A KINETIC STUDY OF THE ELECTRON '

EXCHANGE BETWEEN THE

12-TUNGSTOCOBALTA'I‘EUI) AND THE '

12 - TUNGS'FOCOBAL‘FATH II I)

ANIONSIN AQUEOUS SOLUTION :37 I

I I Thesis I00 II" Danna GIT; pI'I D
MICHIGAN STATE UNIVERSITY

Paul G Rasmussen
1964

THESIS

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Michigan Scare

University

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huIIer .II :IAIE. LII-L'a/iR «VIII-Y
DEPARTI‘.‘ IliNT 02': LL' IVIISTRYL—
EAST LANSING, MICHIGAN

ABSTRACT
A KINETIC STUDY OF THE ELECTRON EXCHANGE BETWEEN THE
lZ-TUNGSTOCOBALTATE(II) AND THE lZ-TUNGSTOCOBALTATE(III)
ANIONS IN AQUEOUS SOLUTION

by Paul G. Rasmussen

The kinetics of the electron exchange reaction between the l2-tungsto-

cobaltate(II) and the lZ-tungstocobaltate(III) anions was studied at 00 in
aqueous solution. The ions were separated by the selective precipitation
of the lZ—tungstocobaltate(III) ion with (Bu)4NI at pH = S. The reaction
was found to be first order with respect to each ion, and second order
rate constants of 10-2 to 101 Mflsec.‘1 were found depending on experi-
mental conditions. The rate of exchange was found to be a function of
the cation present, in particular, the rate in the presence of potassium
ion.was much greater than that for lithium. The rate was also found to
depend on the ionic strength (adjusted with LiCl), and these results were
interpreted with the theoretical equations of R. A. Marcus. At constant
ionic strength however, the rate was found not to depend on the hydrogen
ion concentration for the acid-salt pair of HCl—LiCl. The temperature
dependence of the rate was studied and the parameters of activation were
obtained. The reaction was also studied in dioxane-water mixtures and

the effect of dielectric constant on the rate was determined. In the
light of the data obtained, an outer—sphere mechanism is postulated for
the system and the results are compared to the theoretical predictions
of R. A. Marcus.

A preliminary study was made of the electron paramagnetic resonance
spectra of the l2-tungstocobaltate(II) and lZ—tungstocobaltate(III) ions,
diluted in a host-lattice of potassium lZ-tungstosilicate. To obtain a
signal, it was necessary to go to very low temperatures (o/‘l°K). No

hyperfine splittings were observed in the spectrum of either compound.

A KINETIC STUDY OF THE ELECTRON EXCHANGE BETWEEN THE
lZ-TUNGSTOCOBALTATE(II) AND THE 12-TUNGSTOCOBALTATE(III)

ANIONS IN AQUEOUS SOLUTION

By

Paul G. Rasmussen

A THESIS

Submitted to
Michigan State University
in partial fulfillment of the requirements
for the degree of

DOCTOR OF PHILOSOPHY

Department of Chemistry

l96h

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.‘v. \ J.
x I I I

RV‘gg-B

ACKNOWLEDGMENT

Th author is pleased to gratefully acknowledge many helpful sug—
gestions and encouragement of Professor Carl H. Brubaker, under whose
guidance this research was conducted; the patience and sustaining con-
fidence of Mrs. P. G. Rasmussen, wife of the author; and the financial

aid from the Atomic Energy Commission.

ii

TABLE OF CONTENTS

I. INTRODUCTION . . r . . . . . . . . . . . . . . . . .
II. HISTORICAL . . . . . .
A. Heteropolyacid Chemistry . . . . . . . .
B. Electron Exchange Studies . . . . . . .
III. THEORETICAL . . . . . . . . . . . . . . . . .
A. The Rate Law of the Exchange Process . . .
B. Theoretical Prediction of the Rate Constant
IV. EXPERIMENTAL .
A. Preparation of Reagents . . . . . . .
B. Analytical Methods . . . . . . . . . . . . .
C. Procedure . . . .
D. Errors . . . . .
V. RESULTS . . . . . . . . .
A. Determination of Kinetic Order .
B. Dependence of the Rate on Hydrogen Ion Concen-
tration . . . . .
C. Dependence of the Rate on Ionic Strength .
D. Dependence of the Rate on Temperature
E. Specific Cation Effects . . . . .
F. Dependence of the Rate on Dielectric Constant
VI. DISCUSSION .
VII. LITERATURE CITED .
APPENDIX I. Original Kinetic Data

APPENDIX II.

APPENDIX III.

Visible Region Absorption Spectra of the 12—
tungstocobaltate(II) and the lZ—tungstocobaltate
(III) ions in Aqueous Solution . . . . . . .

A Preliminary study of the electron paramag-
netic resonance spectra of the l2—tungsto-
cobaltate(II) and lZ-tungstocobaltate(III) ions

Page

\OU'LU'L

12

12
1b

18
18

20
21

23
25
25
28
29
33
38
AZ
LB

51

63

65

LIST OF TABLES

TABLE Page

I. Dependence of the rate on reactant concentrations . . . . 26
II. Dependence of the rate on HCl concentration . . . . . . . 28
III. Dependence of the rate on ionic strength . . . . . . . . . 30

IV. Dependence of the rate on temperature and ionic strength.
The parameters of activation. . . . . . . . . . . . . 3b

V. Dependence of the rate on various cations . . . . . . . . 36

VI. Dependence of the rate on dielectric constant . . . . . . NO

iv

LIST OF FIGURES

Figure Page
1. The "Keggin" structure for heteropoly ions . . . . . . . . h
2. Exchange curve for some typical rate data . . . . . . . . 27

. Graph of log kr X§.\fir . . . . . . . . . . . . . . . . . . 31

. Graph of log kr vs. e"Kr . . . . . . . . . . . . . . . . 32

3
u _
5. Graph of log kr 103/T vs. 103/T . . . . . . . . . . . . . 35
6

Graph of rate constant XE' [MCl] . . . . . . . . . . . . . 37

7. Weight percent dioxane vs. dielectric

constant at O0 . . . . . . . . . . . . . . . . . . . . 39

8. Graph of log kr vs. i/DS . . . . . . . . . . . . . . . . . L1

I. INTRODUCTION

Kinetic and mechanistic work in inorganic chemistry has been
greatly stimulated in recent years by the availability of radioisotOpes(l).
The very high efficiencies with which radioactivity may be detected, al—
lows the use of tracer methods, and such methods have been extensively
applied to self-diffusion studies and isotopic exchange reactions as
well as to the more conventional kinetic systems. Some of the isotope
exchange reactions that have been studied have been of the type where
simple electron exchange is a possible mechanism. An example is the

manganate—permanganate system:

 

MnO4_ + MnO4= < MnO4- + Hno4=

studied by Sheppard and Wahl (2). There is good evidence in this case
that there is no interpenetration of the coordination spheres (27). An
example of a similar cationic system is that of the iron l—lO phenanthro—

line complexes (N7):

Ee(phen)3++ + Ee(phen)3+++ <+——j> Fe(phen)3++ + Ee(phen)3+++

which has also been studied by Wahl and his students. Electron exchange
in these systems is considered as a barrier penetration phenomena and
the activated complex is designated as an outer-sphere complex. Un-
fortunately, relatively few of the systems which have been studied can
be unambiguously classified as being of this simple mechanistic type,
and additional data regarding the effect of various ligands and ion

'atmospheres is desirable.

2

This investigation is concerned with what is believed to be a new
example of an electron exchange reaction which proceeds.by an "outer—
sphere" mechanism. The reactants involved are the lZ—tungstocobaltate(II)
and the l2-tungstocobaltate(III) anions. These heteropoly acid anions
are known to have a "Keggin" (3) type of structure in which the central
cobalt atom is surrounded by tungsten-oxygen octahedra sharing corners
and edges. (See Fig. 1) Such species are substitutionally inert, and
have large formation constants in acid solution (10).

In addition to kinetic considerations, a study of this system seemed
desirable because there were no previous reports of electron exchange
work on heterpoly complexes in the literature, and because the ability
of the central atom electrons to penetrate the tungsten—oxygen "cage"
was unknown. Certainly the heteropoly complex cannot aid electron ex—
change by "conduction" through conjugated ligands like the cyanide or
phenanthroline complexes can. If there is not strong overlap between
the reactants in the activated complex, then a recent theory due to R.
A. Marcus (3,h) should be applicable for predicting the rate of electron
exchange, and a test of the theory is possible.

Experimentally, the study was carried out by using 60Co and the
usual tracer methods. A major experimental problem in an investigation
of this sort is the determining of a rapid, effective means of separat-
ing two ions that differ only by one unit of charge. Various methods
have been cited in the literature on electron exchange, including sol—
vent extraction, ion exchange chromatography, precipitation, and diffu-
sion. In this system, tetra-n-butyl ammonium ion was found to be an

effective reagent for the selective precipitation of the cobalt(III)

3

anion in the presence of the cobalt(II) anion and an acetate buffer.
Although this method leads to various amounts of separation induced ex—
change, it was found to be reproducible for a given set of conditions,
so the induced exchange would not affect the results.(23) The counting
of the precipitated samples was done with a windowless gas flow counter.
A McKay plot (23) of the data so obtained yields a straight line with

slope proportional to the rate of electron exchange.

A

Figure l. The ”Keggin” structure for heterpoly ions
(All figures to the same scale.)

    
  
  
   
    

 

/ «x
/

  

.3
b
A“ “‘1’
karat

A. Spatial diagram of atoms; (3 Oxygern o Tungsten, C Cobalt.

B. Cubo-octahedron formed by tungsten atoms at the vertices.

 

C. Polyhedral diagram, note central tetrahedron.

II. HISTORICAL
A. Heteropolyacid Chemistry

The discovery of the large group of compounds designated hetero-
polyacids was due to Marignac (6),who found in 1861.that tungstic and
silicic acids reacted in solution to form stable compounds which could
readily be crystallized out. The heteropoly acids and their salts are
a unique group of compounds in which vanadium, molybdenum, tungsten or
other heavy addenda atoms combine with oxygen and various hetero atoms
to form complex, high molecular weight anions. Heteropoly compounds
are commonly classified by the ratio of addenda atoms to hetero atoms.
Although all the integral Species have been reported from 12:1 to 1:1,
Species with ratios of 12:1, 9:1 and 6:1 are by far the most numerous.

The determination of the structures of these compounds has evolved
slowly over the years and is still the subject of current research.
Around the turn of the century, Rosenheim, and later Miolati, attempted
to write systematic formulae for the heterpoly ions by considering them

as derivatives of a parent acid as follows:

 

+n — +
H12_n[X 06] + 611207 2 > le_n[x n(wzo7)61

This system was appealing because it applied the Werner coordination
theory to what was considered to be an ion analogous to the dichromate
ion. Although it leads to systematic formulae, this method does not

predict the correct basicities for the anions and real knowledge of the

1
Spatial structure was still lacking. In the early thirties, Pauling

applied his empirical bonding rules to the problem and proposed a

S

6
structural model for the 12-tungstophosphate(V) ion. Although the gross
features of Pauling's structure were correct, the X—ray diffraction work
of Keggin(3) showed in detail the oxygen sharing properties of the octa-
hedra involved. Recently, his model has been verified and exact struc-
tural parameters found for the compounds used in.the present study by
Eriks and Yannoni (37). The "Keggin" structure (See Fig. l) locates
the tetrahedrally coordinated hetero atom in the center of a cubo—
octahedron with tungsten atoms at the vertices. Each tungsten atom is
in a slightly distorted octahedron of oxygen atoms. Four clusters of
0N3013) tetrahedrally surround the central atom. Each cluster is made
up of three octahedra sharing edges in such a way that an oxygen of the
central metal atom is common to all of them. The four clusters are
bound together by sharing corner oxygens. The over—all formula is then
[XnO4leO36]8_n for a heteropoly anion with the"Keggin"structure.
Some general properties common to this class of compounds are: (7)
1. Very high molecular weights — to over AOOO.
2. Extraordinary solubility in water.

3. Precipitated by cations such as tetra—alkyl ammonium,
rubidium, cesium, and guanidinium ions.

h. Colors which range through the spectrum as well as color—
less compounds.

5. Strong oxidizing and reducing properties in many anions.
6. High degree of hydration in the crystalline state.
7. Degradable by alkali.
To the general class of heteropoly anions, belong several members
first prepared and characterized by Baker and McCutcheon in 1956 (8).
These authors reported the following preparative scheme for the 12—tungsto-

cobaltates.

 

 

Co+2
hot neutral
\ungstate sol'n.
+
[(30+2C3‘3‘I2I"I1204.2I“8 < H > ICO+2IW207)5]-1°
(blue-green) (NH4)2C03 (blueegreen)
with decomp.
ox.‘ rred. ox.J rred.
+
(dark brown) (yellow)

They note that the relatively late discovery of these compounds is prob-
ably due to the unusual circumstance of their formation in neutral solu-
tion, whereas most heteropoly tungstates form in acid. In later papers
(9,10), Baker and Simmons report further on the structures of the mono-
cobalt compounds. The main features of the "Keggin" structure are pre-
served; the central cobalt atom is tetrahedrally coordinated, and the

formulae are revised to the following:

[C0+ZO4W12036]-6 <§Z§§iZZE [COIL‘IOAH’I1.2036I-5
I red' 11

They found by careful dehydration experiments that all the water of
crystallization could be removed without destroying the heteropoly ion.
This would seem to indicate that there is little vacantspace within the
ion, and indeed, a study of the ionic radii shows that the oxygen atoms
in the heteropoly structure are very nearly close—packed. This feature
as well as other structural aspects is discussed elsewhere (2h).

The oxidation and reduction were found to take place reversibly in
1N H2804 at a potential of -l.O7 volts indicating small stabilization

for the tripositive oxidation state. The bulk magnetic susceptibility

was also determined (25) and the applicability of the Curie—Weiss law

8

rigorously verified. The reported values for magnetic moment were b.25
and 5.07 Bohr Magnetons for I and II respectively. Finally, they found,
in work pertinent to this investigation, that there was no evidence for
dissociation of either monocobalt species as determined by cryoscopy in
fused NaZSO4°1OHZO (33). It is the electron exchange between the mono-
cobalt redox pair (hereafter referred to as I and II respectively) that
is the subject of this thesis. Recently, several review articles have

appeared that may serve to apprise the reader of modern developments in

the area of heteropoly chemistry (ll, 12, 13).

9

B. Electron Exchange Studies

Since the late nineteen forties, electron exchange kinetics has been
a remarkably active area of research. The early work is well summarized
up to 1950 by A. C. Wahl and N. A. Bonner (1h). At that time there were
insufficient data available, however, to allow a quantitative explana—
tion of the factors controlling the rates of electron exchange. Exper—
imental work since 1950 has shown that no single mechanism is generally
Operative in inorganic redox reactions and that ligands and media play
important roles. The system most extensively studied has been Fe(II)—
Fe(III). The effects of acid, fluoride ion, sulfate ion, phosphate ion,
organic acid anions, temperature, heavy water, alcohol, and other agents
on the rate have all been examined. (38,39,hO,h1,b2) From these studies
and those on a number of other systems including many net reaction sys-
tems, it appears that the possible mechanisms for redox reactions fall
into two broad categories. Under the title of "strong over-lap mechan—
isms” are the atom transfer and some ligand bridge mechanisms. Although
these two possibilities have been distinguished in favorable cases,
there is often an ambiguity. Atom transfer can be ruled out in the so-
called ”conduction" mechanisms in which ligands such as para-phthalic
acid greatly accelerate the rate, due to the ease with which the electron
can traverse the conjugated set of n orbitals. The strong over-lap
mechanisms, in either case, are viewed as directly involving the primary
coordination spheres of the reactants. In weak—overlap mechanisms, this
is not true, and the reactants are inert to ligand substitution. In these
cases, the electron transfer is considered as a quantum mechanical bar-

rier penetration. The activated complex is designated as an outer—sphere

lO

complex although its exact nature is difficult to assess. Relatively
few of the electron exchange systems studied have been assigned weak-
overlap mechanisms and most of these are very fast. Examples of weak-
overlap systems include MnO4—-MnO4=, Fe(phen)3++-Fe(phen)3+++, Os(dipy)3++-
Os(dipy)3+++, Fe(CN)6_3-Fe(CN)5-4, and IrCiG‘Z-Irc16‘3.(1)

Theoretical approaches to the problem are varied and contradictory.
Libby (15) has stressed the importance of the Franck-Condon principle
and the ligand rearrangements necessary in electron exchange. R. J.
Marcus, Zwolinski and Eyring (16) have treated the barrier penetration
aspects of the problem. Their results require calibration,however, with
an experimental system. R. A. Marcus (A, 5) has dealt with the weak-
overlap case and has derived rate expressions with no adjustable para-
meters. He has included in his derivation, terms that account for the
rearrangement energy on going from reactant to product in cases where
this is appropriate. Hush (25) derives in his so-called adiabatic theory,
results that are in substantial agreement with those of R. A. Marcus by
using an electron reaction coordinate parameter. Laidler (26) proposes
a "non-adiabatic" theory which is similar in approach to that of Marcus,
Zwolinski and Eyring. Agreement between experimental data and any of
the theories is only fair, and a thorough critique of the various approaches
as well as additional experimental data would be desirable. The work of
R. A. Marcus has gained the widest acceptance at present and his model
will be applied to the system involved in this investigation in a follow-
ing section. Electron exchange between heteropoly anions should be an
excellent example of the weak-overlap type of mechanism because of the

presumed non-lability of the discrete species in solution and their large

11
size. Unfortunately, however, a detailed knowledge of the transition
state is not obtainable in general from rate studies alone.‘ Several

review articles have summarized the available experimental work in

recent years (17,18,19,20,21).

III. THEORETICAL
A. The Rate Law For the Exchange Process

For a stable homogeneous phase, in which all the atoms of a partic-
ular oxidation state are chemically equivalent, and the half—life of the
tracer is long compared to the exchange time, the McKay equation (23)
gives the rate of growth of radioactivity in an initially unlabeled species
as:

(1 - x/Xm) = epr _ a ;bb Rt] (1)

 

t = time
a b = total concentration of each reactant
X = activity in initially inactive species at time t

X0, = activity in initially inactive species after infinite
time

R = rate of exchange process.
This law is followed, regardless of the mechanism of the reaction between

the species involved. The equation may be rewritten as:

2.303 mm - x/xm) = - a gbb Rt (2)

 

From which it is apparent, that a plot of log (1 - X/Xa,) versus time
should be linear. The half-time, tl/Z is easily obtained from a graph

and R is calculated:

aab b (0.693/t1/2) (3)

To determine the reaction order with respect to a reactant, the

rate R is equated to a general rate law:

12

13

R = kria]“[b]B[c1Y (u)
then:

lo R
= (5)
31°9 8 [b],[c] ‘1

therefore a plot of log R versus log [a] for fixed [b] and [c] will have
a slope of a, the order with respect species a. The rest of the ex-
ponents may be similarly determined. A simpler.but less rigorous pro-
cedure, is to assume a value for a, B, Y etc. and then calculate kr,

the rate constant,using the rates obtained for various values of a, b,

c etc. If kr remains constant within experimental error, the assumed
order is adopted. While this method works well for simple systems,
clearly it must be applied with some caution.

Frequently, the reactants in an electron exchange reaction system
can only be incompletely separated, or a rapid exchange may occur dur-
ing the separation. These effects result in an apparent zero time ex-
change. Zero time exchange can occur to the extent of 100% in which
case one does not know whether the reaction itself is rapid or whether
the separation method is at fault. If the apparent zero time exchange
is less than 100%, however, Prestwood and Wahl (22) have shown that
the slope of log (1 - X/Xa,) versus time is unaffected and the rate may
still be obtained. This result depends of course on the apparent zero

time exchange remaining constant for a given set of conditions.

1h

B. Theoretical Prediction of the Rate Constant

0f the current theories mentioned in.the introduction of this thesis,
the author feels that the one due to R. A. Marcus fits the experimental
data the most satisfactorily and is the most complete. Therefore it will
be used as a model for comparison with the present experimental system
of heteropoly anions. The Marcus theory predicts the rate of electron
exchange for the weak-overlap case 1.2. when orbital overlap between
the reactants is small in the activated complex. Experimentally this
situation obtains when the first coordination sphere of the reactantS»
is non-labile and when the ions are large and weakly.solvated. These
conditions appear applicable to l2—tungstocobaltate ions.

The following equations summarize the results of the Marcus theory.
For a detailed development, the reader is referred to the original papers
(h,5). The factor exp[-H.r] will not be found in these papers, however.
Its inclusion was suggested to the author (A9) as a correction for the
effects due to the large sizes of the ions. For heterpoly ions it is
a poor approximation to compute the energy of interaction on the basis
of point charges even at quite low ionic strengths. In the limit of
zero ionic strength, the exp[—FKr] term approaches unity and the original

equation of Marcus is obtained.

 

kr = Z exp[ - AF*/RT] (6)
y 8192
AF" = D r epr-Kr‘] + mzx (7)
s
where: kr = rate constant
Z = collision number in solution equals about 1011
T = absolute temperature (Kelvin)
D = dielectric constant of solvent

15

AF = free energy of activation (Gibbs)

e1,e2 = charges on the reactants
r = distance of closest approach of the reactants
m = derived constant equal to -0.5 for electron

exchange reactions

H, = "Debye-Hfickel kappa" equal to (5.03 x 109)\lu;DST

u = ionic strength.
I = A0 + 1, (8)
_ 1 1 _ l l l 2
K' K'p
h = 2 . _'.L+‘L.'p (“h-”2 (10)
J KJ K J
where: X0 = ”lambda outer"; contribution to X of surroundings

of activated complex

I. = "lambda inner"; contribution to X of changes in
first coordination sphere.

r = a1 + a2
n = refractive index of the solvent

Ae = amount of charge transferred as a result of the
reaction

K K = force constants of jth vibrational coordinate for a
j? j species as reactant and product respectively.

Aq.0= change in the bond distances and angles in the inner
J coordination sphere of each reactant.

The summation is over the j normal modes of each reactant, and over all
bonds involved in a particular mode.

To apply the above to the l2-tungstocoba1tate system, we must esti-
mate several of the terms in the equation for Ni. For a regular tetra—

hedron such as the central Co(II)04 — Co(III)O4 tetrahedra of these

16

heteropoly ions, only one of the j normal coordinates involves signifi-
cant deformation on going from reactant to product. (28) This is the
one that corresponds most closely to bond stretching. For the others,
the quo's are very nearly zero in a case where the reactants and products
are so similar. In order to estimate the contribution to A, of this
singly degenerate normal mode, we need to know the equilibrium bond
length change Aqo as well as the force constant of the bond in each
reactant. Using the crystal radii of Pauling (29) and his ligancy cor-
rection from six to four, we obtain:

Co(II)—O 1.97 A

Co(III)-0 1.87 R.
It is reassuring to note that X—ray diffraction work has found the
Co(III)-0 distance to be 1.88 H in K5[Co(III)O4W12036] in good agreement
with the above. (37) For consistency we use the two calculated values
and obtain for Aqo the value of 0.10 X. To evaluate the force constants,

we make use of an empirical rule of Badger (30):

1/3 _ _
(c/k) re dij (11)
where: C = an empirical constant equal to 0.125 for a bond be-
tween a first row atom and a third row atom.
di' = an empirical constant equal to 1.06 for a bond between
J a first row atom and a third row atom.
r6 = equilibrium bond length in Angstroms.

K = force constant of bond in megadynes per centimeter.

According to the above rule, the force constants are thus:

Co(II)-0 K 1.66 x 105 dynes/cm.

II

Co(III)-O K 2.27 x 105dynes/cm.

1?

Substitution of the above factors into the equation for Ni yields:
—13
I, = 7.68 x 10 ergs (l2)

Proceeding to evaluate the equations for 10 and A, we note that a1 = a2
for electron exchange reactions, and that r=;:10 H from X-ray dif-
fraction work on K5[C0(III)04W12036].(37) At 00 Ds = 88.0 and n = 1.333
for water, and elez = (8)Ih.8 x 10-10)2 e.s.u.. The value of 8 for the
charge product is obtained from the data on ionic strength variation.

(See Results section of this thesis). Substitution of these values gives:

13

10 = 12.7 x 10‘ ergs (13)
AF*= 10.h kcal/mole (1h)
kr = h.7 x 102 M-lsec._l (15)

This value for the rate constant is for infinitely dilute solution, and

it will be compared to the extrapolated experimental value in the Discus—

sion section of this thesis.

IV. EXPERIMENTAL
A. Preparation of Reagents

The potassium salt of anion I was prepared with reagent grade
chemicals* used without further purification in the method of Baker.(8)
Approximately 0.5 millicurie of 5.2y 60Co was added to the Co(acetate)2
- NHZO. The tracer was obtained from Oak Ridge Nationa1.Laboratory as
the chloride, and was converted to the acetate by repeated evaporations
with acetic acid. Anion II was prepared by electrolysis of a 0.02M
solution of anion I in the anode compartment of an "H" cell. An "H"
cell was used to prevent Hydrogen formed at the cathode from reducing
anion II. The two compartments of the "H" cell were separated by a
glass frit which allows the electrical contact to be maintained. The
electrolysis was carried out at a platinum anode with an applied poten-
tial of h.5 volts. The entire cell was immersed in water at 80°. This
procedure avoids the contamination of the stock solution by chemical
oxidizing agents such as Na28208 or Pb02 which were previously used.(8)
The color change from blue—green to yellow that accompanies the oxidation
serves as a convenient indicator of the extent of reaction. The visible
region absorption spectra of the compounds so prepared agree with those
obtained by Simmons. (33) To convert the stock solutions prepared from

the potassium salts to solutions of the lithium salts, they were passed

 

IKCl; Co(acetate)2°bH20 from J. T. Baker Co.
Na2W04oH20 from Mallinckrodt Chemical Works.

18

19
over a column of Dowex 50 X 12 ion exchange resin which had been converted
to the lithium form. The effectiveness of this procedure was checked by
flame photometric analysis which indicated that less than 10 ppm Na+ and
K+ were present. In addition to exchanging K+ for Li+, this Operation
has the desirable effect of removing any cationic cobalt which may be
present.

The quenching reagent was prepared by mixing 75% by volume of a
solution saturated at 250 with (Matheson, Coleman and Bell) (Bu)4NI and
25% by volume sodium acetate—acetic acid buffer solution 0.1 M in acid
and salt, plus 0.1% by weight "Celite" filter aid. It is of interest
to note that without the presence of the buffer, (Bu)4NI precipitates
both of the heteropoly ions, and in the runs in which extra acid was
added, additional sodium acetate was required also. Early attempts at
using (CH3)4NI in the quench led to higher percentages of induced ex-
change and very finely divided precipitates. The solutions of various
acidities and ionic strengths were prepared by appropriate dilutions
of the stock solutions along with the addition of HCl and LiCl. (Baker
"Analyzed” Reagents) The l—b dioxane used in the dioxane—water mixtures
was distilled from LiAlHq, to remove the peroxides, and stored under
nitrogen. Since the density of dioxane is 1.03%?, volumetric dilution
was suitable for the preparation of the solutions. All water that was
used in this investigation had been distilled from a tin lined still
and passed over a mixed resin bed. Conductivity measurements indicated
a metal ion concentration of less than one part per million in the

effluent water.

20

B. Analytical Methods

The stock solution of anion II was analyzed by potentiometric titra-
tion with ferrous ammonium sulfate which was itself standardized against
potassium dichromate. (N8) The reduction is quantitative and rapid as
can easily be demonstrated by rapidly adding excess ferrous ion to a
solution of anion II and noting the color change. The color is not
sufficiently intense,however, to serve as an indicaton Therefore,a
platinum indicator electrode and a saturated calomel reference electrode
were used to follow the reaction potentiometrically. The anion I solu-
tions were similarly analyzed by first oxidizing an aliquot electrolytic-
ally as described in the preparative section, and then titrating as for
the anion II solutions. This method has several advantages over anal—
ysis on the solid material. It requires no knowledge of the somewhat
uncertain number of cations or protons present in these acid salts, or
of the number of waters of crystallization, but gives the concentrations
of the anions (I or II) directly.

The radiochemical assay of the separated precipitates was done with
a windowless gas—flow counter using helium-isobutane as the counting gas
and an applied potential of lbOO volts.

Although anion II is a strong oxidizing agent, its solutions were
found not to decompose with time, as long as dust and other easily ox-
idized materials were carefully excluded, and the solutions were not
heated in the presence of chloride ion which may then reduce them. The
solutions of anion II were easily checked by determining their absorbancy
at 638 mu where anion I has a substantial absorption maximum and anion II

is tranSparent. (See Appendix II)

21

C. Procedure

The reaction was initiated by the addition of 1 ml. of anion II solu-
tion to 1 ml. of anion I solution. The pipets used were stored at 0°
by placing them in a cylinder which was immersed in a bath at 0.0° i 0.1
The solutions were delivered in blackened test tubes to prevent photo-
catalysis, although no evidence for such catalysis was found. Each
point in a run represents an individual reaction mixture since some of
the half-times were short and several minutes were required to carry out
the separation. Anion II was selectively precipitated from the reaction
mixture by the rapid addition of 1 ml. of quenching reagent stored
at 0°. (The composition of the quenching reagent is described above.)
Two-tenths of a minute were allowed for the complete formation of the
precipitate, and then the mixture was filtered by suction on a 2 cm.
circle of Whatman #SNO filter paper in a funnel.with a removable chimney.
At temperatures above 0° slightly less quench was required in order to
prevent the precipitation of both anions. The quality of the separation
can readily be judged by the color of the precipitate, which should be
light yellow. If coprecipitation of anion I is extensive (as it is if
the separation is attempted at too low a pH) the color of the precipitate
becomes green. The devising of a rapid, reproducible, separation of
anion I and anion II, which are large and identical except for one unit
of charge, was the most difficult experimenta1_part of this investiga-
tion. No separation has previously been reported. After filtration,
the precipitates were air dried on the filter papers, which were kept
from curling by placing small lengths of 2 cm. diameter copper tubing

over them. The radioactivity of the dried samples was counted. The

22
infinite-time samples were allowed to equilibrate, and then quenched and
filtered after more than ten half-times had elapsed. Separation induced
exchange to the extent of about 20% was observed in most series. All
operations were carried out in as standard a manner as possible in order
to minimize random errors. Weighing of the precipitates in order to
determine the absolute activity did not lead to increased precision in

the results and this procedure was abandoned.

23

D. Errors

The sources or error in this investigation are of two types, the
usual random errors of laboratory work and the errors associated with
radiochemical assay. Radiochemical assay errors or counting errors
are dealt with in detail in texts on radiochemistby. (31,32) The sali-
ent points are these: 1) radioactive decay follows a Poisson distri-
bution law; 2) for such a process the probable error is determined by
the number of counts rather than by the length of the counting period;
3) the probable error can be reduced below and given value by observing
a sufficient number of counts. In this study, approximately 10,000
counts were measured for each sample. It can be shown that this leads
to a standard deviation in the number of counts observed of i 1%.

Unfortunately, random errors from another source were larger and
more difficult to control. The most serious source of error was the
separation method itself. Both positive and negative errors are pos-
sible since the precipitation of anion II may be incomplete, or the
precipitate may be contaminated by occlusion and coprecipitation of the
anion I compound. Because of these factors, a single point in a given
kinetic run has a relative standard deviation of about i 7%. Although
attempts by various methods to reduce this scatter were not particularly
successful, the rate constants are not in error by this amount for two
reasons: 1) the rate is calculated from the slope of a straight line on
a semi-log plot and not from individual points; 2) the infinite time
point which does affect the Slope (see part A of this section) was deter-

mined in triplicate for each run.

211

In addition to the two sources of error discussed above, there were
the normal errors of quantitative procedures, and the errors in the
measurements of times and temperature. The timing errors were probably
the least significant, as shown by the linearity of the data for a
series in which the half—time was only 0.73 minutes. In most runs, the
temperature was held at 0.0° i 0.1 with ice in equilibrium with water
at 0° in a dewar flask. Similar control was possible for the runs at
slightly higher temperatures by the use of an insulated refrigerator
bath controlled by a bimetallic probe and an electric relay, in which

the regulation was to at least i 0.1°.

V. RESULTS
A. Determination of the Kinetic Order

An essential first step in any kinetic study is the establishment
of the rate dependence on the concentration variables. Although electron
exchange reactions are most often second order, first order in each
reacting species, there are several other cases. The Ag(I) - Ag(II)
exchange (3b) and the U(IV) — U(VI) exchange (A3) are well known examples

where a simple second order rate law does not hold. By assuming that:
R = kr[I][II] (16)

and determining R from the McKay equation (see section III—A of this
thesis) the rate constants of table I were calculated for various con-
centrations of anion I and II. The half—times used in the McKay equa-
tion to calculate R were obtained from plots such as figure 1 which
shows some typical rate data. From the variation in the rate constants,

it was found that

R = kr[I]l.oi.01[II]l.oi.Ol (17)

For the concentrations and conditions listed in table I, the half-

time was approximately 11 minutes.’

25

26

Table I. Dependence of the rate constant on reactant concentrations.

 

Ionic Strength Adjusted to 0.6 with LiCl; Temperature 0°; pH = 2;

Li salts of anions I and II

 

 

[I] [II] kr(m?1 sec.'1)
0.0007 0.0018 0.63
0.0009 0.0015 0.6u
0.0011 0.0012 0.63
0.001u 0.0010 0.65
0.0016 0.0007 0.61

 

0.63 i .016

 

 

27

0.5

0.1
O 2 4 6 8 IO 12

TIME WIN.)
Figure 2. Exchange curve for some typical rate data.
Temperature 00; 9.h1 LiCl; [1] = 1.18 x 10‘3M;
[11] . 1.22 x 10 3m; kr = 0.uu (m‘l sec."1 .—

28

B. Dependence of the Rate on Hydrogen Ion Concentration.

Because the heteropoly acid salts used in this investigation can
ionize to varying degrees depending on the ion atmosphere, it is import-
ant to know how the hydrogen ion concentration affects the rate of ex—
change. The hydrogen ion concentration.was varied from 0.025 M to
0.80 M with HCl while the ionic strength was maintained constant at
1.02 with LiCl. The acid-salt pair of HCl—LiCl is a useful one because
the mean activity coefficients for these two materials are nearly ident-
ical throughout this concentration range in water solution. Within
experimental error of i 6% in the rate constant, the exchange rate was
independent of hydrogen ion conentration at constant ionic strength.
(Table II below). This indicates that protons are not a necessary part
of the activated complex, and that neither a Grotthuss.type of proton
jump mechanism, nor H atom transfer contributes significantly to the rate,
since the ionic mobility of H+ is 3&0 cm./Sec/volt while that of Li+ is

only 37 cm./Sec./volt. (35).

Table II. Dependence of rate on HCl concentration.

 

 

Ionic strength 1.02 adjusted with LiCl; [I] = 1.08 x.10-3M; [II] =
1.22 x lO-sM; Temperature 0°.

 

[HCl] kr(M-lsec.—l)
0.025 1.05
0.200 0.98
0.1100 1.03
0.600 0.93

0.800 1.09

 

29

C. Dependence of the Rate on Ionic Strength

The rate was found to have a marked dependence on the ionic strength.
This is not surprising in view of the large charges of like Sign on the
ions. Simple coulombic considerations would predict that increasing the
density of the ion atmosphere around such ions would aid them in ap-
proaching each other close enough to allow electron transfer to occur.
The ionic strength was varied from 0.015 to 0.81 by the addition of LiCl,
and the resulting effects on the rate constants are given in table III.
Results of this nature may be described by the Br0nsted-Bjerrum equation
(36) if the ionic strength is very low. For large ions at moderate
ionic strengths, a plot of log kr y_s_ Wwill not be linear as would be
the case if the Bransted-Bjerrum equation were followed. (Figure 3).
We note, however, that the deviation from linearity is in the direction
expected if the additive term in \rfi in the denominator of the Debye-
Hfibkel is neglected as is done in the Br6nsted—Bjerrum approximation.
Alternatively, the same data II' may be interpreted with the Marcus
equations (Section II-A). This treatment indicates that a plot of
logkr XE e-Ikr (Figure A) should also be linear but this relationship
will hold for somewhat higher ionic strengths than the limiting Br0nsted—
Bjerrum expression.

From the slope of the line in figure h, a charge product of 2122 = 8
can be calculated. This would seem to be a reasonable value for ions

in solution that have formal charges of five and six.

 

30

Table III. Dependence of the rate constant on ionic strength.

 

A.._._

Ionic strength adjusted with LiCl;-Temperature 0°; [I] = 1.11.1 x lO'SM;
[II] = 1.22 X 10 M'

 

 

11 W 'Kr e-Kr kr(_I\_’1.'l sec._1)
0.015 0.122 0.396 0.670 0.03
0.05 0.22h 0.727 0.h81 0.07
0.21 0.h59 l.h89 0.225 0.31
0.u1 0.6h1 2.080 0.125 0.uu
0.61 0.783 2.5hl 0.079 0.6h
0.81 0.901 2.92h 0.053 0.96

 

RATE CONSTANT (M'.‘ 350'!)

[.0

0.5

5

31

 

 

 

.01 I 1 1 1 1 I I I I 1
0.0 0.5 1.0
SQUARE ROOT IONIC STRENGTH

Figure 3. Graph of log kr is «11.

Temperature 0°; Ionic strength adjusted with LiCl;
[I] = 1.11. x 10‘3N; [11] ~— 1.22 x10_3M.

0.5

RATE. CONSTANT (MT'SECI)

.0I

L0
[\7

0.5
-K r-
63

Figure 11. Graph of log kr vs e—Kr.

0.l

Kr == 3.211” for water solutions at 0°; Ionic strength

adjusted with LiCl; [III :

1.1M x 10‘3N; [111

1.22 X 10 PM.

33

D. Dependence of the Rate on Temperature.

The temperature dependence of the rate was studied at several ionic
strengths and the parameters of activation were obtained. The activation
energy is a function of ionic strength and an approximate extrapolation
to zero ionic strength was possible. (Table IV) The dependence of the
energy and entropy of activation on ionic strength for reactions between
ions in solution, is a matter frequently ignored in experimental papers
on the subject, a fact which undoubtedly has lead to fallicious con-
clusions when comparisons between various systems are made. The results
of this investigation demonstrate the futility of any such conclusions
that are based on slight differences in the activation parameters, es-
pecially when they are for solutions of different strength. The data
on the temperature dependence of the rate may be interpreted by the

transition state equation (36):

kr = (kT/h)exp[ASI/R - AHI/R T] (18)
where:

kr = the rate constant

k = Boltzmann's constant

T = the absolute temperature

AS’ = the entropy of activation per mole

AH‘ = the enthalpy of activation per mole

R = gas constant per mole

Ea = activation energy per mole

for a reaction in solution:

AH... = Ea - RT (19)

3h
The activation energies were obtained from the slope of the lines in
Figure h,but the entropies were calculated from the above equation be-
cause the extrapolation of 1/T to zero is very long. The results are

summarized in Table IV.

Table IV. Dependence of the rate constant on temperature and ionic
strength. The parameters of activation.

 

 

Ionic strength adjusted with LiCl; [I] = 1.1h X 10-3M, [II] =

 

 

 

1.22 x 10'3m.
u RIM—lsec.-1) Temp. Ea kcal/mole AS e.u.
0.21 0.31 00.0
0.21 0.u8 12.7 5.0 t 0.5 -uh.0 r 2
0.21 0.72 25.0
0.05 0.07 . 00.0
0.05 0.13 12.7 8.0 t 1 -35.1 t 1
0.05 0.3h 25.0
0.015 0.03 00.0
0.015 0.08 12.7 12.8 1 2 -19.7 r 1
0.015 0.2b 25.0

0.00 -- -- 18(extrapolated)

 

35

.000 L

Upper Curve, ionic strength 0.21
Middle Curve, ionic strength 0.05

Lower Curve, ionic strength 0.015

 

 

800
3
KrIO
T
I.0
0.5
I
0 .l
12 as 18

3
IO/T
Figure 5. Graph of log kw103/T IE 103/1.

Ionic strength adjusted with LiCl; [I] = 1.1L x 10‘3m;
[11] = 1.22 x 10 3H-

36

E. Specific Cation Effects

Substitution of the lithium cations by potassium ions lead to large
increases in the rate of electron exchange, as shown by the data in table V.
At 0°, [I] = 1.1h x 10-°M, [11] = 1.22 x 10'3N; and [KCl] = °-3.H, the
half-time is only 0.73 minutes. This represents the approximate lower
limit of half-times which could be measured by the present techniques.

In figure 6 the effects of lithium and potassium ions are compared. We
note that not only is the rate constant larger in the presence of potas-
sium, but the rate of increase with concentration is greater. Attempts

to.prepare solutions of similar heteropoly ion concentrations to those

 

of table V in the presence of even 0.1 M cesium ion led to extensive

precipitation of the heteropoly ions.

Table V. Dependence of the rate on various cations.

 

 

Temperature 0°; [1] = 1.11 x 10‘3N; [11] = 1.22 x 10’3g; KI as KCl.

A_A_

 

Cation Concentration kr(Mtlsec.-l)

Li 0.015 0.03
Li 0.05 0.07
Li 0.21 0.31
Li 0.11 0.hb
Li 0.61 0.6h
Li 0.81 0.96
K 0.02 0.7A

0.10 2.21
K 0.30 6.67

 

37

$0-2 x mm; .. E: Mamba x 1:4 u E Moo 838683
.x .3 n z .22.: m 2388 88 0o cameo .6 Sara
2: zo_k<¢hzmozoo 20:40
0.0 ¢.o N6

0 .o 0.0

 

 

 

V

lNVISNOO BIVU

.11)

HUHA l m>bflU 9030A

SK 1 05:50 noun:

{939

I

38

F. Dependence of the Rate on Dielectric Constant

In order to determine the effect of dielectric constant on the rate,
several series of experiments were made in dioxane-water mixtures. Di-
oxane is commonly chosen for such studies because it is miscible in all
proportions with water, and it has a very low dielectric constant itself
so that a wide variation may be obtained. The dielectric constant of
the mixtures approximates a linear function of the weight percent di-
oxane over most of the range. For the sake of comparison with other data
from this investigation, it was desirable to work at 0°. The dielectric
constant of dioxane-water mixtures is a function of temperature however,
and experimental values are not available for 0°. It was possible to
estimate them none-the-less. Data at 15°, 25°, 35°, and A50 for dioxane-
water mixtures are tabulated for mixtures of various concentrations. (Ah)
By extrapolation of the values for each mixture from 15° down to 0°, a
set of data were obtained for 0° from which a plot could be made of
weight percent dioxane vs. dielectric constant. (Figure 7) AS may be
seen from the graph, the known value for the dielectric constant of-
water at 00 falls on the same straight line as the extrapolated values.

As one might expect for a reaction between like charged ions, de-
creasing the dielectric constant decreased the reaction rate. More
quantitatively, the Marcus theory predicts that a plot of 1n kr XE'
l/DS should be linear with a slope proportional to 2122, the charge
product of the ions in solution, if the slight dependence of A on Ds

is neglected. From Figure 8, the charge product is found to be 30.

 

Dielectric Constant

'39

 

 

 

 

4 l 1.
0 20 NO 60
Weight Percent Dioxane
Figure 7. Weight percent dioxane XE' dielectric constant at 0°,

 

 

to

Table VI. Dependence of the rate on the dielectric constant.

 

 

Temperature 0°; Ionic strength 0.6, adjusted.with LiCl; [I] = 1.1h x
10 .Hi [11] = 1.22 x 10'3g.

 

 

 

Volume % Dioxane DS l/DS kr(M-lsec.-l)
0.0 88.0 0.011h.. - 0.63
h.0 8h.3 0.0119 0.52
20.0 69.3 0.01Ah 0.36

32.0 57.0 0.0175 0.26

 

A1

 

.ma-6a x NN.H u mHHL Mme-oa x;aawa u fiHP woo oeoooeooan
maoaa new: oopmohom 6.0 eaoooaem oaooH . g\a .ma a moa Go eamao .w panama

m
o\
wao.o NaO.o I oao.o mao.o _..aao.o mao.o mao.o Hao.o

J . . I q . O

 

 

d.a

 

 

 

VI. DISCUSSION

It has been found, (33) by cryoscopic studies in Na2504'10 H20
(Glauber's salt) that the heteropoly ions used in this investigation
show very little tendency to dissociate into W04= or other simple units.
This fact, plus the number of bonds between groups implies substitutional
inertness with respect to W04= groups and leads one to believe that
electron exchange between such ions can only take place via an outer-
sphere activated complex in acid solution. If the pH is raised above 7,
degradative dissociation does occur with any heteropoly compound. The
postulate of an outer—sphere activated complex is.supported by the lack
of hydrogen ion dependence of the rate, by the negative entropies of
activation (~19 to ~NN e.u.), and by the simple kinetic order, although
individually these factors are not conclusive. Taube (17) has pointed
out the difficulties involved in attempting to make parameters of acti-
vation diagnostic of mechanism. In this sytem, however, the evidence
is not contradictory and the assignment of an outer—sphere mechanism
seems appropriate.

The magnitude of the rate constants for this system as well as the
possible mechanism is deserving of comment. Most outer-sphere reactions
which have been investigated are very rapid, particularly in anionic

systems. For example the reaction: (N5)

W(CN)e-3 + 11(c11)...’4 = W(CN)8_3 + w(c1\1)8'4
has a 105 < kr < 108 M-1 sec.-1 as determined from electron spin reson-

ance experiments. The reaction: (N6)

N2

 

N3

* -3 -4 -3 * -4
Fe(CN)6 + Fe(CN)5 = Fe(CN)5 + Fe(CN)5

has been studied by flow methods and a k = 3.5 x 102 yflsec._1 was
found.

In the light of the above, it is noteworthy that the 12—tungsto-
cobaltate (II - 111) system exchanges slowly enough to allow study by
conventional techniques. Therefore, a comparison with the theory of
Marcus, which is derived for a weak-overlap type of mechanism, is of
interest. A summary of the necessary calculations may be found in the

section on theory50f'this thesis. Several approximations were made in

 

order to obtain a numerical result. From the Marcus theory: (for
infinitely dilute water solution)

kr = N.7 x 102 M_'lsec.-l

By extrapolation of the log kr XE: e-KI‘ to infinite dilution one obtains
kr(eXpt.) = N.5 x 10":"'I‘_’I"lsec.-l

The discrepancy between the calculated and the experimental values is
substantial, and may be caused by a number of factors. The estimate

made of the rearrangement energy necessary for reaction is probably on

the low side, since in the approximation used, distortion of the ions

as a whole was neglected and attention was concentrated on the central
atom and its nearest neighbors. Furthermore, the Marcus treatment as—
sumes that the transmission probability for the electron is unity within
the activated complex. This condition may not hold in heteropoly ions
where a large number of oxygen and tungsten atoms are present,(see Figurel)

which may serve to "insulate" the central cobalt atoms from each other.

NN
Finally, there is some uncertainty in what value to use for the charge
product of the ions in solution. This matter is discussed in more detail
in the latter part of this section. Although the mechanisms may not
be the outer-sphere type, it is of interest to note that the following

reactions are very slow: (1)
00(NH3)6+2 + Co(NH3)6+3 = 00(NH3)6+3 + Co(NH3)6+2

- - "" -4 -3
Co(C204)3 3 + Co(C204)3 3 = Co(C204)3 + Co(C204)3

This may be the result of a large change in the cobalt ligand bond
length necessary upon changing oxidation state, or a consequence of the
change in the number of unpaired electrons. (Octahedral cobaltous com-
plexes are generally spin free while octahedral cobaltic complexes are
spin paired.) This situation does not occur in the present system where
both complexes are tetrahedral and spin free. (33) The observed ex-
change rate is probably affected by several of the factors discussed
above and a determination of the contribution of each is not possible

at the present time.

While the gross magnitude of the rate is determined by the consid—
erations discussed above, variation from kr = 10-2 to 101M_lsec.—1 was
possible as a result of changes in experimental conditions. In particu-
lar, changing the cation from lithium to potassium leads to a large

increase in the rate of exchange. In the range studied, the rate is

linear in cation concentration, (see Figure 6) so that one might write:
R = krIM+IIIIIIII

Such a rate law implies, however, that exactly one cation enters the

activated complex, which is not necessarily true for an outer-sphere

 

NS

mechanism where the activated complex is of uncertain structure. The
effect appears to be specific for each cation which might lead one to
attribute the results to differences in ionic mobilities in the ions.

The ionic mobility of the potassium ion is 73 cm/Sec./volt, while
that of lithium is 37 cm/sec./volt. From Table V, however, we see that
for a given concentration of cation, the rate constant for potassium is
always at least ten times that of lithium. Thus, the rate increase is
not linearly associated with the ionic mobilities. It seems preferable
to attribute the effects of different cations to the greater ion pairing
abilities of the larger ions. Ion pairs can facilitate electron exchange
by decreasing coulombic repulsion, and the number of ion pairs will in-
crease rapidly with concentration, as does the rate.

The result of ionic strength variation on the rate may be interpreted
in several ways. The Brbhsted-Bjerrum equation is often applied to pre-
dict the effects of ionic strength for ionic reactions. In the case of
large ions and only moderately low ionic strengths, however, this equa-
tion is not followed. (Figure 3) The equations of Marcus provide
another means of interpreting the data, and the requirements are not as
stringent. Figure N shows that the linear relationship predicted for

er is followed quite well, and from the slope of the line,

log kr gs. e-
a charge product in solution, zlzz = 8, may be computed.

The charge product can also be calculated from the rate data obtained
as a function of dielectric constant. (See Figure 8). In this case,
a value of zlzz = 30 is obtained. This value does not agree with the

value of eight, obtained from the ionic strength variation data. It

seems probable that the value of eight is the more reliable one, since

 

N6

the validity of the Debye-chkel treatment in solutions of low dieleetric
constant and moderate ionic strength is questionable.

An interpretation consistent with the results in water solution is
that several cations are closely associated with heteropoly ions even
in very dilute solution. These ion groups may then act as a strong
electrolyte of somewhat lower charge. Such a conclusion is corroborated
by the fact that even unsymmetrical 3-1 electrolytes of complex ions
are not "strong" in several cases which have been examined. (50,51) The
Situation in the mixed solvent system of water—dioxane is more complicated,

and defies a clearcut explanation, at present.

 

In summary, this investigation has led to the following results

and conclusions. The rate of electron exchange between the l2-tungsto-
cobaltate (II) and the 12-tungstocobaltate (III) anions is moderately
rapid but slow enough to be followed by classical techniques. The ex-
changing electron is able to penetrate the tungsten-oxygen "cage" of-the
heteropoly ion, and since the "cage" itself is very stable, the reaction
must proceed by an outer-sphere mechanism. The rate law for the exchange
is second order, first order in each ion. Hydrogen ions were found to
have no special function in the activated complex, as evidenced by the
fact that the rate was unaffected by the hydrogen ion concentration at
constant ionic strength. In solution the heteropoly ions probably exist
in close association with cations so that their high charge is partially
neutralized. Thus the presence of K+ which tends to form ion pairs more
readily than Li+ greatly accelerates the exchange rate. The temperature
dependence of the rate is summarized in the thermodynamic parameters of

activation which are within the range of values typically observed for

11?

exchange reactions that proceed via an outer-sphere mechanism. The
large negative values for the entropies of activation and the catalytic
effects of cations, suggest that the closeness of approach of the react-
ants is very important in determining whether or not electron exchange
occurs.

The comparison of the experimental data with the predictions of
the Marcus theory is inconclusive. The plots which are predicted to
be linear, particularly log kr _\_I_s_. e- Kr are fit fairly well by the data,

but the theoretical and experimental rate constants are not in agree-

ment. The decision as to whether this is the result of: substitution

 

of incorrect values for some of the parameters into the equations, mis—
application of the theory altogether, or simply an inadequate theoretical
treatment, will have to be deferred until data from more experimental

systems have been compared to the Marcus theory.

10.

ll.

12.

13.

IN.

15.

VII. LITERATURE CITED

Lewis, J., and R. G. Wilkins, ed. "Modern Coordination Chemistry,"
Chap. 2, by D. R. Stranks, p. 78-173. Interscience, New York,
1960.

Sheppard, J. C. and A. C. Wahl, J. Am. Chem. Soc. 72, 1020 (1957).
Keggin, J. F., Proc. Roy. Soc. lNN, 75 (193N).

Marcus, R. A., J. Chem. Phys. 2N, 966 (1956).
J. Chem. Phys. 26, 867 (1957).
J. Chem, Phys. 25, 872 (1957).

Marcus, R. A., J. Phys. Chem. 61, 853 (1963).

Sidgwick, N. V. "Chemical Elements and Their Compounds," Vol. 2,
p. 10N2. Oxford, 1950.

Kline, C. H., "Heteropoly Anions in Science and Industry,” Symposium
on Structure and Properties of Heteropoly Ions, 130th National
Meeting American Chemical Society, Atlantic City, N.J., Sept.

17, 1956.

Baker, L. C. W., and T. P. McCutcheon, J. Am. Chem. Soc. 78, N503
(1956). "

Baker, L. C. W., and V. E. Simmons, J. Am. Chem. Soc. 81, NTNN (1959).

Kirschner, 3., ed. "Advances in the Chemistry of Coordination
Compounds," p. 60N. Proceedings of the 6th I.C.C.C. Wayne
State U., Detroit, Mich., Sept. 1961. Macmillan, New York,
1961.

Cotton, F. A., ed. "Progress in Inorganic Chemistry," Vol. N,
Chap. N, "Isopolytungstates", by D. L. Kepert, p. 199.
Interscience, New York, 1962.

Souchay, P. Pure Applied Chem. 6, 61 (1963).

Pascal, P., ed. "Nouvou Traite De Chemie Minerale," V01 XIV,
"Heteropolyacids" by L. Malaprade, p. 903. Masson et Cie,
Paris, 1959.

Wahl, A. C., and N. A. Bonner, ed. "Radioactivity Applied to
Chemistry," Chap. 1, by 0. E. Meyers and R. J. Prestwood, p. 6.
Wiley, New York, 1951.

Libby, w. P., J. Phys. Chem. 56, 863 (1952).
J. Chem. Phys. 3E, N20 (1963).

N8

 

16.

17.

18.
19.

20.
21.
22.
23.
2N.

25.
26.

27.

28.

29.

30.

31.

32.

33.
3N.
35.

N9

Marcus, R. J., B. Zwolinski, and H. Eyring, J. Phys. Chem. 58, N32
(195N) . '—

Emeleus, H. J. and A. G. Sharpe, ed., "Advances in Inorganic Chem-
istry and Radiochemistry," Vol. I, Chap. 1 by H. Taube, p. 1-53.
Academic Press, New York, 1959.

Stranks, D. R., and R. G. Wilkins, Chem. Rev. 51, 7N3 (1957).

"Oxidation-Reduction Reactions in Ionizing Solvents" -- A General
Discussion, Discussions Faraday Soc., Vol. 29, 7-169 (1960).

Halpern, J., Quart. Rev. (London) Vol XV, No. 2, 207 (1961).
Brubaker, C. H., Record Chem. Progr. EN, 181 (1963).
Prestwood, R. J. and A. c. Wahl, J. Am. Chem. Soc. _7_1, 3137 (19119).

Mckay, H. A. C., Nature 1N2, 997 (1938).

 

Baker, L. C. W., Chemical Data Series, Bulletin Cdb-l2, Climax
Molybdenum Co., 1960.

Hush, N. 8., Trans, Faraday Soc. 57, 557 (1961).

Laidler, K. J., Can. J. Chem. 31, 138 (1959).
Sacher, E., and K. J. Laidler, Trans. Faraday Soc., 52, 396 (1962).

Symons, M. C. R., J. Chem. Soc. 3676 (195N).

Herzberg, G. "Infrared and Raman Spectra of Polyatomic Molecules,"
p. 100. D. Van Nostrand Co., New York, 19N5.

Pauling, L. "Nature of the Chemical Bond," p. 518,538. Cornell
University Press, Ithaca, New York, 1960.

Badger, R. 11., J. Chem. Phys. 2, 128 (1931.).

J. Chem. Phys. 3, 710 (1935).

Friedlander, G. and J. W. Kennedy "Nuclear and Radiochemistry,"
John Wiley and Sons, New York, 1955.

Overman, R. T. and H. M. Clark "Radioisotope Techniques," McGraw-
Hill Co., New York, 1960.

Simmons, V. E., Ph.D. Dissertation, Boston University, 1963.
Gordon, B. M., and A. C. Wahl, J. Am. Chem. Soc. 89, 273 (1958).

Hamer, W. J., ed. "The Structure of Electrolyte Solutions," Chap. 5,
by M. Eigen and L. Dtheyer, p. 66. Wiley, New York, 1959.

36.

37.
38.
39.
N0.
N1.
N2.
N3.
NN.

N5.
N6.
N7.
N8.

N9.
50.

51.

52.
53.
SN.

50

Frost, A. A. and R. G. Pearson "Kinetics and Mechanisms," 2nd Ed.,
p. 100. Wiley, New York, 1961.

Eriks, K. and N. F. Yannoni, gt;_§l. Acta. Cryst. 23, 1139 (1961).
Silverman, J. and R. W. Dodson, J. Phys. Chem. 56, 8N6 (1952).
Hudis, J. and R. W. Dodson, J. Am. Chem. Soc. 12, 911 (1956).
Hudis, J. and A. C. Wahl, J. Am. Chem. Soc. 15, N153 (1953).
Sheppard, J. C. and L. C. Brown, J. Phys. Chem. 61, 1025 (1963).
Sutin, N., J. Phys. Chem. 62, 1766 (1960).

Rona, E. R., J. Am. Chem. Soc. 12, N339 (1950).

Harned, H. S. and B. B. Owen. "The Physical Chemistry of Electrolytic
Solutions", 3rd. Ed., Reinhold, New York, 1958, p. 713.

 

Weissman, S. I. and C. S. Garner, J. Am. Chem. Soc. 22, 1072 (1956).
Wahl, A. c. and c. D. Deck, J. Am. Chem. Soc. lg, N05N (195N).
Eichler, E. and A. C. Wahl, J. Am. Chem. Soc. 29, N1N5 (1958).

Blaedel, W. J. and V. W. Meloche. "Elementary Quantitative Analysis,"
2nd Ed., p. N73. Harper and Row, 1963.

Marcus, R. A., Private communication.

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82, 1111 (1960).

Groves K. 0., J. L. Dye, and C. H. Brubaker, J. Am. Chem. Soc. 82,
NNNS (1960).

Griffiths, J. S. and J. Owen, Proc. Roy. Soc. A 22g, 96 (195N).

Ebsworth, E. A. V. and J. A. Neil, J. Phys. Chem. é3, 1890 (1963).

Booth, H. 8., Ed., "Inorganic Synthesis," McGraw—Hill, New York,
1939, p. 129.

 

APPENDIX I

ORIGINAL KINETIC DATA

51

52

Dependence of rate on reactant concentrations.

 

[I] = 1.1N x 10’3M; [II] = 1.22 x lO'SM; pH = 1.6; 0.6g LiCl; Temper-
ature 0°; tl/é = 8.0 min.

 

 

 

t (min.) X (c/6 min.) X0) - X*
0.12 1571 10926
1.50 3279 9218
3.01 3999 8N98
N.51 5155 73N2
6.01 5833 666N
8.03 7130 5367

10.00 7338 N659

12.02 7981 N516

1N.Ol 9307 3190

00 12N97 --

 

 

[I] = 9 x 10-?M; [II] = 1.5 x lO-SM; pH = 1.6; 0.6M LiCl; Temperature

 

0.15 1N82 10917
0.21 1610 10789
1.0N 2550 98N9
3.01 3635 876N
5.02 5507 689N
7.02 6N78 5921
00 12399 --

 

[I] = 7 x 10-IM; [II] = 1.8 x lOTIM; pH = 1.6; 0.6 M LiCl; Temperature
0°; tl/2 =,8‘O min.

 

0.15 1225 8506
2.02 2516 7215
N.ON 3338 6393
6.03 N52N 5207
8.01 SN25 N306
10.00 SN96 N235
12.01 5532 N199
1N.01 7315 2N16
00 9731 --

 

*The value of Xd) is experimentally determined. The listed value

is the average of triplicate determinations.

 

53

Dependence of rate on reactant concentrations. (Cont.)

 

 

t (min.) X (c/6 min.) X0) - XII

 

[I] = 1.N x 10'3M; [II] = 1.0 x lOTSM; pH = 1.6; M LiCl; Temperature 0°;
tl/2 = 7.7 min.

 

 

 

0.13 1939 11236
2.01 NN27 87N8
N.01 S998 7177
6.01 7123 6052
8.01 7370 5805
10.00 801A 5161
12.01 93N2 3833
1N.00 10578 2597
00 13175 _-
[1] = 1.6 x 10'3m; [11] = 7 x io'fm; pH = 1.6; g LiCl; Temperature 0°;
t1/; = 8.2
2.02 NO0N 8226
5.02 5199 7031
8.26 7393 N837
10.N2 8158 N072
12.82 8725 3505
CD 12230 --

 

 

SN

Dependence of rate on hydrogen ion concentration.

W

[1] = 1.08 x 10'3g; [11] = 1.22 x io‘ém; 0.025 N HCl; 1.0 g LiCl;
Temperature 0°; tl/Z = N.8 min. _

 

 

t (min.) X (c/6 min.) . Xm - X
0.13 2N32 15697
0.72 3826 1N303
1.57 5972 12157
2.91 8507 9622
5.03 10889 72110

.10. 17 111265 38611

CD 18129 --

 

[I] = 1.08 x.10‘3M; [II

_ 1 = 1.22 x 10‘3m; 0.8 N HCl; 0 2.! LiCl;
Temperature 0°; tl/2 = N.6

 

 

0.12 11393 12100
0.N8 11881 11612
1.03 1N691 8802
1.53 1N930 8563
2.02 13928 9565
3.05 1N558 8935
5.07 l7N7N 6019
8.03 19779 371N
a, -_

23N93

 

[I] = 1.08 x lO‘éM; [11] = 1.22 x io‘ém; 0.6 N HCl; °-b.N Lic1; Temper-
ature 00; tl/Z = 5.11 min.

 

0.15 3653 12692
0.52 N52N 11821
1.06 5N3? 10908
2.59 6933 9N12
5.0N 9118 7227
7.23 11558 N787
10.12 12872 3N73
on 163N5 --

 

SS

Dependence of rate on hydrogen ion concentration. (Cont.)

 

t (min.) f X (0/6 min.) XGD - x

 

[I] = 1.08 x 10'3M; [II] = 1.22 x 10*?3; 0.1 M HCl; 0.6 g 1101; Temper-
ature 0°; tl/2 = 5.9 min.

 

0.18 1988 13538
0.62 3267 12259
1.11 1109 11117
2.09 5129 10097
3.07 6765 8761
1.09 7259 8267
5.08 9077 6119
5.11 9106 6120
6.95 10316 5210
8 . 59 10596 1930
00 15526 --

 

 

[I] = 1.08 x 10'31; [II] = 1.22 x 10'33; 0.2 1 H01; 0.8 y 1101;
Temperature 0°; tl/Z = 5.1.

 

0.12 1820 15189
0.52 2865 11511
1.05 3516 13833
1.55 1518 12831
2.06 5076 12303
3.12 6930 10119
1.10 7191 10185
5.08 9567 7812
6.08 9123 8256
8.05 10983 6396
00 17379 -—

 

Dependence of rate on ionic strength

 

 

[I] = 1.11 x 10'§y; [II] = 1.22 x 10'31; 0.8 y 1101; Temperature 0°;
tl/Z = 5.1 min.

 

0.12 1928 13107
0.51 3210 12095
2.08 5286 10019
1.01 7718 7587
7.01 10115 5190
10.02 13133 2202
00 15335 ~-

 

56

Dependence of rate on ionic strength. (Cont.)

 

 

“_1 (min.) X (c/6 min.) Xa, - X

 

[I] = 1.11 x 10-311; [II] = 1.22 x ‘10-32’1; 0.6 _N_I LiC1; Temperature 0°;
tl/Z = 7.6 min.

 

A 0.12 2193, 13311
0.65 3010 12191
1.07 3320 12211
2.08 1502 11032
3.07 1826 10708
5.01 6506 9028
7.01 7517 7987
9.07 9897 5637

11.57 11807 3727

13.71 11690 3811

Go 15531 --

 

[I] = 1.11 x 10'31; [11]

1.22 x 10'31’1; 0.1 fl LiCl; Temperature 0°;

 

tl/2 = 11.0 min.
0.15 2181 12895
0.77 3019 12027
1.52 3611 11162
3.06 1910 10136
5.06 6021 9052
7.05 6558 8518
10.07 8261 6812
13.85 9826 5250
18 . 92 10752 1321
a) 15076 --

 

[I] = 1.11 x 10‘§g; [II] = 1.22 x 10‘31; 0.2 g 1101; Temperature 0°;

t1/2 = 15.7 min.

 

0.12 2330 12067
1.11 3238 11159
3.08 1329 10068
5.01 1579 9818
8.02 6655 7712
11.07 7376 7021
11.02 8102 6295
00 11397 ..

 

 

5?

Dependence of rate on ionic strength. (Cont.)

 

f

w— v

t (min.) x (0/6 min.) XCID - X

 

[I] = 1.11 x 10'8M; [II] = 1.22 x 10'91; 0.0 fl LiCl; Temperature 0°;

tlfi = 180 min.

 

0.13* 3293 6221
2.51 3599 5918
7.62 3666 5851
15.07 3876 5611
25.06 3619 5868
30.01 1111 5106
16.11 1362 5155
60.05 1818 1699
92.78 5000 1517
131.01 5987 3530
CD 9517 --

 

[I] = 1.11 x 10'31; [II] = 1.22 x 10'31; 0.05‘1 1101; Temperature 00

tl/z = 66.11 min.

 

1.18 3325 8218
5.31 1206 7367
10.31 5051 6519
25.22 5585 5988
12.06 6586 1987
61.10 7220 1353

CD 11573 ——

 

Dependence of rate on temperature.

 

 

[I] = 1.11 x 10'§M; [11]

tl/Z = 20.3 min.

1.22 x 10‘31; 0.0 M 1101; Temperature 250;

 

0.26 1331 10525
5.02 1015 7811
10.10 5121 6132
15.02 6261 5595
25.02 6907 1919
31.91 8935 2921
a) 11856 -—

 

 

58

Dependence of rate on temperature. (Cont.)

 

 

t (min.) X (c/6 min.) Xm - X

 

[I] = 1.11 x 10'33; [II] = 1.22 x 10'33; 0.6 g 1101; Temperature 250;
t1/‘2 = 11.5 min.

 

0.17 2002 13617
0.77 2183 13136
1.56 2775 12811
3.02 3262 12357
5.03 1638 10981
7.52 5281 10338
10.53 7310 8309
13.97 6967 8652
00 15619 ..

 

 

[I] = 1.11 x 10‘31; [11] = 1.22 x 10‘31; 0.2 y 1101; Temperature 250;
tl/Z = 6.77 min. .

 

0.10 1623 11352
0.11 2356 13619
0.81 3057 12918
1.32 3702 12273
2.02 1159 11516
3.02 5802 10173
1.02 6118 9527
5.03 7250 8725
7.00 9161 6811
00 15975 ‘-

 

[1] = 1.11 x 10’?!; [11] = 1.22 x 10'31; 0.2 y 1101; Temperature 12.70;
t1/2 = 10.3 min.

 

0.15 1168 13796
0.62 2181 12780
1.18 3091 12170
2.01 3351 11913
3.01 3991 11270
5.17 5650 9611
7.09 5631 9633
10.02 8631 6631
13.72 9656 5608
00 15261 --

 

59

Dependence of rate on temperature. (Cont.)

 

 

t (min.) X (0/6 min.) Xm -X

 

[1] = 1.11 x 10'§y; [II] = 1.22 x 10'?fl; 0.05 g 1101; Temperature 12.70;
tl/Z = 37.0 min.

 

1.01 1862 13700
2.55 2978 12581
5.05 1108 11151
8.09 1101 11158
12.85 5086 10176
17.05 5219 10313
23.08 6218 9311
30.07 7726 7836
12.83 10137 5125
60.07 11151 1111
on 15562 --

 

[I] = 1.11 x 10‘31; [II]

1.22 x 10'?m; 0.0 g 1101; Temperature 12.70;

 

tl/2 = 60 min.

2.01 3119 10777
10.01 5991 8205
15.05 6808 7388
20.01 7610 6556
30.57 8003 6193
§§.0g 89E? 5259

.0 91 3 O 3
co 111196 --

 

[1] = 1.11 x 10'81; [11]
t1/2 = 58 min.

1.22 x.10'§y; 0.0 M 1101; Temperature 12.70;

 

2.70 1130 10785
7.00 6736 8179
11.70 8753 6162
25. 12 9786 5129
35.10 10068 1817
52.20 11191 3121
66.10 11765 3150

00 11915 __

 

 

6O

Dependence of rate on specific cation.’

 

 

[I] = 1.33 x 10'?g; [11] = 1.79 x 10'31; (X+ salts) 0.0 1 KCl; Temper—

ature 0°; tl/2 = 5.0 min.

 

 

t hum) X(q@1MnJ Xm-X
0.15 1216 12251
0.15 5693 10807
1.30 7251 9219
3.00 8521 7979
5.01 10771 5726

00 16500 --

 

[I] = 1.11 x 10fég; [II]
tl/Z = 2.2 min.

1.22 x lO-éfl; 0.10 3 KCl; Temperature 0°;

 

0.13
0.58
1.11
2.02
3.03
CD

2278
3653
’5181
7229
7153
13570

11292
9917
8389
6311
6117

 

[I] = 1.11 x 107§fl3 [II]
tl/2 = 0.73 min.

= 1,22 x 10-7M; 0.30 M KCl; Temperature 0°;

 

0.08
0.09
0.17
0.29
0.11
0.61
0.80
1.11
1.12
CD

3015
3137
3785
5368
6053
7278
8177
10565
10127
11083

11068
10616
10298
8715
8029
6805
5606
3518
3656

 

 

61

Dependence of rate on dielectric constant.

 

 

[1] = 1.11 x 10'31; [11] = 1.22 x 10'81; 0.6 M 1101; Temperature 00;
0 % Dioxane; tl/2 = 8.0 min.

 

 

t (min.) X (c/6 min.) Xa, - X
0.12 1571 10926
1.50 3279 9218
3.01 3999 8198
1.51 5155 7312
6.01 5833 6661
8.03 7130 5367

10.00 7838 1659

12.02 7981 1516

11.01 9307 3190

(D 12197 ——

 

 

[I] = 1.11 x 10-3M; [II] = 1.22 x 10_§M; 0.6 M LiCl; Temperature 0°;
1 % Dioxane; tl/Z = 9.6 min.

 

0.13 1161 10937
2.02 3858 8510
1.01 1153 7915
6.03 5196 7202
8.01 6115 5953
10.03 7116 1872
12.03 8191 1201
11.00 8126 3972
03 12398 __

 

[1] = 1.11 x 10'313 [11] = 1.22 x 10‘31; 0.6 1 1101; Temperature 00;
20% Dioxane; tl/2 = 11.0 min. '

 

0.10 1123 11206
2.02 2661 9965
1.02 3970 8659
6.02 1728 7901
8.03 5217 7112
10 02 5597 7032
12.01 6398 6231
11.00 6953 5676
18.22 7381 5215
22.10 8962 3667
a) 12629 --

 

62

Dependence of rate on dielectric constant. (Cont.)

 

 

t (min.) X (0/6 min.) XGD -X

 

[1] = 1.11 x 10‘31; [11] = 1.22 x 10'31; 0.6 y 1101; Temperature 00;
32% Dioxane; tl/2 = 19.1 min.

 

2.01 1885 7016
1.01 2630 6301
6.02 3120 5811
8.02 3587 5311
10.00 1122 1809
12.01 3917 1981
11.03 1131 1197
16.01 1965 3966
18.02 1888 1013
20.00 5561 3367
22.01 5653 3278
21.02 5621 3307
26.02 5537 3391
28.00 6609 2322
00 8931 --

 

 

APPENDIX II

Visible Region Absorption-Spectra of Anion I.[l2r tungstocobaltate (11)]
and Anion 11 [l2-tungstocobaltate (111)] in Water Solution. pH =’ 2.

abscissa: wavelength in millimicrons

ordinate: absorbance

63

61

 

 

 

 

 

 

 

 

 

 

0.5

0.4.

0.3

0.2 1

0.1

0'0.

800

700

600

500

400

 

 

APPENDIX III

A preliminary study of the electron paramagnetic resonance spectra

of the l2-tungstocobaltate (II) and l2-tungstocobaltate (III) ions.

65

66

An investigation of the electron exchange between the l2-tungsto-
cobaltate (II) and the l2-tungstocobaltate (III) anions has been described
in this thesis. The results of this investigation show that an electron
that is formally associated with the central or hetero atom of such a
complex, can readily pass through the "cage" of addenda atoms and associ-
ate with another ion. This suggests that there is extensive electron
delocalization in such a complex, and further elucidation of the electronic
structure would be of interest. One of the few methods for the direct
determination of electron density is from measurement of the hyperfine

coupling constants in electron paramagnetic resonance (EPR) experiments.

 

A well known example of such a study is the IrClS-2 ion, which has one
unpaired electron. (52) Hyperfine splittings were found both from the
Ir and from the C1 nuclei. From the magnitude of the splittings, it

was deduced that the unpaired electron spends about 70% of its time on
the Ir nucleus and about 5% on each chlorine. Ebsworth and Neil have
observed the EPR signal from [(N113)5C0-02-C0(N1-I3)5]+5 (53) in water solu—
tion, and from the hyperfine structure proved the equivalence of the

two cobalt atoms within the ion. Both the l2—tungstocobaltate (II) and
the lZ-tungstocobaltate (III) ions are paramagnetic, and have been shown
to have three and four unpaired electrons respectively. Further, the
graphs of magnetic susceptibility vs. temperature show no evidence of
spin aligning interactions down to liquid nitrogen temperatures. There-
fore, a preliminary study was undertaken to determine whether or not EPR
experiments might provide new data about the electronic structure of
these complexes. Experimental work was done on two instruments; a com—

mercial model, varian 1502, which could be operated down to liquid

67'
nitrogen temperature, and an apparatus constructed by the Michigan State
University physics department which is designed for use down to liquid
helium temperature. Both instruments employ X-band microwave frequency
sources. (pr-9000 megacycles).

Examination of the EPR spectra of the powdered potassium salts of
the ions gave no detectable absorption lines either at room temperature
or at liquid nitrogen temperature. Solutions of these salts, examined
in a specially designed quartz solution cell also gave negative results.
Single crystals of these salts showed no resonance absorptions down to
78°K.with the possible exception of the C0(II) compound which gave a
broad baseline depression. In order to lengthen the spin-lattice relax-
ation times, the search was carried on at approximately 10K, obtained
by pumping on liquid helium. At this temperature, the cobalt (II) com-
pound gave a resonance absorption line over 2000 gauss wide with no
fine structure. It was decided that spin—spin relaxation was the cause
of the line broadening, and therefOre a diamagnetic host-lattice was
sought for these compounds. The colorless potassium salt of l2—tungst0-
silicic acid proved to be suitable for this purpose. (The synthesis of
this acid has been described. (51) Large hexagonal prisms of this com-
pound could easily be grown from solutions containing 1% by weight of
either of the heteropbly cobalt compounds. Visual inspection of the
crystals so formed, indicated that they were homogeneous. From the
colors of these crystals it appeared that they contained approximately
as much of the cobalt compounds as the solution from which they were
grown. A single crystal of potassium l2—tungstosilicate containing 1%
of the l2—tungstocobaltate (II) ion gave a single resonance line at 10K.

This line appeared very close to the DPPH reference line, and its position

68
changed slightly as the static magnetic field was rotated, although its
intensity remained approximately the same. No fine structure was ob-
served. A similar crystal containing the 12-tungstocobaltate (III) ion
showed three resonance lines. The center line appeared at a field
strength of about 3000 gauss and did not shift as the static magnetic
field was rotated. The other two lines varied in distance from the
center line depending on the angle of the static field with respect to
the hexagonal crystal axis. Rotation through ninety degrees caused
these two lines,if initially coalesced with the center line, to diverge
from it to a maximum distance of less than 1000 gauss each way, and
to again coalesce with it. For no angle of the static field did any of
the lines show resolvable fine structure.

It seems unlikely that the spectra of these compounds will yield
sufficient detail to make further study of them worthwhile, at least
from the point of view of a chemist. The hoped for comparison of the
nuclear hyperfine splittings for the cobalt atoms in the two different
oxidation states is out of the question since no hyperfine structure was ob-
served. The reason for the failure to observe these splittings is not
clear, although it is probably related to the number of unpaired electrons

in these complexes.

 

 

 

in! . I
.Jw. .w.:’.’N.. !.I.

WEMK§TrY LIBRARY

111M

 

 

 

 

 

   

111117117[[1111771717S