0
L
4
Klapotke and Passmore
Figure 9. View of the polymeric cations (SesI+), in (SesI),.n( AsF,),
representing weak intercationic selenium-selenium
contacts (3.591 (3) A).
0*
Se
.=I
Figure 10. Se81,2+ cation in SesIz(AsFs)z.2SOz.
AsF6 was prepared quantitatively according to eq 16 in
liquid SOz solution.50 Large crystals of Se61AsF6can
3 months
32Se + 21z + 9AsF5
~ S ~ G I A S+FSe8(AsF6):,
G
+ 3AsF3 (15)
12Se + I2 + 3AsF5
1 week
2Se&'k#6
+ AsF,
(16)
be prepared by thermally cycling the reaction mixture.
The crystals appeared ruby red in transmitted light and
have a golden appearance in reflected light.
The successful synthesis of the polymeric (Se61),n(Ad?,) suggested that Se612(AsF6)2
might be preparable.
Subsequently, selenium and iodine were reacted with
AsF5 in liquid SOz according to eq 17.51
6Se + Iz i- 3AsF,
S0,(1)
to that in [Phzses2'],53 which has an Se6ring with a boat
conformation. The iodine substituents are in the 1,4
axial positions.
The tricoordinate selenium atoms in the Se6IZ2+
cation are positively charged, but there is delocalization
of charge into the ring, resulting in bond alternation51
(2.482 (2) A, 2.227 (2) A) and the formation of 4pv4pn
bonds. Each of the two iodine atoms makes two intraionic contacts with both the dicoordinate, but partially charged, selenium atoms within the ring and the
contacts (3.719 (2) A and 3.709 (2) A) that are substantially shorter than the sum of the van der Waals
radii of Se and I (4.15 A). Thus the Se6IZ2+
has a definite distorted cube cluster-like geometry, which it
probably retains in solution (see below).
Identification of Se412+and Se612+by 77SeNMR
Spectroscopy in Solution. The characterization of
the sulfur-iodine cations is seriously hindered by the
lack of a suitable spectroscopic technique, but 77Se
NMR can be used in the selenium system. We therefore systematically searched for selenium-iodine cations
in SOz solution, by natural abundance 77SeNMR. As
part of this investigation, we followed the reaction of
Se42fwith varying amounts of I2 and found that the
equation proceeded according to eq
Se4(AsF6)z 212 --* Se414(ASF&
(18)
The major features of the spectrum were two peaks
of equal intensity and satellite peaks showing TeJ7Se
couplings consistent with an AXX'A' spectrum. These
data are consistent with an IzSe+SeSeSe+12
formulation
for the cation. In addition to the two major peaks attributable to Se412+,there were three other less intense
peaks present, one attributable to Se13+and the other
two due to Se612+,whose integrated areas were in a
ratio of 1:2. The latter two peaks have satellites due
to 77Se-77Secouplings consistent with an AX2X2/Af
spectrum and therefore attributable to a symmetric
Se6IzZ+.Se4142+was shown to be in equilibrium with
%I3+ and Se612+according to eq 19, and AHo and ASo
have been estimated to be 20 kJ mol-' and 60 J K-'
mol-1.55
SeGIz(AsF6)z-2S0z
+ AsF3
(17)
The 77SeNMR of this solution showed the presence
of about 11 different selenium cations (see below), and
many attempts to produce crystals at room temperature
2se4142++ 2Se13++ Se6IZ2+
(19)
were unsuccessful. However, when the solution was
cooled to -70 "C for 1 h and left at room temperature,
The 77SeNMR spectra of solutions of Se612(AsF6)2
at
then 80% of the selenium crystallized out as highly
various
temperatures
show
that
Se6IZ2+
itself
is
in
crystalline Se61z(AsF6)z~2S0z.51~52
with SeJ+ and Se4142+
according to eq 20.62
Structure O f (Se61)n.n(ASF,) and S ~ ~ I Z ( A S F ~ ) Zequilibrium
-~so2.The structure of (Se61),.n(AsF6)consists of AsFG2Se612+* SeS2++ Se41d2+
(20)
anions and polymeric strands of [Se61+],cations with
It is also in equilibrium with several other species, which
some cation-cation and cation-anion interactions50
are presently under investigation. Se412+and Se612+
(Figure 9). The cation contains hexaselenium rings in
undergo
Se+-I and Se-Se bond redistribution reactions,
a chair conformation similar to that of cyclohexaand since the various combinations do not differ greatly
selenium.48c The rings are joined to two neighboring
in enthalpy, the formation of the large number of
hexaselenium rings by two weak (2.736 (3) A) exocyclic
species
is probably entropy driven. It is likely that the
1,4 diaxial Se-I bonds (Figure 9) of bond order ca. 0.5;
sulfur-iodine
cations also give complex equilibrium
each tricoordinate Se atom carries a charge of 0.5. The
mixtures in solution.
[Se6I+Incation was the first example of a derivative of
Chloro a n d Bromo Cations of S u l f u r and Selea selenium ring; and it is also polymeric, unlike the
nium.
All binary chalcogen-chlorine cations of the type
known sulfur-iodine cations.
XC13+ (X = S, Se, Te) have been prepared, and the
The discrete centrosymmetric Se6IZ2+
cation contains
hexaselenium rings of chair conformation (Figure 10)
(53) Faggiani, R.; Gillespie, R. J.; Kolis, J. W. J . Chem. Soc., Chem.
similar to those in Se648b
and [Se61+],,wbut in contrast
Commun. 1987, 592.
(51) Passmore, J.; White, P. S.;Wong, C.-M. J . Chem. SOC.,Chem.
Commun. 1985, 1178.
(52) Wong, C.-M. Ph.D. Thesis, University of New Brunswick, 1988.
(54) Burns, R. C.; Chan, W. L.; Gillespie, R. J.; Luk, W. C.; Sawyer,
J. F.; Slim, D. R. Inorg. Chem. 1980, 19, 1432.
(55) Carnell, M. M.; Grein, F.; Murchie, M. P.; Passmore, J.; Wong,
C.-M. J . Chem. Soc., Chem. Commun. 1986, 225.
Ace. Chem. Res., Vol. 22, No. 7, 1989 239
Sulfur and Selenium Iodine Compounds
3Br2
- i7 I- P
O = s e
6/gs8(Or
-
6Se) -k 3AsF5
Br2XXXBr(AsF6)+ AsF3 (23)
less than 1 (Se(l)-Se(2), 2.554 (6) A) to ca. 1.5 (Se(2)-Se(3), 2.211 (6) A). This implies substantial 4pn4pa bonding between Se(2) and Se(3) and, in valencebond terms, suggests that the bonding may be represented by valence-bond structures C and D.
Figure 11. Se9C1+cation in Se9C1SbCls.
Br
Br
I
Se+-Se
/
/
Se
Br
I
Br
Figure 12. Se2Br6+cation in Se2Br5AsFs.
C
I
s" 8"
Br
Se*
I
Br
D
The cluster-like geometry of this cation maximizes
intracationic contacts, charge delocalization,the number
of Se+-Br bonds (cf. MeSeSe+(Me)SeMe)61,Se-Se bond
alternation, and P bonding.
(56) Edwards, A. J. J. Chem. SOC.,Dalton Trans. 1978, 1723.
(57) Stork-Blaisee, B. A.; Romers, C. Acta Crystallogr. 1971, B27,386.
(58) Faggiani, R.; Gillespie, R. J.; Kolis, J. W.; Malhotra, K. C. J.
Chem. Soc., Chem. Commun. 1987,591.
(59) Murchie, M. P.; Passmore, J.; White, P. S. Can. J . Chem. 1987,
Conclusions
A New Class of Compounds Discovered: The
Chalcogen Iodine (and Bromine) Cations. As a result of our unsuccessful attempts to prepare S13(MF6),
a large number of novel, stable sulfur and selenium
iodine and bromine cations have been prepared quantitatively and their structures have been determined
(Table I). Thus, whereas stable neutral sulfur iodides
and selenium iodides either are not known (selenium)
or can only be prepared at low temperatures (sulfur),
sulfur-iodine and selenium-iodine cations have been
shown to be unexpectedly numerous. These simple
cations have novel structures and bonding arrangements, and it may be argued that the goals outlined at
the beginning of this Account have been achieved.
The Stable >X+-I (X = S, Se) Bond. The crystal
lattice component is likely not the only factor responsible for the stability of the X+-I bonds in the salts
described in this Account, as (CH3)2SISbF6,(CH3)2SISbC16, and (CH3)(CF3)SIMF632
are stable only at low
temperatures. The S+-I (Br) and Se+-I (Br) bond
distances in the more complex chalcogen halide cations
are all slightly shorter than the sums of the corresponding covalent radii and the corresponding bond
lengths in neutral S(I1)-I (Br) and Se(I1)-I (Br) containing compounds.lOcJ1They are also shorter than the
observed or predicted8aX-Hal bond lengths in XHa13+
(X = S, Se; Hal = I, Br).8a The S+-I (Br) and Se+-I
(Br) bonds in the chalogen halide cations are therefore
presumably stronger than those in corresponding neutral compounds, or simple MHa13+salts and their simple derivatives. In addition, the more complex chalcogen halide cations, as a whole, are probably stabilized
by charge delocalization, bond alternation, and halogen-chalcogen intercationic contacts, as a consequence
of the presence of X+-I (Br) bonds (much less extensive,
or not possible, in Se13MF6,(CH3)2SISbF6,(CHJ2SISbC16, and (CH3)(CF3)SIMF632).Neutral sulfur and
selenium bromides and chlorides are stable, consistently; XhHal n+ cations containing an Se-C1 bond (in
Se7+SeSeC1)&and X-Br bond (in Br2X+XXBr,X = S,
Se)60are observed whereas iodine-containing analogues
are not.
(60)Passmore, J.; Tajik, M.; White, P. S. J . Chem. SOC.,Chem. Commun. 1988,175.
(61) Laitinen, R.; Steudel, R.; Weiss, R. J . Chem. Soc., Dalton Trans.
1986, 1095.
0.S e
@ = 8r
Figure 13. SeSBr3+cation in Se3Br3AsF6.
structures of the SC13+&and SeC13+57cations are similar to that of Se13+(Figure 1). Gillespie et al. prepared
and characterized by X-ray crystallography the Se7+Se2Clcation in Se9C1SbC&.58This compound contains
the first example of a structurally characterized seven-membered selenium ring. The chair conformation
of the Se7 ring with the Se2Cl in the endo position is
shown in Figure 11. The geometry of the Se7+SeSeC1
is similar to that of Br2Se+SeSeBr(see below) with the
Se7+replacing the Br2Se+in Br2Se+SeSeBr.
The bromine cations SBr3+
and SeBr3+l k p d are
also of interest as these are the simplest binary bromine
cations which contain the X(IV)+-Br (X = S, Se) bond.
The X-ray Crystal StI'UctureSOf SBr3MF6 and SeBr3MF6
confirm the identity of the pyramidal cationssa (Figure
1). S7BrMF6 and ( S ~ B ~ ) ~ S ~ ( Awere
S F S prepared
)~
quantitatively by routes similar to those of the iodine
counterpart^^^ (see above). The structure of S7Br+in
(S7Br)4S4(ASF6)6is very similar to that of s71+
(Figure
2).
The first example of an X2Ha15+species (X = chalcogen) is the cation Se2Br5+,which was prepared
quantitatively according to eq 21 and 22.59 The Se2Br5+
4Se + 5Br2 + 3AsF5 2Se2Br5AsF6+ AsF,
(21)
8a~14aib
-+ -
Se4(AsF6)2 5Br2
2Se2Br5AsF6
(22)
cation (C2,, symmetry) contains two trans SeBr2units,
linked by a bridging bromine atom at an inversion
center (Figure 12). Recently Br2X+XXBr(AsF6-)[X
= S, Se] have been prepared quantitatively according
to eq@
,3'2
containing the Br2X+XXBrcation (Figure
13). The Se-Se bond orders vary from substantially
65,1584.
240 Ace. Chem. Res., Vol. 22, No. 7, 1989
Klapotke and Passmore
Table I1
M-M Bond Distances and Bond Orders
M-M shortest
bond order
cation
S7Br+
S2142+
S71+
(s71)z13+
Se9Cl+
Se3Br3+
Se21t+
SeJ+
Se6IZ2+
bond distance,
A
1.92 (2)
1.828 (1)
1.906 (5)
1.897 (10)
2.223 (5)
2.211 (6)
2.841 (2)
2.292 (4)
2.227 (2)
(ref 31)
1.7
2.3
1.8
1.8
1.5
1.6
0.2
1.2
1.5
Examples of Thermodynamically Stable n pan pa (n L 3) Bonds. Charge delocalization and bond
alternation lead to the presence of particularly short
S e s e and S-S bonds in the homopolyatomic sulfur and
selenium halogen cations (Table 11)except for Se2142+.
Thus these cations can be regarded as containing examples of thermodynamically stable 3pa-3pa bonds
S2142+
is particularly remarkable
and 4 p t - 4 ~bonds.
~
in that it maximizes a bonding and contains 3u and
3 n p r n p t bonds (n 1 3) (Figure 5). The bond order
of the S2unit corresponds to 2.33, the highest observed
for any isolated non second row element containing
compound. It is thermodynamically stable with respect
to an all u bonded isomer and also with respect to addition of 12(s)(eq 24).24 Se2142+
is also thermodynamS ~ I ~ ( A S F J ~+( SI )~ ( s ) ~SI~ASFG(S)(24)
ically stable with respect to an all u bonded isomer. It
consists of two Se12+units, joined by a weak t*-a*
interaction (Figures 6 4 , and the dimer contains one
4pa-5pt bond delocalized over the four Se-I bonds. A
similar situation is found for many homopolyatomic
chalcogen and halogen cations (e.g., X42+,X = S, Se,
Te; Hal2+,Hal = Br, I, 142+)7
which also contain thermodynamically stable npa-npa (n I3) bonds. This is
in contrast with the neutral group 14 and 15 compounds
containing npn-npt (n I3) bonds which are kinetically, but not thermodynamically, stable.I2 This is in
part because there is charge localization on adjacent
positively charged atoms in the alternative a-bonded
isomer [12X+X+12](X = S, Se). The S+-S+ u bond
dissociation energy will be significantly less than that
in a normal sulfur-sulfur bond, and in addition, there
will be an energy loss accompanying charge localization.
Thus the energetics of the u versus a bonds in the
cations are very different from those in neutral molecules.62
Presence of T*--?T*Bonds in S2142+and Se2142+and
Its Implications for Bonding in Related Species.
The dimers in S2142+
are joined by two weak naturally
perpendicular sets of a*-t* bonds (Figure 5). This
situation is similar to that found in Id2+and in (NO),
dimers in the solid state, and in a variety of other
+
(62) Schmidt, M. W.; Truong, Phi. N.; Gordon, M. S. J. Am. Chem.
SOC.1987, 109, 5217.
0 .Se
e . I
Figure 14. The structure of SeBzt compared with that of Sez142+.
Figure 15. T h e HOMO-1 of S2'.
sulfur-containing dimer^.^,^^ Similarly, the two Se12+
units are joined via a six-center two-electron t*-t*
bond (Figure 8), similar to the bonding in S2042-.42
The long Se-Se bonds in Se2142+(2.841 A) and Se82+
(2.83 A)7 are similar in length. In addition, both Se21,2+
and, in Set+,the two tricoordinate formally positively
charged selenium atoms and the four selenium atoms
to which they are joined have the same eclipsed geometries (Figure 14). Thus the six selenium atoms in Se+
:
are also joined by a six-center two-electron a*-a* bond.
The geometry of sg2+ is also similar, and the HOMO-1
has been shownMto have a*-a* characteristics (Figure
15). Therefore, the presence of a*-a* bonds of both
the four-center two-electron type and the six-center
two-electron type are likely to be found in a variety of
compounds of the electron-rich elements (e.g., in S4N4,
which may be viewed as containing a six-center twoelectron bond about each of the two S-S interactions).
Where they have been measured, the strengths of the
t*-t*bonds are weak (less than 40 k J
In
contrast, a high bond energy is associated with t
bonding within the monomer. Thus the a*-a* bond
formation follows that of the t-bonded fragments.
This Account would not have been possible without the fine,
dedicated experimental work of former graduate students and
postdoctoral fellows (whose names are given in the references),
but especially Dr. Peter Taylor, whose outstanding work opened
u p this field, and Dr. Peter White, for the determination of many
X - r a y structures, which have often been complicated by MF6disorder problems. W e thank Dr. Neil Burford for many enjoyable discussions and ideas on bonding in related cations (see
ref 7), U N B and N S E R C (Canada) for financial support, and
the Humboldt Foundaton for a L y n e n Fellowship (T.K.). W e
also thank Simon Parsons for his help i n improuing the manuscript.
(63) Banister, A. J.; Clarke, H. Y.; Rayment, I.; Shearer, H. M. M.
Inorg. Nucl. Chem. Lett. 1974,10,647. Awere, E. G.; Burford, N.; Mailer,
C.; Passmore, J.; Schriver, M. J.; White, P. S.; Banister, A. J.; Oberhammer, H.; Sutcliffe, L. H. J. Chem. Soc., Chem. Commun. 1987, 66.
(64) Burford, N., private communication.
Vol. 12, 1979
Chemistry of Electronegative Elements
211
Recent Developments in the Chemistry of Some
Electronegative Elements
KONRAD
SEPPELT
Arzorganisch-Chemisches Institut der Uniuersitat, 0-6900 Heidelberg, West Germany
Received August 18, 1978
In 1962 a frontier of chemistry was opened by the
discovery of noble gas compounds.l Since then this
field has experienced vast development, and many
reviews cover the first ten hectic years of worka2
Soon it became obvious that noble gas chemistry had
to be seen in close relationship to halogen chemistry and
to the chemistry of fluorides and oxyfluorides of certain
other, heavier nonmetal elements such as selenium,
tellurium, and antimony, especially in their higher
oxidation states. Although the total number of original
papers dealing with the chemistry and structure of
nonmetal fluorides and related species exceeds several
thousand, there is still no review available that covers
the complete literature. But the structural principles
of those compounds are now well-known, especially if
the coordination number six is not exceededa3
Xenon hexafluoride is one of the molecules with a
formally higher coordination number, if one counts the
nonbonding electron pair. Thus, its structure has
caused much confusion; the most recent theory, based
on physical measurements, is presented here. A more
chemical way of solving such a problem would be to
prepare other hexavalent xenon compounds like XeL6
and to investigate their structures. Out of many choices
of suitable ligands L, only -OTeF5 and perhaps -OSeF5
have been found to be appropriate. Fluorine stabilizes
high oxidation states (PtF6,XeF6, IF7,ReF7) for several
reasons: its small size and strong electron withdrawal
produce a strong, partly ionic bond. This effect can be
achieved by the ligand -OTeF5 as well. And fluorine
is not likely to be eliminated as F2,since the dissociation
energy of Fz is very small. On the other hand, the
ligand dimers F5TeOOTeF5,F5SeOOSeF5,FS0200SOzF, and others are readily formed, as the peroxide
bonds are quite strong. Thus it was to be expected that
no ligand would be better than fluorine, but research
in this direction has produced much new information.
The Xenon Hexafluoride Structure Problem
It is now well established that XeF6 is the highest
fluoride of ~ e n o n . ~In, ~contrast to XeF2 and XeF4, its
structure has prompted 15 years of discussion, but is
now almost clear. The fundamental problem must be
separated into several parts, as the structure differs in
all three physical states. This behavior is different from
that of other known hexafluorides, 15 altogether, which
are plainly octahedral.6
In the gaseous state XeF6 is monomeric7 but not
octahedral, though the deviation from octahedral
Konrad Seppelt is Associate Professor of Chemistry at the University of
Heidelberg. He was born in Leipzig, Germany, in 1944. Following undergraduate
studies at the University of Hamburg, h e went on to graduate work at Heidelberg,
where he received the Ph.D. degree in 1970. His research interests are In the
synthetic chemistry of small molecules and in the main group chemistry of the
chalcogens, halogens, and noble gases.
symmetry is smaller than predicted by the VSEPR
model.* The structure is best described in terms of a
mobile electron pair that moves over the faces and edges
of the octahedron and thus distorts it in a dynamic
manner.g From its vapor pressure it can be concluded
that XeF6 is associated in the condensed phase. The
association in nonpolar solvents like F5SOSF5,n-C5F12,
CF2Cl2,and S02C1F has been verified by 19Fand 129Xe
nuclear magnetic resonance. The latter NMR method
was first used only recently, but is now the best analytical probe for any kind of xenon chemistry in solution. 129Xehas a natural abundance of 26.4% (xenon
enriched up to 60% is commercially available), a nuclear spin of 1/2, and a gyromagnetic factor very close
to that of 13C. The 129XeNMR spectra of many xenon
compounds have now been observed.’@14 The results
of XeF6 in solution’’ a t temperatures below -100 “ C
have caused much skepticism; however, repetition of the
measurements has not allowed any conclusions other
than those derived from the first NMR spectra.12J3
These spectra (Figure 1) have to be interpreted in
terms of a tetramerization of XeF6 in which both the
xenon and the fluorine atoms are magnetically
equivalent. Each xenon atom couples to 24 equivalent
fluorine atoms and each fluorine atom to 4 equivalent
xenon atoms. Thus, the fluorine atoms in the cluster
Xe4F24are bound in a nonrigid manner and are
equilibrated by a scrambling mechanism that is best
named a cogwheel mechanism (Figure 2). Attempts
to freeze out the fluorine migration have been made but
could not be confirmed.lZJ4
Undoubtedly this cluster has the largest known
number of nonrigid bonded ligands, namely 24. The
low-temperature configuration of Xe4F24is somewhat
like the high-temperature form of Rh4(CO)12.15J6 In
(1) N. Bartlett, Proc. Chem. Soc. London, 218 (1962).
(2) See, e.g., N. Bartlett and F. 0. Sladky, “The Chemistry of Krypton,
Xenon and Radon”, in “Comprehensive Inorganic Chemistry”, Vol. 1,
Pergamon Press, Oxford-New York, 1976 p 213.
(3) K. Seppelt, Angew. Chem., Int. Ed. Engl., 18, 181 (1979).
(4) B. Weinstock, E. E. Weaver, and C. P. Knop, Inorg. Chem., 5,2189
(1966).
(5) One can estimate that a xenon octafluoride would be thermodynamically unstable toward loss of fluorine atoms. This would indicate
kinetic instability, except at cryogenic temperatures.
(6) SF,, SeF,, TeF,, MoF,, WF6, TcF,, ReF,, RuF,, OsF,, RhF6, IrF6,
PtF6, UF,, NpFcj, PUF,.
(7) W. E. Falconer, M. J. Vasile, and F. S. Stevie, J . Chem. Phys., 66,
5335 (1977).
(8) J. Gillespie and R. S. Nyholm, Q. Reu. Chem. SOC.,11,339 (1957).
(9) K. S. Pitzer and C. S. Bernstein, J . Chem. Phys., 63, 3849 (1975);
U. Nielsen, R. Haensel, and W. H. E. Schwarz, ibid., 61, 3581 (1974); S.
Y. Wang and C. C. Lohr, Jr., ibid., 60, 3901 (1974); 61, 4110 (1974).
(10) K. Seppelt and H. H. Rupp, Z. Anorg. Allg. Chem., 409, 331,338
(1974).
(11)H. H. Rupp and K. Seppelt, Angew. Chem.,Int. Ed. Engl., 13,613
(1974).
(12) K. Seppelt and N. Bartlett, 2. Anorg. Allg. Chem., 436,122 (1977).
(13) G. J. Schrobilgen, J. H. Holloway, and P. Granger, J . Magn.
Resonance, in press.
(14) G. J. Schrobilgen, J. H. Holloway, P. Granger, and C. Brevard,
Inorg. Chem., 17, 980 (1978).
0001-4842/ 79/01 12-0211$01.00/ 0 0 1979 American Chemical Society
Seppe 1t
212
I
Accounts of Chemical Research
15
I
0
1
3
Y
325 Hz
Figure 1. (Above) 19FNMR spectrum of XeF, in CF3C1at -130
"C (natural isotopic abundance, lZ9Xe= 26.4%). (Middle) '$F
NMR spectrum of XeF, in CF3C1at -130 OC with an enriched
sample of 12$Xe(62.5%). (Bottom) lZ9XeNMR spectrum in
CF2Cl2at -120 "C. The interpretation of these spectra is shown
on the right-hand side. Under the assumption of a tetramerization
to Xe4F, and the equivalency of four xenon and 24 fluorine atoms,
five different species are in solution: '29Xe4FZ4,
lZ9Xe3XeFZ4,
'29(e2XeZFu,lBXeXe3FZ,and Xe4F,. As only
has a nuclear
spin of l/z, the first species shows in the '$F NMR a 1:4:6:4:1
quintet (Qt), the second a 1:3:3:1 quartet (Q), the third a 1:2:1
triplet (TI, the fourth a 1:l doublet (D), and the fifth a singlet
(S). Many of those lines overlap within the line width, and the
intensity of the resulting nine lines depends upon the '*$Xe
concentration. In the case of 62.5% the line intensities are
(calculated) 3.8:18.8:49.8:84.5:100:85.5:49.818.3:3.8 and (observed)
3.1:18.1:48.883.2:10083.450.1:17.8:2.9. The lZ9Xespectrum shows
a binominal function of 24th order: (calculated down from the
center line) 100:92:72.5:48.3:27.2:12.8:4.9:1.6 ...; (observed)
100:93.3:68.8:45,8:29.1:14.5:4.1
... The fit between experimental
and calculated lines is striking. As cannot be shown here, no other
kind of multiplicity can explain these spectra.
the dynamic rhodium carbonyl there are no problems
concerning the binding forces of such a cluster, as there
are bonds between the metal atoms, in addition to
mobile carbonyl bridges. In Xe4F24,however, the mobile
fluorine bridges alone are responsible for the existence
of the cluster. While any xenon-xenon interaction is
highly speculative,ls one can, just by counting, find
enough orbitals and electrons to form four three-center
two-electron bonds, each on one plane. A comparable
bonding model is that of B4C14. Certainly the general
chemistry of xenon does not call for electron-deficient
bonds, though the charge of the xenon in XeF6 is found
indeed to be some 1.5 positive units.ls Attempts have
failed to measure the Xe-Xe distances in the cluster
in solution, mainly because of experimental difficulties.
I t is interesting to compare the structure of the
cluster in solution with the solid-state structure of XeF@
The cubic modification is the only one (out of four) that
has been fully ana1y~ed.l~
The unit cell contains 24
tetrahedral (XeF,+F-), and 8 octahedral (XeF5+F-)6
units. The tetrahedras are comparable to the dynamic
cluster in the dissolved state. However, the Xe-Xe
distances are quite large (4.2 A) here and do not favor
(15) F. A. Cotton, C. Kruczcynski, and B. C. Shapiro, J . Am. Chem.
SOC.,94, 6191 (1972).
(16) J. Evans, B. F. G. Johnson, J. Lewis, J. R. Norton, and F. A. Cotton,
J . Chem. SOC.,Chem. Commun., 807 (1973).
(17) The very new Xe2+ion is the first one to have an unquestionable
xenon-xenon bond: L. Stein, R. Norris, A. J. Downs, and A. R. Minihan,
J. Chem. SOC.,Chem. Commun., 502 (1978).
(18) T. X. Carroll, R. W. Shaw, T. D. Thomas, C. Kindle, and N. Bartlett,
J . Am. Chem. SOC.,96, 1989 (1974)
0
l
Figure 2. The cluster Xe4FZ4in two of its configurations. In the
first configuration each xenon holds six fluorine atoms, and in
the second four F atoms bind the Xe atoms by Xe-F-Xe bridges.
The latter picture resembles very much the (XeFS+F-)4tetramer
in the solid state.lg All fluorine atoms take part in this dynamic
process, which is best explained as a cogwheel mechanism.
(Reprinted with permission from ref 10, 11, and 13. Copyright
1974, Johann Ambrosius Barth.)
the speculative bonding model in solution. On the other
hand, in the crystal the fluorine atoms are in fixed
positions.
For a simple, binary compound the structural behavior of XeFs is without parallel. The species
chemically most similar in behavior is SeF4,because it
has (including one electron pair) an odd number of
ligands, is highly associated in the liquid phase, and
easily forms SeF3+and SeF,- ions, just as XeF, forms
XeF6+,XeFa2-,and XeF7-. However, even a t -140 " C
SeF, is monomeric in solution, and a freezing-out effect
occurs in the fluorine NMR spectrum.20
Xe(OSeF5)2,Xe(OTeF5)2,Xe(OTeF5)r,
O=Xe(OTeF5)4, Xe(OTeF&
There has been hope that ligands other than fluorine
could form single covalent bonds to xenon. Indeed, the
(19) R. D. Burbank and G. R. Jones, J . Am. Chem. SOC.,96,43 (1974).
(20) K. Seppelt, 2. Anorg. Allg. Chem., 416, 12 (1975).
Chemistry of Electronegative Elements
Vol. 12, 1979
Figure 3. Crystal structure of Xe(OSeF&. This structure
determination is complicated by a threefold disorder along the
Se-Xe-Se axis.
-
first ligand of this type was fluorosulfate.21 However,
the compounds FXeOSOzF and Xe(OS02F)2decompose a t room temperature or slightly above, and the
fluorosulfate of XeF6 turned out t o be ionic:
XeF5+S03F-.zz The situation changed when the
substances Xe(OTeF5), and Xe(OSeF5)zwere made.z3t24
These do not decompose below 100 "C, and their covalent character is well established by their lz9XeNMR
spectralo and their crystal structuresz5(Figure 3). Once
it was shown that the ligands -OSeF5 and -OTeF5 have
the capacity of stabilizing high oxidation states, almost
like fluorine, a large amount of their chemistry was
studied (see below). The process of getting higher
valent xenon linked to these ligands turned out to be
very difficult and was successful only in the case of the
-OTeF5 ligand. Only recently Xe(OTeFJ4, Xe(OTeF,),,
and O=Xe(OTeFj)4 as stable, oxygen single-bonded
xenon(1V) and xenon(V1) compounds have been made
for the first time.26
+
XeF6 + 2B(OTeF5)3
2BF3 + Xe(OTeF5)6,decomp -20
"c
,
213
-~ -
~~~
..~ ______..
.
2 3 2 > C l 5s:
Figure 4. lZ9XeNMR spectrum of O=Xe(OTeF,),, solution in
C2F4ClZ,55.33 MHz. The chemical shift of fi = -5121 against
atomic xenon is typical for six-valent xenon. The splitting into
quintets ( J x ~=- 4~ Hz), 13 lines out of a multiplet of 17 (JXe-F
= 55 Hz), and the satellites (Jx,125Te = 1281 Hz) give an almost
complete structure description. (Reprinted with permission from
ref 26. Copyright 1978, Verlag Chemie GMRH.)
be converted into Kr(OTeF5)2. Only the decomposition
products Kr and F5TeOOTeF5 (F5SeOOSeF5) were
observed, indicating the intermediate formation of an
oxygen-bonded krypton compound, as these peroxides
are the typical decomposition products of all xenon
compounds as well. Thus KrFz and its ionic derivatives
KrF+ and Kr2F3+remain the only krypton compounds
known.z8 On the other hand, xenon is known to bind
even to nitrogen in FXeN(S02F)2.z9 XeC12 can be
made only a t cryogenic temperature^.^^ Despite many
attempts, a xenon-carbon bond has not yet been
achieved.
The lz9XeNMR of O=Xe(OTeF5)4 is shown in Figure
4. I t gives a full structural description of the molecule
by the coupling to axial and equatorial fluorine atoms
and to the lZ5Teisotopes.
The most interesting species, Xe(OTeF5)s,is a dark
red, almost violet crystalline solid. The deep color
indicates that the molecule is monomeric, like the
yellow XeF6. The color has to do with the nonbonding
electron pair on xenon, as Te(OTeF5)6lacks this
electron pair and is colorless.z7 Due to the sensitivity
of Xe(OTeF5)6 to light and elevated temperatures,
structural investigations are difficult but are in progress.
The ligands -OSeF5 and -OTeF6 were not found to
form stable compounds with krypton. KrF2 could not
Group Electronegativity of the -OSeF5 and
-OTeF5 Groups
The reason why the -OSeF5 and -OTeF5 groups form
the most stable xenon compounds after the simple
fluorides has not been fully answered. One may define
group electronegativities for these ligands. Certainly
electronegativity has no clear meaning, as there are
several definitions. The concept becomes even more
confusing if electronegativity is extended to whole
groups. Even so, a qualitative argument is possible with
the valence shell electron pair repulsion model.* This
predicts that in a trigonal bipyramid the two axial
positions are occupied by the more electronegative atom
or ligand, whereas in the case of a square pyramid of
the IF, type, the equatorial positions are occupied by
the more electronegative ligands. There are numerous
examples for this substitutional behavior in case of the
trigonal bipyramid. A substitution with -OSeF5 or
-OTeF5 on a trigonal bipyramid has not yet been
achieved. However, we have replaced fluorine atoms
of IF5 with those ligands. All other groups like -OCH3
or -CF3 on IF, indeed prefer the axial p o ~ i t i o n , as
~~-~~
(21) N. Bartlett, M. Wechsberg, F. 0. Sladky, P. A. Bullinger, G. R.
Jones, and R. D. Burbank, Chem. Commun., 703 (1969).
(22) D. D. DesMarteau and M. Eisenberg, Inorg. Chem., 11,2641 (1972).
(23) F. Sladky, Angew. Chem. Int. Ed. Engl., 8,523 (1969); F. Sladky,
Monatsh. Chem., 101, 1559 (1970).
(24) K. Seppelt, Angew. Chem.,Int. Ed. Engl., 11,723 (1972);K. Seppelt
and D. Nothe, Inorg. Chem., 12, 2727 (1973).
(25) L. K. Templeton, D. H. Templeton, N. Bartlett, and K. Seppelt,
Inorg. Chem., 15, 2718 (1976).
(26) D. Lentz and K. Seppelt, Angew. Chem., Int. E d . Engl., 17,356
(1978); D. Lentz and K. Seppelt, ibid., 18, 66 (1929).
(27) D. Lentz, H. Pritzkow, and K. Seppelt, Angew. Chem., Int. Ed.
Engl., 16, 729 (1977); Inorg. Chem., 17, 1926 (1978).
(28) H. Selig and R. D. Peacock, J . Am. Chem. Soc., 86, 3895 (1964);
D. E. McKee, C. J. Adams, A. Zalkin, and N. Bartlett, J. Chem. Soc., Chem.
Commun., 26 (1973); R. J. Gillespie and G. J. Schrobilgen, ibid., 90 (1974);
B. Frlec and J. H. Holloway, Inorg. Chem., 15, 1263 (1976).
(29) R. D. Le Blond and D. D. DesMarteau, J. Chem. Soc., Chem.
Commun., 555 (1974).
(30) C. Y. Nelson and G. C. Pimentel, Inorg. Chem., 6, 1758 (1967);
W. F. Howard and L. Andrews, J . Am. Chem. SOC.,96, 7864 (1974).
(31) G. Oates and J. M. Winfield, Inorg. Nucl. Chem. Lett., 8, 1093
(1972); J. Chem. Soc., Dalton Trans., 119 (1974).
(32) G. Oates, J. M. Winfeld, and 0.R. Chambers, J. Chem. SOC.,
Dalton
Trans., 1380 (1974).
(33) D. Naumann, M. Schmeisser, and C. Deneken, J . Inorg. Nucl.
Chem., H. H. Hyman Mem. Vol., 13 (1976).
XeF4 + 4 / 3 B(OTeF5)3
4 / 3 BF,
Xe(OTeF5)4,mp 72 "C (dec)
-
XeOF4 + 4/3B(OTeF5)3
4/3BF3+ O=Xe(OTeF5)4, mp 53 "C
S e p p e 1t
214
observed
a x i a l region
Accounts of Chemical Research
equatorial regior
F3i;
not observed
F ~ < F
F
F
F
F
c
'
.
;=:
O
A
O
c
F
010
F
0
o&o
0
0
Figure 5. 19FNMR spectra of F,I(OTeF5)5_,and F,I(OSeF5)5_,,schematically with the tellurium and selenium bonded fluorine atoms
omitted. The assignment is easily made by the appearance of the axial-bonded fluorine atom in a downfield area. The ligands -OTeF,
and -OSeF,, as indicated by 0, do not go into the axial position, where normally the less electronegative ligands have to go. (Reprinted
with permission from ref 35. Copyright 1978, Verlag Chemie GMBH.)
predicted. The surprise is that in the compounds
F,I(OTeF5)5-, and F,I(OSeF5)5-, the substitutional
behavior is inverted.
These species were made according to the equations:
IC& + 3C10TeF5
- +
3Cl2
I(OTeF5)334
+ Xe(OTeF5I2 I(OTeF5)5+ Xe3j
I(OTeF5), + IF5 F,I(OTeF5)5-,35
IF5 + X POF2-OSeF5 F,I(OSeF5)5-,35
I(OTeF&
The conformational analysis of the mixed substituted
molecules was done by the I9F NMR of the iodinebonded fluorine atoms, as shown in Figure 5 . Fluorine
keeps the axial position, meaning that in the sense of
the VSEPR model the ligands -OTeF5 and -OSeF5 are
more electronegative than fluorine. Kinetic factors
cannot explain this abnormal substitutional behavior,
as prolonged reaction time and heating do not change
the results. Nor can it be a steric effect: there is no
reason why F41-OTeF5(eq) should be sterically favored
over F4FOTeF5(ax),and X-ray structures of the closely
related compounds trans-FzTe(OTeF5)t6and T e ( 0 TeF6)627show that there is no important steric hindrance between the quite bulky ligands (Figure 6).
The high formal electronegativities of those groups
are certainly a result of the inductive effect of five
fluorine atoms, possibly enhanced by some (pd),
backbonding between oxygen and selenium or tellurium. The latter effect is supported by the structural
investigation of F5SOSF5,F5SeOSeF5,and F5TeOTeF5.
(34) K. Seppelt, Chem. Ber., 106, 1920 (1973).
(35) D. Lentz and K. Seppelt, Angew. Chem., Int. Ed. Engl., 17, 355
(1978).
-,
\ - -
(36) H. Pritzkow and K. Seppelt, Angew. Chem., Int. Ed. Engl., 1 5 ,
771 (1976); Inorg. Chem., 16, 2685 (1977).
-2
Figure 6. Molecular structure of Te5O4Fzz.Most tellurium(V1)
oxyfluorides have the general formula TenOn-1F4n+2.Here the
species is best formulated as trans-F,Te(OTeF,),. The environment around the tellurium atoms is always octahedral.
(Reprinted from ref 36).
Here indeed a sterically unfavorable eclipsed configuration is maintained, probably by some (pd), double
bonding.37 Once the high electronegativity of these
groups was established, it became clear that the
chemistry of those ligands could become as extensive
as that of fluorine (Table
Only helium,
1233,34338-46).
(37) H. Oberhammer and K. Seppelt, Angew. Chem., Int. Ed. Engl.,
17, 69 (1978); Inorg. Chem., 17, 1435 (1978).
(38) A. Engelbrecht and F. Sladky, Monatsh. Chem., 96 159 (1965).
(39) A. Engelbrecht, W. Loreck, and W. Nehoda, Z. Anorg. Allg. Chem.,
360, 89 (1968).
(40) F. Sladky, H. Kropshofer, and 0. Leitzke, J . Chem. SOC.,Chem.
Commun., 134 (1973).
(41) F. Sladky and H. Kropshofer, J . Chem. SOC.,Chem. Commun.,
fino (1973).
.,
_ _ . / _ .
(42) E. Mayer and F. Sladky, Inorg. Chem., 14, 589 (1975).
(43) K. Seppelt, Chem. Ber., 108, 1823 (1975); 109, 1046 (1976); 110,
1470 (1977).
Chemistry of Electronegative Elements
Vol. 12, 1979
215
h
P
S
Figure 7. I9F NMR spectrum of HOSFSat -60 "C in CH2C1,
solution. This is a typical AB4 pattern. The two strong groups
belong to the B4 part: the other lines represent the A part. The
line S (singlet) represents the exact chemical shift of the A part.
(Reprinted with permission from ref 46. Copyright 1977, Johann
Ambrosius Barth.)
neon, argon, and krypton are definitely not able to be
bound to these groups. The difference from fluorine
is seen in the case of krypton, and in some cases the
extremely high oxidation states, as in IF7 and PtF6,
could not be established with those groups so far.
HOSF5, HOSeF5, and HOTeF5
The main sources of the -OSeF5 and -OTeF5 groups
are the acids. Certainly the most simple derivatives are
the most interesting ones, especially in terms of their
synthetic value. HOTeF5 is readily made by reaction
of telluric acid or its salts with fluorosulfuric a ~ i d , ~ ~ , ~ ~ , ~
Te(OH)6 + 5FS03H
3Se02Fz+ 4HF
ClOSF5 + HCl
20
-
-
5HzS04
+ HOTeF5
HZSeO4+ 2HOSeF5
-90 "C
Cla
+ HOSF5
whereas the selenium compound is made from SeOzF2.45,46
The sulfuric acid HOSFShas been prepared very
recently by the method of combining partially positive
chlorine (in ClOSFJ with partially negative chlorine (in
HC1).46r47 Because HOSF6 decomposes above -60 " C ,
the chemistry of the -OSF5 group is quite limited, in
contrast to that of -OSeF, and -OTeF5. All the materials are strong acids for which several physical
properties have been r n e a ~ u r e d . ~ Of
~ , special
~ , ~ ~ value
is 19F NMR spectroscopy, as the square-pyramidal
arrangement of the fluorine atoms gives rise to a typical
AB4 pattern. This is true for all derivatives as well
(Figure 7).49~50 The chemistry of HOSFE,(the only
six-coordinated sulfuric acid) is limited to its decomposition reaction.
HOSF5
-60 "C
HF
+ SOF4
The driving force of the reaction is the coordination
G. Mitra and G. H. Cady, J . Am. Chem. Soc., 81, 2646 (1959).
K. Seppelt, Angew. Chen. Int. Ed. Engl., 11, 630 (1972).
K. Seppelt, 2. Anorg. Allg. Chem., 428, 35 (1977).
K. Seppelt, Angew. Chem., Int. Ed. Engl., 15, 44 (1976).
W. Porcham and A. Engelbrecht, Monatsh. Chen., 102,333 (1971).
P. Bladon, D. H. Brown, K. D. Crosbie, and D. W. A. Sharp,
Spectrochim. Acta, Part A , 22, 2221 (1970).
(50) K. Seppelt, 2. Anorg. Allg. Chem., 399, 65 (1973).
(44)
(45)
(46)
(47)
(48)
(49)
Seppel t
216
Figure 8. Molecular structure of Se202F8,as determined by
electron diffraction. The structure of Te202FBis almost identical.
The Se-Se or Te-'re distances are surprisingly short. (Reprinted
with permission from ref 59. Copyright 1974, Verlag Chemie
GMBH.)
decrease on sulfur by formation of a sulfur-oxygen
double bond. This reaction is very similar to the decomposition of trifluoromethanol, which was made very
re~ently.~~,~~
CFBOCl
-120 o c
+ HC1-
-20 "C
HOCF,
C12 + HOCF3
CF20
+ HE'
In general, the SF5group closely resembles the CF3
group.
-
SF5NHz
-
CF3NH2
N=SF3
N=CF
+ 2HE3
+ HF52~54
HOSeF5, however, is stable to 290 "C and then decomposes in a quite different way:
HOSeF5
290 "C
HF
+ 1/202
+ SeF4
The oxidation state +6 for selenium is less favored than
in the two neighboring elements S and Te (see below).
This effect is certainly a result of the incomplete
shielding of the nuclear charge by the first filled d shell
(transition-metal contraction) and can be seen as related
to the late discovery of the p e r b r ~ m a t e and
s ~ ~A s C ~
about 150 years after corresponding compounds of their
neighbors in the periodic system. Thus HOSeF5 is a
highly oxidative, aggressive fluorinating agent and is
preferably handled in materials like Kel-F and Teflon
FEP. Only in the absolute absence of H F does it fail
to attack stainless steel. (HOTeF5 even exceeds the
stability of HOSeF5 and is by far a weaker oxidizer.)
The anhydrofluoride of HOSeF5, SeOF4, was very
difficult to synthesize. Absolutely pure NatOSeF5decomposes in vacuo a t 200 OC.j7
K. Seppelt, Angew. Chem., Int. Ed. Engl., 16, 322 (1977).
G. Kloter and K. Seppelt, J . Am. Chem. Soc., 101, 347 (1979).
A. F. Clifford and L. C. Duncan, Inorg. Chem., 5, 692 (1966).
G. Kloter, W. Lutz, K. Seppelt, and W. Sundermeyer, Angem
Chem., I n t . E d . Engl., 16, 707 (1977).
(51)
(52)
(53)
(54)
Accounts of Chemical Research
-
NatOSeF5NaF + SeOF4
SeOF4 probably has the same structure as SOF4trigonal bipyramidal with the double-bonded oxygen
in the plane. The unusual coordination number 5 can
be considered the driving force for its dimerization to
Se202F8,a molecule with a four-membered ring ~ n i t ~ ~ p
(Figure 8).
The analogous compound of tellurium, Te202F8,was
prepared by pyrolysis of Li'OTeFL or B(OTeF,), and
studied by electron diffraction as well.57-59The fourmembered-ring unit seems to be a general building
principle for the heavier nonmetal elements. Ions like
Te208(OH)26or 1208(OH2-,as well as As2OzF82-,all
exhibit a four-membered-ring unit.60-62The recently
determined crystal structure of I02F3shows it to be a
dimer, 1204F6,
with a four-membered ring.63 Such a
structure maintains the high coordination number 6 and
avoids double-bonded oxygen.
The monomer TeOF4 has not yet been detected.
Here the necessity of the coordination number 6 is
strict. Generally, double bonds on Te(V1) are not
known, as arithmetically they would lead to lower
coordination numbers. It is unnecessary to use orbital
theory to explain this effect. In IOF5 a double bond is
possible while the CN of 6 is retained. All other tellurium(V1) oxide fluorides have octahedrally coordinated telluriumz7 (Figure 6). (The more extensively
discussed problems of double bonds on silicon can be
explained similarly: any silicon double-bonded species
would have CN 3, and the latter seems to be impossible
in stable compounds under normal conditions. If the
favored CN 4 is retained, as in POFBor S02F2,double-bonded species are present.)
Conclusion
My intention has been to show that in the classical
field of the electronegative elements there has been
some exciting progress, despite the fact that this field
has largely been neglected in recent years in favor of
transition-metal chemistry. I regret that many other
significant discoveries in this field, such as the development of the halogen fluorides and oxyfluorides, could
~ , not
~ ~ be mentioned in the limited space of this Account.
I acknowledge t h e Deutsche Forschungsgemeinschaft and t h e
Fonds der Chemischen Industrie for grants i n aid of research.
T h e Deutsche Forschungsgemeinschaft is thanked f o r scholarships for several graduate students.
(55) E. H. Appelmann, J . Am. Chem. Soc., 90, 1900 (1968).
(56) AsC1, is made from AsC1, and Clz at -100 "C by UV irradiation.
It decomposes at -30 "C. The oxydechloride AsOC1, is more stable. K.
Seppelt, Angew. Chem., Int. Ed. Engl., 15, 3 i 7 (1976); 15, 766 (1976); 2.
Anorg. Allg. Chem., 434, 5 (1977); 439, 5 (1978).
(57) K. Seppelt, Angem. Chem., Int. Ed. Engl., 13,91 (1974); Z. Anorg.
Allg. Chem., 406, 287 (1974).
(58) H. Oberhammer and K. Seppelt, Inorg. Chem., in press.
(59) K. Seppelt, Angew. Chem., Int. Ed. Engl., 13, 92 (1974).
(60) 0. Lindquist, Acta Chem. Scand., 23, 3062 (1969).
(61) H. Siebert and H. Wegener, Angew. Chem. Int. Ed. Engl., 4,523
(1965).
(62) W. Haase, Chem. Ber., 106: 734 (1973); 107, 1009 (1974).
(63) L. E. Smart, J . Chem. SOC.,Chem. Commun., 519 (19771, and
literature therein.