HIGH-PRESSURE MOSSBAUER RESONANCE STUDIES OF THE CONVERSION OF Fe(III) TO Fe(II) IN FERRIC HALIDES
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HIGH-PRESSURE MOSSBAUER RESONANCE STUDIES OF
THE CONVERSION OF Fe(III) TO Fe(II) IN FERRIC HALIDES*
BY G. K. LEWIS, JR., AND H. G. DRICKAMER
DEPARTMENT OF CHEMISTRY AND CHEMICAL ENGINEERING
AND MATERIALS RESEARCH LABORATORY, UNIVERSITY OF ILLINOIS, URBANA
Communicated July 24, 1968
In several recent papers1-3 it has been shown by high-pressure Mbssbauer
studies that ferric ion reduces to ferrous ion at high pressure, and that this is an
essentially reversible process. In this paper we present data on FeCl3, FeBr3,
and KFeCl4 to 175 kilobars (kb). FeCl3 and FeBr3 each have an octahedral
arrangement of the Cl- ions around the iron, while the FeCl4- ion exhibits
tetrahedral symmetry.
The compounds were synthesized from iron enriched to 77 per cent in Fe57
by using methods available in the literature.4-8 Since FeCl3 and FeBr3 are
extremely hydroscopic and KFeCl4 is moderately so, all handling was done in a
dry box in a dry argon atmosphere. Each material gave the appropriate X-ray
spectrum upon analysis. The high-pressure M6ssbauer techniques have been
previously described.9 At each pressure the spectrum was run until at least
200,000 counts had accumulated in each channel. (Where the effect was over
20 per cent only 150,000 counts per channel were obtained.) The spectra were
computer-fit with pairs of Lorentzian peaks. A minimum of five loads was run
on each substance at 2940K. In addition, at several pressures isobars were
run to 4180K, and several high-temperature isotherms were obtained.
As discussed in detail in reference 1, it is easy to identify high-spin ferrous
and ferric states from M6ssbauer spectra. The ferrous state exhibits a rela-
tively low electron density with isomer shifts in the range 1.2-1.4 mm/sec below
bec iron and high quadrupole splitting (2.0-3.0 mm/sec). The ferric state ex-
hibits isomer shifts in the range 0.3-0.5 mm/sec below iron and small quadrupole
splitting (0-0.7 mm/see) caused entirely by a lack of cubic symmetry at the
Fe(III) center.
It is difficult to determine highly precise values for the change of isomer
shift and quadrupole splitting with pressure on systems which exhibit a large
change in the ferrous-ferric ratio, as these do (see below). However, Table 1
lists the average values for these systems. Both ferric and ferrous ions show an
increase in electron density at the nucleus with increasing pressure, as do most
nonreacting systems.1 The major cause is probably changes in the 3d-3s
shielding as discussed in reference 1. There is a relatively large increase in
quadrupole splitting for the ferric systems, whereas the ferrous quadrupole
splitting is relatively independent of pressure. This is not unreasonable in
first order as the splitting for Fe(III) results directly from lattice distortions
(qiat), while the Fe(II) splitting is due to asymmetric distribution of the 3d
electron (qvai).
The most striking feature of the spectra is the large conversion of the ferric ion
to the ferrous state even at pressures as low as 10 kb. Figure la-d shows typical
spectra. Figure 2 exhibits a plot of ln K versus ln P, where K = CII/CIII is
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414VOL. 61, 1968 CHEMISTRY: LEWIS AND DRICKAMER 415
z
o .000
a-
cr .030
0
aI .060
(a) 3-5 kbar
< .090
z
0
p .120
c: .150
LL
FIG. la, b.-Typical spectra,
FeC1h. (a) 3-5 kb; (b) 11
kb. .000
z
°.050
0-
° .100
(a
Ji .150
z
°.200
( .250
-2.0 -1.0 0 1.0 2.0 3.0 4.0
DOPPLER VELOCITY IN MM/SEC
FIG. ic, d.-Typical spectra,
FeCh3. (c) 127 kb; (d) after
release of pressure.
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TABLE 1. Isomer shifts* and quadrupole splittingt of the ferric halides.
Pressure
(kb) FeCl3 FeBr3 KFeC14
Fe(III) isomer shift 4 0.35 0.35 0.32
40 0.32 - -
100 0.27 0.32 0.31
125 0.26 - -
Fe(II) isomer shift 10 1.30 1.30 1.28
40 1.27 -
100 1.24 1.22 1.26
125 1.22 - -
Fe(III) quadrupole splitting 4 0.88 0.80 0.50
40 1.20 0.95 0.80
100 1.40 1.05 1.10
125 1.49 1.20
Fe(II) quadrupole splitting 10 2.15 2.30 1.95
100 2.10 2.17 1.95
* Mm/sec relative to bcc iron metal.
t Mm/sec.
an equilibrium constant. It is seen that these data follow quite accurately the
form:
K = APB, (1)
where A and B are constants. An analysis of the data of Champion'0 and of
Vaughan" shows equally good agreement with this form of equation. The
constants are included in Table 2. It may be mentioned parenthetically that
the data for a variety of other ferric systems now under study in this laboratory
appear to follow this same relation.
From Table 2 it may be seen that B 0.5 for more ionic compounds and is
-
measurably larger for more covalent materials. It was difficult to obtain good
high-temperature data for the halide systems because of the high conversion,
but, within our accuracy, B appears to be independent of temperature for
these materials. This is not universally true. Figure 3 shows the results of
isobars obtained at 11 kb. The heats of reaction are listed in Table 3. They
increase measurably with increasing temperature for FeCl3 and FeBr3. Nor-
mally samples were diluted with boron, but no significant difference in behavior
was observed for samples diluted with aluminum oxide or with boron carbide.
It should be pointed out that what is observed is an equilibrium phenomenon,
not the result of slow kinetics. Typical runs at a single pressure took 8-48
hours, and consecutive runs at the same pressure showed no change in con-
version. However, when the pressure was increased, a definite increase in con-
version was noted in the spectrum displayed on the oscilloscope as soon as
enough counts were available to define a spectrum-usually within five minutes
or so.
From a molecular standpoint, there are two questions to be answered:
(1) How are the electronic levels of Fe(III) lowered with pressure vis-A-vis
the ligand levels to give thermal electron transfer?
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6.0
4.0 -
FeC0 3
2.0-
.0
0.6-
0.2
0.1.
5 10 20 30
1
60 100 200
P(kbor)
FIG. 2.-Ln equilibrium constant versus In pres-
sure, FeC13, FeBr3, and KFeC14.
(2) Once this lowering has taken place, why do we not observe complete re-
duction of the ferric ion within a short pressure range?
It is straightforward to identify the most probable energy levels involved.
For an octahedral complex (see Ballhausen and Gray,'2 p. 103) the transfer
must be from the ligand nonbonding t2. level to the predominantly metal anti-
bonding t2,. For the tetrahedral case, the corresponding levels are the tj(7r)
and e(7r). Broad charge transfer peaks are observed in most such complexes at
3-6 ev, with tails extending sometimes through the visible and even into the infra-
red region of the spectrum.'3
One can imagine at least two mechanisms which could contribute to a red
shift of these peaks (i.e., a lowering of the t2. t20 energy difference). The bond-
ing for ionic complexes is largely a bonding, but it is possible for the t2, orbital to
bond with the lr* orbitals of the ligand.'4 For these complexes one expects a
markedly smaller r-ir overlap than u-a- overlap at 1 atmosphere, and with
increasing pressure one would anticipate a larger increase in the former than in
the latter. This overlap should tend to stabilize the t2g orbitals as they are
bonding with respect to the ligand -r* orbitals. The reduction of Fe(III) ion
takes place with a wide variety of ligands in various symmetries, so the possible
causes, including ir bonding, must be independent of the details of structure.
As a prototype, the Fe-Cl system was analyzed in both octahedral and tetra-
hedral symmetry. Group overlap calculations were made as a function of
interatomic distance in accordance with the procedures outlined in reference 12
and in references contained therein. Clementi's1' wave functions were used.
Figure 4 shows the change in 7r and a- overlap with Fe-Cl distance for FeI+-Cl-
in octahedral symmetry. Calculations for tetrahedral symmetry and for Fe'+-
Cl- did not differ qualitatively. Indeed, the wr overlap increases faster than the
a, which should be a stabilizing factor. However, the difference is not great
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TABLE 2. Parameters A and B for K = AP*.
Compound A B
FeCli 0.265 0.564
FeBr3 0.076 0.426
KFeCh4 0.091 0.497
Ferric phosphate1 0.787 0.457
Ferric citrate'0 0.112 0.350
Phosphate glass* 0.486 0.312
KFe(CN)6'0 0.109 2.06
Ferric acetate (3820K)* 0.22 X 10-6 3.05
Ferric acetate (4180K)* 0.022 0.986
* These data will be discussed elsewhere.
enough to be a major factor in any large red shift of the charge transfer peak.
The-second factor that would tend to decrease the tku -a tag energy difference
is the radial spreading of the 3d orbitals with pressure. It is well established
from high-pressure optical studies16' 17 that the interelectronic repulsion (Racah
or Condon-Shortley parameters) of transition metal ions decreases with in-
creasing pressure-by as much as 8-10 per cent in 150 kb. This can best be
explained as involving an increase in the radial extent of the 3d orbitals by attrac-
tion of the ligand nuclei. Using Watson's'8 calculations of the Condon-Shortley
parameters and of the energies associated with various electron configurations
and ionic radii, Vaughan'9 has shown that a red shift of several electron volts
could easily be consistent with the above decrease in interelectronic repulsion.
The effect of pressure on charge transfer peaks has been measured,20 and red
shifts in the order of 0.5-1.0 ev in 100 kb have been observed. Also, the absorp-
2.5 3.0 3.5
I/T x 1030K'
FIG. 3.-Ln equilibrium constant versus 1/T
at 11 kb, FeC13, FeBr,, and KFeCl4.
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TABLE 3. AH of reaction.
AH (ev)
500C 1200C
FeCi3 0.12 0.18
FeBr3 0.20 0.32
KFeC14 0.07 0.07
tion edge of TICI, which involves a not unrelated transition, shifts red by about
1.0 ev in 100 kb.2' This, of course, is not enough to move the optically observed
energy differences to zero.
We must keep in mind that optical transitions are subject to the Franck-
Condon principle and take place vertically on a configuration coordinate diagram
(Fig. 5). The transitions involved in the reduction of Fe(JII) observed here are
thermal transitions and occur sufficiently slowly that the coordinates can assume
their new equilibrium positions. Figure 5 shows that the thermal energy can be
of the order kT while there is still substantial optical gap.
The temperature coefficient can be explained also from this diagram. There
is a Boltzmann factor for electron transfer from the ground state to the excited
electronic state. In addition, with increasing temperature there will be increased
occupation of excited vibrational levels of the ground electronic state, which will
further increase In K at higher temperatures.
The probable explanation of the pressure dependence is more subtle. When
an Fe(IJI) ion reduces to Fe(JI), the radius of the metal ion increases, and one
forms a radical from one of the ligands. (In particular cases one may get
combinations of ligands into ion-radicals such as C12- or even molecules like C12.
It is also possible for the hole to be smeared out over all four or six ligands). In
1.8
G/Go 1.4 0ma(
1.0 .98 .96 .94 .92 .90 .88 .86 .84
R/R0
FIG. 4.-Relative sigma and pi overlap versus relative FeCl
distance for Fe3+ - C1-. Solid line, CL2 = 0.1 (a is covalency
parameter); dashed line, a2 = 0.3.
One-atmosphere group overlap integrals (Go):
a2 = 0.1 - 7-r = 0.0835
a-oa = 0. 3043
a 2 = 0.3 - -r = 0.0836
a-a = 0.3169
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any case, there will be locally a decrease in volume, a
displacement of charge, and a lattice strain or distortion.
These will set up a stress field which will interact with
neighboring complexes. This can distort the shapes
o 0 of the potential wells and change their positions enough
TS X to inhibit electron transfer on these complexes. In-
creased pressure further lowers the excited state vis-
FIG. 5.-Schematic con- A-vis the ground state and increases conversion. This
figurational coordinate di- urn creates
agram. 0, optical trans- in t
i
more stress fields which give further dis-
ition; T, thermal transi- tortion. The process is thus a continuing one and is
tion. consistent with the form of pressure dependence ob-
served experimentally.
The argument for local strain around each converted site is consistent with
the observation of considerable hysteresis for some systems upon release of
pressure. In some cases the spectrum returns completely to Fe(III) directly
upon release of pressure. In others the reverse transformation runs only
partially. When the pellet is powdered, one obtains substantially complete
reconversion.
One can analyze the results thermodynamically:
K = exp (-RT4 ' (2)
la In K PAV P(VIII - VII)
kblnP IT RT RT ' (3)
where VI,, and VI, refer to the ferric and ferrous ions with their associated
ligands.
The concentrations are taken as proportional to the areas under the Lorentzian
peaks. They are nominal since they do not take into account self-absorption,
or differences in (x2) at different sites. Nevertheless, changes in relative area
should give reasonable measures of concentration changes.
Equation (3) can be rearranged:
a In Cr_ P (V"1 - V") C (4)
bin P RT C1.(4
Thus, the fractional increase in conversion with fractional increase in pressure
is proportional to the concentration of sites available for conversion. The
proportionality coefficient is the work to create a ferrous site, measured in
thermal units. The fact that this is constant is a reasonable first approximation.
Apparently higher terms are negligible.
The authors wish to acknowledge very helpful discussions and correspondence with
R. W. Vaughan and C. P. Slichter, whose analyses have contributed a great deal to our
understanding of the problem. G. K. L. would like to acknowledge financial assistance
from a Chrysler Corporation fellowship.
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* This work was supported in part by the U.S. Atomic Energy Commission under contract
AT(11-1)-1198.
1 Champion, A. R., R. W. Vaughan, and H. G. Drickamer, J. Chem. Phys., 47, 2583 (1967).
2 Champion, A. R., and H. G. Drickamer, J. Chem. Phys., 47, 2591 (1967).
3 Champion, A. R., and H. G. Drickamer, these PROCEEDINGS, 58, 876 (1967).
4Maier, G. G., U. S. Bur. Mines, Tech. Papers, No. 360 (1925), p. 40.
5 Schafer, H., Angew. Chem., 64, 111 (1952).
6 Gregory, N. W., and B. A. Thackery, J. Am. Chem. Soc., 72, 3176 (1950).
7Gregory, N. W., J. Am. Chem. Soc., 73, 472 (1951).
8 Friedman, 11. L., and H. Taube, J. Am. Chem. Soc., 72, 2236 (1950).
9 Debrunner, P., R. W. Vaughan, A. R. Champion, J. Cohen, J. A. Moyzis, and H. G.
Drickamer, Rev. Sci. Instr., 37, 1310 (1966).
10Champion, A. R., Ph.D. thesis, University of Illinois (1967).
11 Vaughan, R. W., Ph.D. thesis, University of Illinois (1967).
12Ballhausen, C. J., and H. B. Gray, Molecular Orbital Theory (New York: W. A. Benjamin,
Inc., 1964).
13 Orgel, L. E., Quart. Rev., 8, 422 (1954).
14Gray, H. B., and Beach, J. Am. Chem. Soc., 85, 2922 (1963).
15 Clementi, E., "Tables of Atomic Functions," a supplement to IBM J. Res. Develop., 9, 2
(1965).
16Stephens, D. R., and H. G. Drickamer, J. Chem. Phys., 34, 937 (1961); 35, 424, 427, 429
(1961).
17 Zahner, J. C., and H. G. Drickamer, J. Chem. Phys., 35, 1483 (1961).
18 Watson, R. E., Technical Report No. 12, Solid State and Molecular Theory Group,
Massachusetts Institute of Technology (1959).
1 Vaughan, R. W., private communication (many of the ideas concerning the theory of the
relative shift of the t2h and t2,,tlevels are due to Dr. Vaughan).
20 Parsons, R. W., and H. G. Drickamer, J. Chem. Phys., 29, 930 (1958).
21 Zahner. J. C., and H. G. I)rickamer, J. Phys. Chem. Solids, 11, 92 (1959).
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