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Proton contact shifts of some magnetically anomalous cobalt (11) complexes

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Proton contact shifts of some magnetically anomalous cobalt (11) complexes
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Vaughn, Joseph Benjamin, 1942-
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English
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vii, 73 leaves : ill. ; 28 cm.

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Anions ( jstor )
Electrons ( jstor )
Ions ( jstor )
Ligands ( jstor )
Magnetism ( jstor )
Metal ions ( jstor )
Nitrogen ( jstor )
Orbitals ( jstor )
Protons ( jstor )
Symmetry ( jstor )
Cobalt -- Magnetic properties ( lcsh )
City of Gainesville ( local )
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bibliography ( marcgt )
theses ( marcgt )
non-fiction ( marcgt )

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Thesis:
Thesis--University of Florida.
Bibliography:
Includes bibliographical references (leaves 70-72).
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Typescript.
General Note:
Vita.
Statement of Responsibility:
by Joseph Benjamin Vaughn, Jr.

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University of Florida
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Copyright [name of dissertation author]. Permission granted to the University of Florida to digitize, archive and distribute this item for non-profit research and educational purposes. Any reuse of this item in excess of fair use or other copyright exemptions requires permission of the copyright holder.
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PROTON CONTACT SHIFTS OF SOME MAGNETICALLY
ANOMALOUS COBALT(II) COMPLEXES







By

JOSEPH BENJAMIN VAUGHN, JR.


A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL
OF THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY





UNIVERSITY OF FLORIDA


1977
















TABLE OF CONTENTS

Page

LIST OF TABLES iii

LIST OF FIGURES iv

ABSTRACT v

INTRODUCTION 1

EXPERIMENTAL 14
Preparations 14
Elemental Analysis 24
Mass Spectrum Analysis 24
Spectrometers 24

RESULTS 26

DISCUSSION 34

APPENDIX 64

BIBLIOGRAPHY 70

BIOGRAPHICAL SKETCH 73















LIST OF TABLES


Table Page

I Selected Infrared Absorptions and
Assignments for the Co(II) and Fe(II)
Compounds Studied .27

II Magnetic Moments and g Values for the
Co(II) and Fe(II) Compounds Studied 30

III Observed and Relative Isotropic Shifts
and Half-Height Peak Widths for Cobalt(II)
Compounds 32

IV INDO Calculated and Experimental Proton
Spin Densities for [Co(GdH)-3+] 39

V INDO Calculated and Experimental Proton
Spin Densities for Cationic and Anionic
Radicals 42

VI Selected INDO Molecular Orbital Symmetries
and Energies of Ligand Radicals 53

VII Proton Chemical Shifts of the Substituted
2-Pyridinalphenylimine Ligands 65

VIII Proton Chemical Shifts of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II)
Complexes .66

IX Effective Magnetic Moments of the Sub-
stituted Tris(2-pyridinalphenylimine)
Cobalt(II) Compounds 68















LIST OF FIGURES

Figure Page

1. Ligands whose cobalt(II) complexes were
successfully prepared 10

2. Ligands whose cobalt(II) complexes could
not be prepared 11

3. Possible spin transfer mechanisms 37

4. Possible Oh--D correlation diagrams and
electron confi rations for a d7 ion under
D3 symmetry 50

5. Representation of ligand T MO and metal ion d
orbital parity upon rotation about the C2 axes. 52

6. Representation of spin polarization effects
in a fragment of a [M(GdH)r+] complex 56

7. Possible Oh--D2d correlation diagrams and
electron configurations for a d7 ion under
D2d symmetry 60











Abstract of Dissertation Presented to the
Graduate Council of the University of Florida
in Partial Fulfillment of the Requirements
for the Degree of Doctor of Philosophy



PROTON CONTACT SHIFTS OF SOME
MAGNETICALLY ANOMALOUS COBALT(II) COMPLEXES

By
Joseph Benjamin Vaughn, Jr.

August 1977
Chairman: R. Carl Stoufer
Major Department: Chemistry

A series of cobalt(II) complexes known to lie near the
E( 2G)-- T1( F) cross-over point was prepared. The first

reported proton contact shift investigation of such magnetic-

ally anomalous cobalt(II) complexes was carried out in order

to obtain information about the metal--ligand bonding proper-

ties. These complexes contain ligands with the a-diimine

linkage (glyoxaldihydrazone, GdH; biacetyldihydrazone, BdH;

biacetylbismethylimine, BMI; 2,6-pyridinediacetyldihydrazone,

2,6-PdAdH), which provides relatively large ligand field

splitting. The symmetry of the complexes containing these

symmetric, bidentate ligands is D3 and that of the terdentate

ligand complex is D2d. Under D3 symmetry 2Eg(Oh) transforms

as 2E(D3) and 4Tl(Oh) transforms as 4A(D3) and E(D3). Under
2 2 2
D2d symmetry Eg(Oh) transforms as Al(D2d) and 2Bl(D2d) and
2 4 4 4
Bl(D2d) and Tl(Oh) transforms as B2(D2d) and E(D2d).

Because these terms are isotropic, the dipolar shift








contribution to the contact shifts is neglected. The magni-

tudes of these shifts (100 ppm and less) are typical of those

observed for other complexes.

Spin densities were calculated using a modified INDO

program. The approach commonly used in such calculations is

to treat a single ligand as a radical cation or anion, using

only the "T" symmetry atomic orbitals in some instances, pre-

sumably due to the limited capacity of the available programs.

In contrast, the program used for these calculations is capable

of handling the complete valence basis set (ninety-nine

atomic orbitals) for [Co(GdH)3 ], the simplest complex in

the series. That is, metal penultimate d and ultimate s and

p orbitals and all the ligand valence orbitals were included.

But, meaningful results for even this simple complex were

not obtained, presumably because the number of unpaired spins

(three for high-spin complexes and one for low-spin complexes)

is sufficiently small that the spin densities calculated over

ninety-nine basis orbitals represent only "noise." Thus it

is concluded that such calculations do not yet have the degree

of sophistication necessary to provide meaningful results
++
even for such a simple complex as 1Co(GdH)3 ]. Consequently

each ligand was treated both as a radical anion and a radical

cation, as is commonly done, including the complete valence

basis set. Evaluation of these INDO calculated and experi-

mentally determined spin densities leads to the following con-

clusions.

(1) Because the best fit was obtained for the GdH+1








BdH+1 and BdH1, BMI- and no fit was obtained for 2,6-

PdAdH, it might be concluded that the spin transfer mechanism

is:

(a) L M in the complex [Co(GdH) ].

(b) L M and/or M L in the complex [Co(BdH)4].

(c) M L in the complex [Co(BMI)34.

(d) Undetermined in the complex [Co(2,6-PdAdH)214].
However, such conclusions are somewhat tenuous because of the

possibility of spin polarization, which adds further ambigu-

ity by placing opposing spins simultaneously at the same

position in the ligand.

(2) More than one molecular orbital is involved in the

transfer mechanism in each complex, making definitive conclu-

sions somewhat more difficult. Comparison of the INDO molec-

ular orbital calculations and the experimental data provides

evidence that spin transfer onto the hydrazone amine protons

occurs via a a molecular orbital with an opposing u molecular

orbital contribution, in keeping with the usual assertion

that d7--p7 back bonding is involved in these complexes.

(3) There is increasing lack of agreement within the

series of ligands, between INDO calculated spin densities and

the experimentally determined spin densities,which corresponds

to increasing contribution of the low-spin configuration to

the ground state of the complex. This trend is taken as an

indication that there is a contribution to the contact shifts

attributable to the low-spin configuration, the first such

contribution reported for an octahedrally coordinated d ion.


vii















INTRODUCTION


During the past century or so a great deal of research

has been directed toward the investigation of transition

metal coordination compounds. A significant portion of this

work has involved octahedral cobalt(II) compounds, the ma-

jority of which are high-spin, Veff = 4.7 5.2 BM, having

a 4Tg( F) ground term. There are relatively few octa-
hedral cobalt(II) coordination compounds which are low-spin,

=eff = 1.8 2.2 BM, having a 2E (2G) ground term. Since
g
1956 (40), there have been reported an increasing number of

magnetically anomalous cobalt(II) coordination compounds,

i.e., compounds which exhibit both room temperature magnetic

moments intermediate between high-spin and low-spin moments

and anomalous Curie-Weiss behavior.

Variable temperature magnetic studies led Stoufer and

coworkers (41) to postulate that these anomalous compounds

lie near the "cross-over" point and that Boltzmann distri-

bution over high-spin and low-spin states is attained, i.e.,

AE(2E --4 T1g) kT. They noted that the majority of these

compounds contain ligands, with an a-diimine linkage

(-N=C-C=N-), which provide relatively large ligand field

splitting due in part to their q-acid capabilities. Research

interest in these compounds still exists as evidenced by








the recent work of Simmons and Wilson (39), who investigated

solution and solid state magnetic and spin lifetimes of some

six-coordinate bis(N-R-2,6-pyridinedicarboxaldimiine) cobalt(II)

complexes. For the R=NH(CH3) derivative, they report solution

rates for E(s) 4T(hs) of kI = 3.2 x 106 s-1 and k- =
6 -1 k-1 2
9.1 x 106 s- with corresponding spin state lifetimes of T(2E)

= 3.1 x 10-7 s and ( T) = 1.1 x 10-7 s.

Information about bonding properties in paramagnetic

transition metal complexes as inferred from unpaired spin

distribution is traditionally obtained by electron spin

resonance techniques, where the isotropic couplings are mani-

fest as nuclear multiple structure. The sign of the cou-

pling constants cannot be determined directly from the ESR

experiment. However, the sign is obtained directly from

the NMR contact shift, and thus the sign of the spin density

is gotten directly. Because it is the sign of the spin den-

sity which gives the most useful information about bonding,

the major advantage is with the NMR technique. Furthermore,

the spectral complexity is greater in the ESR experiment.

That is, for a system with n coupling constants, ESR spectra

will exhibit 2n lines; whereas NMR spectra will exhibit only

n lines. However, the two techniques are more complementary

than competitive, since it is rare that both can be usefully

employed for the same system. This is because of the re-

quirements of very short (! 10-1 s) electron-spin relaxation

time, Tle, for observing NMR spectra and long Tle's (> 10-9 s)

for observing ESR spectra.









In their review article, Lewis and Morgan (26) make

a general distinction between species containing one un-

paired electron (s = ) and those with more than one (s > 1).

For S = systems, electron-spin relaxation is very ineffi-

cient, yielding well-resolved ESR spectra. Exceptions can

occur in systems possessing low-lying excited states, which

reduce Tle considerably. For systems with S > 1 the presence

of orbital degeneracy in the ground state, low-lying excited

states and/or modulation in solution of zero field splitting

can produce rapid electron-spin relaxation and therefore

well-resolved NMR spectra.

Because Tle is typically long for complexes in which

S = it has been predicted (42) that low-spin, 2E octa-

hedral cobalt(II) complexes will not yield observable NMR

contact shift spectra. And to date no NMR data have been

reported for any low-spin, octahedral, d7 ion (42). This is

due, in part, to the fact that there are no proton-containing,

low-spin, octahedral cobalt(II) complexes known. It may be

that the Eg species present in anomalous cobalt(II) com-

plexes investigated in this work will have Tie shortened

sufficiently to observe NMR spectra by the presence of the

low-lying T1g excited state near the cross-over point videe

infra).

Nuclear magnetic resonance spectra of paramagnetic

transition metal complexes typically exhibit very large

(200 ppm and larger for protons) contact shifts. Contact

shifts are measured relative to a diamagnetic standard,









usually the free ligand or, better yet, a homologous dia-

magnetic metal complex. And, the shifts may be related to

the transfer of spin either to or from the ligand. The

ligand is usually a closed shell diamagnetic molecule, where-

as the metal ion may possess some unpaired electrons in in-

completely occupied d or f orbitals. If any of the metal

orbitals containing the unpaired electron(s) is involved in

M-L bonding, the unpaired electron(s) may be transferred to

a new molecular orbital (MO) of the complex. Consequently

the unpaired spin may be partially delocalized onto the

ligand. McConnell and Chestnut (28) showed that for ligands

possessing nuclei with I # 0, this delocalized spin may

interact with these nuclei via the Fermi contact interaction

(8). For a complex possessing m spins, the Hamiltonian for

the Fermi interaction is given by R = A S"mI. A is the

hyperfine coupling constant for the contact interaction

between the electron spin, S, and the nuclear spin, I, and

is given by (36)


A = e gN 4e N -1
where p(r) is the spin density operator at the nucleus, 4 is

the electronic wave function for the molecule, is the

time averaged value of the z component of the electron spin,

and the other symbols have their usual meanings. Thus Fermi

coupling constants arise only from spin in s orbitals, and

A is a linear function of that spin density. The variation

in coupling constants for the various nuclei may lead to









the determination of the distribution of unpaired spin density

in the ligand, and thus to the characterization of the MO

containing the unpaired spin. Consequently a better under-

standing of electronic structure in transition metal complexes

may be achieved.

Both the sign and magnitude of A can be determined from

contact shift spectra. A is related to the contact shift

Av by (10,28)

2 2
Av Ag S(S + 1) (2)
AV e e (2)
v gN N 3 kT

if the Curie law is obeyed. It is clear from equations (1)

and (2) that a downfield shift in resonance position corre-

sponds to a positive value for A and therefore to positive

spin on the corresponding nucleus.

La Mar has outlined the three general areas of considera-

tion for the elucidation of the observed coupling constants

in terms of the M-L bonding: "(i) The relationship between

the observed coupling constant for a given nucleus and the

probability (density) of finding the unpaired spin in an

atomic orbital at the nucleus of interest and/or at a neigh-

boring nucleus, particularly as to how the coupling constant

reflects the type (a or 7) of ligand molecular orbital con-

taining the spin; (ii) the evaluation of the mechanisms by which

the unpaired spin on the metal is transferred to the ligand;

and (iii) the qualitative and, where possible, the quantitative

relationships between the fraction of total spin transferred

and the M-L covalency"(23:87).









There are several possible problems which may introduce

uncertainties into the interpretation of contact shift data.

Three major ones are: (i) Separation of the dipolar or

pseudocontact shifts from the measured isotropic shifts. These

dipolar shifts occur only in complexes which are magnetically

anisotropic, and are produced by the dipolar interaction

between the electron's magnetic moment, W, and the nuclear

magnetic moment h yN I. The dipolar shift averages to zero

by tumbling in solution unless the electronic magnetic moment

is anisotropic. It can, under appropriate conditions, be

separated from the isotropic shift by (36)


Adip (3 cos 2 1
--- (g3 g1) (3)
r

where D is a positive quantity whose value depends on the

temperature and various relaxation times (17,21,22,29) and the

other symbols have their usual meaning. Due to uncertainties

in D, most published calculations of absolute dipolar shifts

are open to question (7). However, if the g value is iso-

tropic or nearly so (gll gl = 0) dipolar shifts can be neg-

lected. (ii) Electron correlation effects can cause spin

density to be induced on L in the M-L fragment, even if the

metal orbital containing the spin is symmetry restricted

from interacting with any orbital on L. These effects can be

accounted for by use of an appropriate quantum mechanical

model for calculated spin densities. (iii) The possibility

of spin-polarization effects can also cause uncertainties in

determining the mechanism of spin delocalization. Again the









appropriate choice of quantum mechanical model can help over-

come this problem.

The mechanism by which the unpaired spin is delocalized

onto the ligand is related to the symmetry of the MO containing

it. These MO's are predominantly metal d orbitals with smaller

contributions from ligand orbitals which may be filled i or-

titals, empty I orbitals, a orbitals, or a mixture of all

three. Symmetry considerations sometimes restrict the possi-

bilities, in complexes of high symmetry for instance. But -

in general there is no easy method available for sorting out

the various contributions. Quantitative evaluation of M-L

covalency is no simple task, and the present development of

the evaluation of spin-delocalization mechanisms does not

permit a definitive presentation (23). "Thus the theoretical

complexities of relating transferred spin density by any

single mechanism to covalency, plus the practical problems

of multiple delocalization mechanisms, suggest that caution

should be exercised in any quantitative interpretations"

(23:120).

The ideal procedure for theoretical analysis would be

a complete, many-electron calculation on the entire complex

with a concomitant breakdown of the isotropic shifts into

their component parts. This would be a formidable undertaking

even for a simple complex. The nonideal approach commonly

employed is to treat a single ligand molecule as a free radical

anion or free radical cation depending on the delocalization

mechanism involved. The spin densities are then calculated









from the MO's containing the unpaired electrons and a corre-

lation is attempted between the calculated spin densities

and the relative spin densities gotten from the observed

contact shifts videe infra). In some cases, experimental

results satisfactorily fit calculated spin densities assuming

only one ligand MO is involved in the delocalization. In

other cases, it appears more than one ligand MO is involved.

In any case, this extreme approximation leaves much to be

desired regarding any quantitative conclusions. One obvious

objection to this approach is that the metal orbital contri-

butions are excluded in forming the MO's. The metal is con-

sidered to be simply an electron source or sink for the

ligand. This obviously assumes the relative energies of the

ligand MO's are unaltered by interaction with the metal.

Certainly this situation is seldom achieved.

There are several theoretical methods available for

treating spin-density calculations in metal complexes. In

this work, the INDO method was chosen because it was readily

available and it is more realistic than the others for the

following reasons (23,34,36):

(i) It is an all valence electron method.

(ii) It is capable of calculating negative spin densities

arising from electron correlation since it uses unrestricted

wave functions. The extended Hickel method, for example, can

calculate only zero or positive densities.

(iii) It accounts for spin polarization by including one-

center exchange integrals.








On the negative side, INDO calculations have "black-box"

aspects which tend to hinder a qualitative understanding of

the resultant calculations.

To date no calculations reported have included the metal

ion orbitals and all the ligand valence orbitals in the basis

set. This is because the present theoretical development is

simply not capable of handling such a large basis set. However,

one complex investigated in this work [Co(GdH) ] does have

a sufficiently small basis set (ninety-nine atomic orbitals)

to include the metal ion orbitals and all the ligand orbitals.

The purpose of the present work is to use NMR proton

contact shifts to investigate the metal-ligand (M-L) bonding

properties of a series of anomalous cobalt(II) compounds.

The bis(2,6-PdAdH) cobalt(II) complex and tris cobalt(II)

complexes of the other ligands in Figure 1 were prepared for

this purpose. Ligands with X = -CN and -NO2 were also pre-

pared, however, they did not coordinate to the metal ion. The

"tris" cobalt(II) complexes of GdH, BdH, BMI and the bis-

(2,6-PdAdH) cobalt(II) complex have been prepared previously

and are known to lie near the cross-over point with a E
g
ground term (5,10,37,41,44). The 2-pyridinalphenylimine

complexes were prepared with the expectation that increasing

the 7-withdrawing capability of the substituent would give a

homologous series of complexes which would exhibit magnetic

moments over the range of high-spin and low-spin values,

the ultimate goal for this series being to relate contact

shift changes to the Hammett acidity function. To accomplish

this goal the ligands in Figure 2 were prepared or their prep-

aration attempted.











H\ H
C-C
H2N-N N- NH2

Glyoxaldihydrazone
(GdH)


H3C / CH3
C- C
H C-N \N-CH3

Biacetylbismethylimine
(BMI)










H

SN-NH2

2-Pyridinalhydrazone
(2-PAH)


H 3C C
HNC
H21'uL-Nd


CH3
-C N
N NH2


Biacetyldihydrazone
(BdH)


C
H2N-N'


CH3
CN_ NH2


2,6-Pyridinediacetyl-
dihydrazone
(2,6-PdAdH)






X = H, CH3, NH2, C1

L '- -1


CN ) X

2-Pyridinalphenylimine
(2-PPI)


Figure 1. Ligands whose cobalt(II) complexes were success-
fully prepared





















H
\r /
4:-C\


H3 C-N


N- CH3


H3C\ C -C/ CH3

X N N X


X = H, NO2


X = H, NO2


R, R
SC-C H
X- N-N N N X


R = H, CH3; X = H, NO2





Figure 2. Ligands whose cobalt(II) complexes could not
be prepared

















H
N N

N N


HX
cX


X = H, OH, OCH3,


NO2


CH3


X = H, CH3, NO2











H3C' CH3
N C
X-- N N x
I i
H H


X = H, Br, CN









The rationale underlying this choice of ligands was to

prepare a homologous series of ligands (containing the a-di-

imine linkage) whose cobalt(II) complexes should lie near

the cross-over point. The substituents, X, have known o-and

n-acid effects. It is known that the r-acid character of

the ligand is important in the formation of the octahedral,

low-spin cobalt(II) complexes. Therefore, it was expected

that a series of ligands with increasing 7-acid effects would

produce a homologous series of complexes some of whose

members have a 4Tg ground term and some of whose members

have a 2E ground term (i.e., span both high-spin and low-

spin sides of the cross-over point). And, it was expected

that the resultant ri-acid capability could be related to the

change in proton contact shifts.















EXPERIMENTAL


In the preparations described below, melting points of

previously prepared compounds are all within experimental

error of those reported. The yields for all the compounds

were rather poor, and varied considerably upon repeating the

preparations. These poor yields were a result of washing or

recrystallizing the soluble product many times in order to

obtain a pure product. Consequently yields are not reported.

For several compounds, mass spectra parent peaks and their

relative intensities are reported in lieu of elemental analy-

sis.



Preparations



Biacetyldihydrazone, BdH. BdH was prepared by the method

of Fisher (10). M. p., 1590C.

Anal.: Calcd. for C4H10N4: C, 42.11; H, 8.77; N, 49.12.

Found: C, 42.23; H, 8.91; N, 49.03.



Glyoxaldihydrazone, GdH. GdH was prepared by the method

of Fisher (9). However, better yields were obtained using

only 40 ml of hydrazine rather than 75 ml. M. p., 990C.

Anal.: Calcd. for C2H6N4: C, 27.91; H, 6.98; N, 65.12.

Found: C, 27.78; H, 7.11; N, 64.85.









2,6-Pyridinediacetyldihydrazone, 2,6-PdAdH. A solution

of 50 ml of methanol and 4.0 ml (0.12 mole) of hydrazine was

prepared and placed in an ice bath. Another solution of

8.1380 g (0.0498 mole) of 2,6-diacetylpyridine dissolved in

50 ml hot methanol was added dropwise with stirring. The

white powdery product was washed with cold methanol and dried

in vacuo over P4 010 M. p., 1790C.

Anal.: Calcd. for C9H13N5: C, 56.54; H, 6.81; N, 36.65.

Found: C, 56.16; H, 6.98; N, 36.47.



2-Pyridinalhydrazone, 2-PAH. To 3.0 ml (0.09 mole) of

hydrazine in an ice bath, 7.2 ml (0.06 mole) of 2-pyridine-

caroxaldehyde (2-PCA) was added dropwise with stirring. The

golden yellow solution was "dried" in vacuo over P4010 leav-

ing a golden yellow oil. B. p., 1810C.

Anal.: Calcd. for C6H7N3: Mol. wt., 121 d.

Found: Mol. wt., 121 d (46%).



2-Pyridinalphenylimine, 2-PPI. To a solution of 9.1 ml

(0.1 mole) aniline in 10 ml hexane was added dropwise with

stirring 10.7 ml (0.1 mole) of 2-PCA. The bright yellow

powdery precipitate isolated was recrystallized from hot

hexane and dried in vacuo over P4010. M. p., 400C.

Anal.: Calcd. for C12H10N2: Mol. wt., 182 d.

Found: Mol. wt., 182 d (14%).









2-Pyridinal-p-erethylphenylimine, 2-PMI. A solution of

10.7155 g (0.1 mole) of p-aminotoluene dissolved in 50 ml

hexane was prepared by heating on a steam bath, and 10.7 ml

(0.1 mole) of 2-PCA was added slowly with stirring. This

yellow solution was allowed to cool to ambient temperature.

The yellow needle-like crystals which precipitated were

recrystallized from hot hexane and dried in vacuo over P4010

M. p., 550C.

Anal.: Calcd. for C13H12N2: Mol. wt., 196 d.

Found: Mol. wt., 196 d (69%).



2-Pyridinal-p-aminophenylimine, 2-PAI. A solution of

10.8982 g (0.1 mole) of p-phenylenediamine dissolved in

100 ml hot deionized water was prepared. One gram of sodium

hydrosulfite was added to this dark brown solution, and some

black particles were removed by filtration. To this beige

solution, 10.7 ml (0.1 mole) of 2-PCA was added dropwise

with vigorous stirring. The gold powdery precipitate prod-

uced was washed with hot deionized water and dried in vacuo

over P4010. M. p., 1540C.

Anal.: Calcd. for C12H11N2: Mol. wt., 197 d.

Found: Mol. wt., 197 d (100%).



2-Pyridinal-p-chlorophenylimine, 2-PCI. With gentle

heating and stirring, 10.7 ml (0.1 mole) of 2-PCA was added

to 12.7613 g (0.1 mole) of p-chloroaniline, and the solution

was allowed to cool to ambient temperature. The resultant








pale yellow powder was recrystallized from hot hexane and

dried in vacuo over P 010. M. p., 650C.

Anal.: Calcd. for C12H9N2Cl: Mol. wt., 217 d.

Found: Mol. wt., 217 d (41%).



2-Pyridinal-p-cyanophenylimine, 2-PCPI. While heating

7.0 ml (0.06 mole) of 2-PCA on a steam bath, 6.0353 g (0.05

mole) of 4-aminobenzonitrile was added slowly. The resultant

yellow solution was heated for one hour, and then extracted

six times with 100 ml boiling hexane. The hexane extracts

were evaporated and the resultant yellow powdery product

was dried in vacuo over P4010. M. p., 1090C.

Anal.: Calcd. for C12H9N3: Mol. wt., 207 d.

Found: Mol. wt., 207 d (51%).



2-Pyridinal-p-nitrophenylimine, 2-PNI. While heating a

slurry of hot hexane and 13.8017 g (0.1 mole) of p-nitro-

aniline on a steam bath, 10.7 ml (0.1 mole) of 2-PCA was

added. The amount of solid appeared to increase and became

more flocculent. This mixture was heated subsequently for

one hour, washed four times with hexane, and dried in vacuo

over P4010. M. p., 122 C.

Anal.: Calcd. for C12H9N302: Mol. wt., 227 d.

Found: Mol. wt., 227 d (9%).



Tris(biacetyldihydrazone)cobalt(II) tetrafluoroborate,

[Co(BdH)3](BF4)9. To a solution of 1.8247 g (0.01 mole) of









Co(N03)2 in 15 ml deionized water, 3.4175 g (0.03 mole) of

BdH was added with stirring. The resultant deep brown solu-

tion was filtered and 20 g NaBF4 was added with stirring.

The light brown powder that separated was filtered, washed

with deionized water, and dried in vacuo over P4010. M. p.,

2460C dec.

Anal.: Calcd. for [Co(C4H o N4)3] (BF4)2: C, 25.06; H, 5.22;

N, 29.25.

Found: C, 25.23; H, 5.29; N, 29.36.


Tris(biacetyldihydrazone)iron(II) iodide dihyrate,

[Fe(BdH)3]I2-2H90. Excess BdH was added with stirring to a

solution of 3.9219 g (0.01 mole) of Fe(NH4)2(SO4)2.6H20 in

15 ml deoxygenated deionized water. This deep red-brown

solution was filtered, and 45 g NaI was added to the 45 ml

of solution. The red-brown powdery product was washed with

deionized water and dried in vacuo over P4010 M. M. p.,

2050C dec.

Anal.: Calcd. for Fe(C4H10 N4)32-2H20: C, 20.91; H, 4.94;

N, 24.46.

Found: C, 20.61; H, 4.46; N, 24.20.


Tris(glyoxaldihydrazone)cobalt(II) iodide, [Co(GdH)3]T2.

All reagents, solvents and equipment were deoxygenated under

nitrogen for twenty-four hours, and the preparation was also

carried out under nitrogen. A solution of 3.6751 g (0.02

mole) of Co(NO3)2 in 10 ml deionized water was made, and a









white slurry of 5.1116 g (0.06 mole) of GdH in 15 ml deionized

water was added with stirring. The resultant brown slurry

was allowed to stir thirty seconds and then filtered. To

the brown solution, a solution of 25 g of NaI in 25 ml de-

ionized water was added with stirring. Thirty seconds later

a black powder was filtered off and dried in vacuo over P4010.

M. p., 1650C dec.

Anal.: Calcd. for Co(C2H6N4)312: C, 12.62; H, 3.15; N, 29.44.

Found: C, 12.29; H, 3.08; N, 28.71.



Tris(glyoxaldihydrazone)iron(II) iodide, [Fe(GdH)3]I2.

A solution of 1.3675 g (0.01 mole) of FeC12 in 25 ml deionized

deoxygenated water was made under nitrogen. To this solution

2.5817 g (0.03 mole) of GdH was added with stirring. The

resultant intense red-brown solution was filtered, and 25 g

of NaI was added with stirring. The red powdery product was

filtered, washed with deionized water, and dried in vacuo

over P4010. M. p., 2020C dec.

Anal.: Calcd. for Fe(C2H6N4)312: C, 12.68; H, 3.17; N, 29.60.

Found: C, 12.79; H, 3.22; N, 29.80.



Tris(biacetylbismethylimine)cobalt(II) iodide,

[Co(BMI)']I2. This complex was prepared following the method

of Clevenger (5) being extremely careful to deoxygenate all

equipment and reagents with nitrogen. M. p., 2200C dec.

Anal.: Calcd. for Co(C6H12N3)312: C, 33.30; H, 5.55; N, 12.95.

Found: C, 33.13; H, 5.56; N, 12.80.









Tris(biacetylbismethylimine)iron(II) tetrafluoroborate,

[Fe(BMI)_] (BF4)2. This compound was prepared by a method

analogous to that used for preparing Co(BMI)3I2 except that

FeCl2 and NaI were used. M. p., 2680C dec.

Anal.: Calcd. for Fe(C6H12N3)3(BF4)2: C, 38.28; H, 6.37;

N, 14.86.

Found: C, 38.32; H, 6.44; N, 14.91.


Bis(2,6-pyridinediacetyldihydrazone)cobalt(II) iodide,

[Co(2,6-PdAdH)2]I2. A solution of 5.3108 g (0.03 mole) of

Co(N03)2 in 50 ml of deionized water was prepared, and

12.1071 g (0.06 mole) of 2,6-PdAdH was added with stirring.

The resultant brown solution was filtered to remove excess

ligand, and 50 g of NaI was added with stirring. The black

powdery product was then removed by filtering, washed with

cold water, and dried in vacuo over P4010. M. p., 2700C.

Anal.: Calcd. for Co(C9H13N5)2 2: C, 31.09; H, 3.75;

N, 20.15.

Found: C, 31.25; H, 3.76; N, 20.23.


Bis(2,6-pyridinediacetyldihydrazone)iron(II) iodide di-

hydrate, [Fe(2,6-PdAdH)2]I 2H20. This compound was prepared

by a method analogous to that used for preparing

[Co(2,6-PdAdH)2]I2 except that Fe(NH4)2(S04)26H20 was used.

M. p., 3000C.

Anal.: Calcd. for Fe(C9H 3N5)212 2H20: C, 29.81; H, 4.13;

N, 19.36.

Found: C, 30.30; HI, 4.10; N, 19.46.









Tris(2-pyridinalhydrazone)cobalt(II) iodide,

[Co(2-PAH)3 I]. A solution of 3.0 ml (0.06 mole) of hydra-

zine in 10 ml of deionized water was prepared in an ice bath,

and 7.2 ml (0.06 mole) of 2-PCA was added dropwise with

stirring. To this straw yellow solution, 3.6592 g (0.02

mole) of Co(N03)2 dissolved in 10 ml deionized water was

added dropwise with stirring. The resultant deep brown

solution was then heated with stirring to 1000C, allowed to

cool to ambient temperature, and placed in a refrigerator

for 6 hours. A beaker containing the solution was then

scratched with a freshly broken glass rod and allowed to

sit 30 minutes. The resultant tan powder was washed with

cold deionized water and dried in vacuo over P4010. M. p.,

2200C.

Anal.: Calcd. for Co(C6 7N3)312: C, 31.96; H, 3.11; N, 18.65.

Found: C, 31.26; H, 3.11; N, 18.09.



Tris(2-pyridinalphenylimine)cobalt(II) tetrafluoroborate,

[Co(2-PPI)3](BF )2. A solution of 1.8175 g (0.01 mole) of

Co(N03)2 dissolved in 30 ml of 50% ethanol and 50% deionized

water was prepared. To this solution, 5.7518 g (0.03 mole)

of 2-PPI was added with stirring. After adding 100 ml de-

ionized water, this red-brown mixture was allowed to stir

for 24 hours. The aqueous layer was separated from a red oil,

and 6 g of NaBF4 was added with stirring to the aqueous mix-

ture. The resultant red-brown oily precipitate and aqueous

mixture was heated with stirring, and a beige powder was








produced. It was washed with deionized water and dried in

vacuo over P4010. M. p., 2800C dec.

Anal.: Calcd. for Co(C12H10N2)3(BF4)2: C, 55.49; H, 3.85;

N, 10.79.

Found: C, 55.33; H, 3.92; N, 10.88.


Tris(2-pyridinal-p-methylphenylimine)cobalt(II) tetra-

fluoroborate, [Co(2-PMI)3](BF,)2. This compound was prepared

by a method analogous to that used for [Co(2-PPI)3](BF4)2.

M. p., 2830C.

Anal.: Calcd. for Co(C3H2N2)3 (BF4)2: C, 57.03; H, 4.39;

N, 10.24.

Found: C, 56.76; H, 4.45; N, 10.33.


Tris(2-pyridinal-p-aminophenylimine)cobalt(II) tetra-

fluoroborate, [Co(2-PAI)3](BF,)2. All reagents, solvents,

and equipment were deoxygenated and all procedures were carried

out under nitrogen in order to obtain this compound. A solu-

tion of 1.8271 g (0.01 mole) of Co(N03)2 in 25 ml deionized

water was prepared, and 6.1028 g (0.03 mole) of 2-PAI was

added and the mixture was allowed to stir for 15 minutes.

This red-brown mixture was then filtered, and 5 g of NaBF4

was added with stirring to the aqueous solution. The result-

ant red-orange powdery solid was washed with deionized water

and dried in vacuo over P4010. M. p., 1700C dec.

Anal.: Calcd. for Co(C12H11N3)3(BF4)2: C, 52.46; H, 4.01;

N, 15.30.

Found: C, 52.18; H, 4.11; N, 15.22.









Tris(2-pyridinal-p-chlorophenylimine)cobalt(II) tetra-

fluoroborate, [Co(2-PCI)3](BF,)2. To a solution of 1.8294 g

(0.01 mole) of Co(N03)2 in 15 ml of deionized water, 6.5026 g

(0.03 mole) of 2-PCI was added and the mixture was allowed

to stir for 24 hours. The resultant orange powder was dis-

solved in 50 ml of hot ethanol, and 10 g of NaBF4 was added

with stirring. This red-brown solution was allowed to sit

at ambient temperature for 72 hours. A red-brown solid

consisting of needle-like crystals (soluble in dimethylsulf-

oxide, DMSO) and round globs (insoluble in DMSO) was obtained.

Enough DMSO was added to just dissolve the crystals, and

the round globs were removed by filtering. The orange DMSO

solution was warmed to about 750C, and deionized water added

until an orange powder just began to precipitate. This solu-

tion was placed in a refrigerator for 6 hours. The orange

product was dried in vacuo over P4010. M. p., 2450C.

Anal.: Calcd. for Co(C12H N2C1)3(BF4)2: C, 48.98; H, 3.07;

N, 9.53.

Found: C, 48.72; H, 3.16; N, 9.39.



The cobalt(II) complexes of 2-PCPI and 2-PNI could not

be prepared. Repeated attempts to prepare these complexes

in deionized water, ethanol, methanol, tetrahydrofuran,

carbon tetrachloride, chloroform, and dimethylsulfoxide met

with failure. The cobalt complexes of GdH, BdH, BMI, and

2,6-PdAdH are not stable in air. However, the remaining

cobalt complexes and the iron complexes are considerably

more stable.








Elemental Analysis



All elemental analyses were performed by Atlantic Micro-

lab, Inc., Atlanta, Georgia.



Mass Spectrum Analysis


Parent peaks for several ligands were obtained, in lieu

of elemental analysis, using a voltage of 70 eV, on an AEI

MS 30 mass spectrometer equipped with a DS 30 data system.



Spectrometers


A Varian Associates XL-100 NMR spectrometer was used to

obtain the NMR spectra.

Infrared spectra were obtained on a Perkin-Elmer 137

sodium chloride spectrophotometer, and were calibrated with

polystyrene. Pressed potassium bromide samples were used to

obtain infrared spectra.

Electron paramagnetic resonance spectra were obtained

at 80 K using a Varian Associates E-3 EPR spectrometer,

V-4502-14.

The magnetic susceptibilities were determined by the

Gouy method. The equipment has been described previously

(5). The magnetic field was calibrated using mercury(II)

tetrathiocyanatocobaltate(II) (25). The maximum field strength

at a current of 4 A was 6750 + 80 gauss.









Solution susceptibilities were determined using a

Varian Associates A-60 NMR spectrometer, and the method of

Lbliger and Scheffold (27). Spectral grade methanol was

used to prepare solutions of 2 x 10-5 mole/ml.

Theoretical spin density calculations were obtained from

a INDO/1 program modified to include calculations on molecules

containing atoms from the first transition series. The ori-

ginal INDO model is not changed from that given by Pople and

Beveridge (34). This modified program was written and devel-

oped by Michael Zerner and Alan Bacon at the Chemistry Depart-

ment, University of Guelph, Ontario.














RESULTS


The positions of the infrared absorptions and their

corresponding assignments for characterization of the compounds

prepared for this investigation are listed in Table I. These

assignments are based on the ranges given in Bassler and

Silverstein (2). In addition, similar v(C=N) assignments

were made by Fisher and Stoufer (11) for Co(GdH)3Br2, Co(BdH)312,

Fe(GdH)312, and Fe(BdH)312.

Average room temperature magnetic moments, veff, cal-

culated from susceptibility data, and g values calculated

from ESR data are reported in Table II. These are average

moments determined at field strengths of 5649 Hz and 6750 Hz

to insure there was no field strength dependence and thus

no ferromagnetic or antiferromagnetic interactions were

present. Pascals constants for diamagnetic corrections to

the measured susceptibilities were obtained from Lewis and

Wilkins (25) and from Selwood (38). It would have been

preferable to determine peff in solution for all the compounds,

however, only Co(2,6-PdAdH)212 was amenable to this technique,

since the other cobalt(II) compounds decompose too rapidly

in solution. Solution measurements for these would yield

Veff values which would be the average for each compound
and its decomposition products. As reported previously (5,10)










Table I
Selected Infrared Absorptions and Assignments for
the Co(II) and Fe(II) Compounds Studied


Compounds

Co(GdH)3I2


Co(BdH) 3 (BF4)2
















Co(2,6-PdAdH) 22


Absorption (cm- )


3200 s

3100s(sh)

3000 s

1590 m(sh)

1570 s

1410 m


3510

3420

3220

3000

1650

1610

1440

1375



3400

3300

3180

3100

3000

2850


S

s

s

s

m(sh)

m


Assignment

V (N-H)

as (N-H)

s (C-H)

)(C=N)

v(C=N)

6(C-H)


s (N-H)

as (N-H)
as (C-H)

v (C-H)

6(N-H)

v(C=N)

as (C-H)

6s(C-H)


s (N-H)

vas(N-H)

vas(ring C-H)

vs(ring C-H)

vas(methyl C-H)

v (methyl C-H)













Compounds

Co(2,6-PdAdH)2I2


Table I continued



Absorption (cm- )


1620

1580

1440

1375

800


Assignment

6(N-H)

v(C=N)

6 as(methyl C-H)

6s(methyl C-H)

6(ring C-H)


Co(BMI)3I2


2900

2850

1590

1420

1375


Fe(GdH)3I2b













Fe(BdH)312 2H20b


w(sh)

w

m

m

m


3225 s

3100 s

3000s

1580 s

1540 s

1416 w



3350 m

3250 m

3150 m

2900 w


vas(C-H)

v (C-H)

v(C=N)

6 as(C-H)

is(C-H)



vs(N-H)

as(N-H)

Vs(C-H)

v(C=N)

v(C=N)

as (C-H)



us(N-H)

Sa(N-H)
as

vas(C-H)
vs(C-H)
s













Compounds

Fe(BdH)312 2H20


Table I continued



Absorption (cm- )


1620 m

1590 m

1440 m

1375 m


Assignmenta

6 (N- H)

V(C=N)

as (C-H)
as(C-H)
6 (C-H)


Fe(2,6-PdAdH) 212
























Fe(BMI)3(BF4)2


3400

3300

3180

3100

3000

2850

1620

1590

1440

1375

800



2950

2850

1610

1430

1375


m

m

m

m

w(sh)

w

m

m

m

m

in



w

m

w

s

(s)


s (N-H)

v a(N-H)

vas(ring C-H)

us(ring C-H)

Sas(methyl C-H)

us(methyl C-H)

6 (N-H)

v(C=N)

as(methyl C-H)

6s (methyl C-H)

6(ring C-H)



vas(C-H)

us (C-H)

v(C=N)

6a(C-H)

6 (C-H)


aBased on ranges given in Bassler and Silverstein (2).
Similar v(C=N) assignments were made by Fisher and Stoufer (11).










Table II
Magnetic Moments and g Values for the
Co(II) and Fe(II) Compounds Studied


Compoundsa


Co(GdH)3I2

Co(BdH))312

Co(2,6-PdAdH) 212

Co(BMI)3T2

Fe(GdH) 12e

Fe(BdH)3I2 2H20e

Fe(2,6-PdAdH) 2 e

Fe(BMI)3 (BF4)2e


Jeff (300 K, in BM)

4.23


4.65b

2.28b,c

3.30d


0.00

0.61

0.57

0.80


g (80 K)


2.15

2.06

2.09

2.13


rascals constants for
susceptibilities were
and Selwood (38).

The reported value is


diamagnetic corrections to the measured
obtained from Lewis and Wilkins (25)


4.17 BM (40).


Determined by NMR in solution (27).

The reported value is 2.91 BM (41).


eNo EPR signal was observed.
No EPR signal was observed.







these anomalous moments for all the cobalt(II) compounds

span the range between low-spin, 1.7 BM, and high-spin, 5.2

BM, values. It is noteworthy, however, that the 2,6-PdAdH

complex has a Ueff sufficiently low, 2.28 BM, that it may be

considered to be a low-spin complex. The observed nonzero

moments for the "diamagnetic" iron(II) compounds are a result

of temperature independent paramagnetism.

The EPR spectra for these cobalt(II) compounds consisted

of a very broad signal centered near a g value of 4 attributable

to the slightly populated E ( T2g) excited term, and, superim-

posed on it, a relatively narrow signal centered upon a g value

in the range 2.06 2.15, very close to the spin-only value of

2.00229 and the isotropic value of 2 expected for the 2E level.

The iron(II) compounds exhibited no ESR signal as expected.

The EPR data presented above and the NMR data which follow

are correlated in the discussion videe infra). The isotropic

shifts, Av, for the cobalt(II) compounds are reported in

Table III together with the relative shifts and the corre-

sponding half-height widths. These relative isotropic shifts,

Avrel, are normalized with respect to the largest observed

shift, which in each case happens to be the proton-con-

taining moiety bound to the hydrazone or imine nitrogen. As

stated previously the GdH, BdH, and BMI complexes decompose

rather quickly in solution. Therefore, for each it was

necessary to make all instrumental adjustments and run a

trial spectrum on one sample and run the final spectrum

quickly on a freshly prepared sample. Within ten minutes

resonances attributable to the decomposition products

began to appear. Thus it was possible to make unambiguous










Table III
Observed and Relative Isotropic Shifts
and Half-Height Peak Widths for Cobalt(II) Compounds


Av (ppm)



2.29

- 103.66



2.72

86.69



10.14

42.90


vrel (ppm)a



- 0.022


1.00


0.031

1.00


0.236

1.00


Width (Hz)



60


500


30

150


60

150


2,6-PdAdH


C-H

N-H


8.97


- 61.91

- 12.09

- 6.06


- 0.144

1.00

0.195

0.098


110

750

60

80


aRelative isotropic shifts are normalized with respect
to the largest observed shift.


Ligand


GdH


C-H

N-H


BdH


C-H

N-H


BMI


C-C-H

N-C-H









assignments based on these observations, the relative peak

intensities, and the fact that all the protons bound to the

hydrazone or imine group would be expected to shift the same

direction. Furthermore, they exhibit the same relative

magnitude (i.e., largest shifts observed in each spectrum)

and all are the broadest peaks in the spectra, which broaden-

ing is due to their proximity to the nitrogen quadrupole.

There was no decomposition problem involved in the

assignments for 2,6-PdAdH. However, the HA and HB isotropic

shifts as assigned by relative intensities are reversed

compared to the expected shifts videe infra). Consequently,

the assignments for HA and HB are based on the following

syllogism. Because the electron-proton dipolar relaxation

mechanism often dominates the line widths in paramagnetic

complexes, the relative line widths for nonequivalent protons

will be determined primarily by their relative values of

R where R is the distance from the central metal ion to

the proton (4,24). Since RH > RH (see the structures in

the introduction), R6 < R6 and the line width for HA should
A B
be less than that for HB.















DISCUSSION


Theoretically it is predicted that proton contact shift

spectra for complexes with s = will not be observable

because Tle is usually too long (42). And as noted in the

introduction no spectra have been reported for low-spin,

octahedrally coordinated d7 ions. As inferred from the

introduction, the purpose of this investigation was, in part,

to determine if proton contact shift spectra are observable

for anomalous cobalt(II) complexes near the cross-over point,

and to determine if the spectra are attributable to the 2E (2G)
g
term. If so, this would mean Tle is sufficiently lowered by

low-lying excited states (42).

There are several approaches which should be considered

in arriving at an interpretation of the isotropic shifts of

these complexes. (i) The site symmetry of these complexes

may be taken to be octahedral. Accordingly, if the concen-

tration of the 4T( F) species is small, and if the 2E --4T1
tration of the species E


interconversion provides a mechanism whereby Tle is sufficient-

ly lowered, then a contact shift should arise from the 2E
g
term. In this case there would be no spin-orbit coupling
9
effects from "E and, because E and A terms are isotropic,
g
there would be no dipolar shift complications to consider (15).

(ii) Under octahedral symmetry the contact shift arises from







the 4T term. This is the origin of all reported cobalt(II)

contact shift spectra (42). In this case there are spin-orbit

coupling and dipolar shift complications. (iii) Both (i)

and (ii) contribute to the observed spectra. In this case,

if the 2E -- T interconversion time is large on the NMR time

scale, two superimposed spectra should result. This was not

observed for these complexes (Table III). (iv) If (i) and

(ii) contribute and the E -- T1 interconversion time is
g 1
small on the NMR time scale, time-averaged spectra should

result. In this case spin-orbit coupling (44) and dipolar

or pseudocontact shifts from the 4T1 term are involved. But,

approaches (i) (iv) are valid explanations only under octa-

hedral symmetry.

Strictly speaking six-coordinate complexes containing

symmetric bidentate ligands are inherently of D3 symmetry;

consequently approaches (i) (iv) are not valid considerations.

Under D3 symmetry 2E (Oh) 2E (D3) and 4Tl(Oh) 4A(D3) +

E(D3). Since all the term levels in D3 symmetry are isotropic

(A and E), the observed isotropic shifts should be due en-

tirely to contact interaction (15,29,43). Therefore pseudo-

contact contributions either should be absent or negligible.

The nonideal approach commonly used to interpret contact

shift data is to treat a single ligand molecule as a free

radical cation or anion depending on the delocalization mech-

anism involved. That is, the delocalization of unpaired spin

is considered to result from L M or M L charge transfer

(23,36). In the discussion which follows, (+) and (+)








represent an electron whose magnetic moment is aligned par-

allel and antiparallel, respectively, to the external magnetic

field.

Figure 3 illustrates the spin transfer mechanisms possible

(23). The ligands contained in these complexes are all dia-

magnetic and thus the ligand MO's, L' either are doubly

occupied or are unoccupied. L M charge transfer is

illustrated in (a) (c). If the metal ion has a singly

occupied d orbital of a symmetry (a) is the only transfer

possible. The spin is transferred-from a ligand a orbital

(OO) to the occupied metal do and paired with the spin in the
do orbital. Since the metal ion unpaired spin will be par-

allel (-) to the external field, this leaves net (+) spin

on the ligand. The metal ion can contain a singly occupied

d" orbital with a low-lying do orbital, (b) and (c). Un-

paired spin (+) and () can be transferred into the do from

the L. However, correlation favors transfer of (t) spin

(Hund's rule of maximum multiplicity) leaving net (+) spin

in eL (32). M L charge transfer is illustrated in Figure

3 (d) (f). For a diamagnetic ligand a vacant L is re-

quired. Vacant o MO's are too high in energy to be im-

portant, so only vacant 7T MO's need be considered. If the

metal d orbital is singly occupied, (d) is the only trans-

fer possible and net (t) spin is delocalized onto the ligand.

However, if the unpaired spin occupies a d orbital, and

the d orbitals are doubly occupied, (e) and (f) spin trans-

fer is possible. In this case, correlation favors transfer

of (+) spin from metal to ligand.



















d 7


(a)


(b)


, --4 d
- __d


(c)


Sd
..-++ 'a -


(f)


Figure 3. Possible spin transfer mechanisms


-- do


L


L


- L


-- L









A theoretical analysis of the [Co(GDH) ]spectrum was

attempted based on a complete, many-electron calculation for

the entire complex. Such a calculation was performed with

the complex considered to be (i) high-spin with sp2 and with

sp3 hybridized amine nitrogen orbitals and (ii) low-

spin with sp2 and with sp3 hybridized amine nitrogen orbitals.

These calculations were accomplished using a modified INDO

program. Since X-ray crystallography data were not available,

the bond lengths used in these calculations are the average

bond lengths given by March (30). This program is capable

of handling the minimum basis set (ninety-nine atomic orbitals)

of the complex, including metal penultimate d and ultimate s

and p orbitals and all of the ligand valence orbitals.

The results of the calculations together with the ex-

perimental results are listed in Table IV. An inspection of

the data will reveal that neither the absolute magnitudes

nor the signs of the theoretical spin densities agree with

the experimentally determined values and signs. Indeed, even

the relative magnitudes do not agree. That is, although the

hydrazone protons are found to have the largest shifts and

therefore the largest spin densities, the INDO calculations

predict the protons attached to the carbons to exhibit the

largest shift. One factor causing this disagreement may be

that these experimentally determined densities were calculated

using Equation (2) which is applicable to complexes having

normal Curie behavior, a condition which does not exist for

this complex videe infra). Another factor may be that,















0
O
0 -r
4-J 4-3
co
)3 -$ 0 -0 N
4> CM D C .N
rJ O 0 0 r4
C 4 0 0 0 0 (U *r2-




b 0 o





4-1 h
00O


L4-4 ,- ,0 Lf Cfr-
0 8 .i1

L > cM CM U
U c CN O Cq C 1 0
0 4J 0 0 0 C C0 aO
U &4 (U nJ
0 "- I I C0 C



















OO 0e nn .
Q o ) o I o N N o
C-4 CN I C N) 0>
*rol S Ir 0
>C 4.




0 C cn CU .C N
*0- > n O



-4 C D O C M 0 0C
co CL 1- r-0 C 0 I r- U



) uo ) P (

o0 m
CI C o no m j-I









P o 0 o 0
0 Lfl 0 iUi 0e














C- 0*i0 o -<4 10

*r- 0 -4 0CD

I-1 M Q1)
-40 CI)


















co x r ap


'-< -1 I I CM MO *i-
tO CO XD r
U-I 0 N


Q co d4
-^c U n )

O r-O O 001
*) C 4 CO CO c 4.

l0 00 C13 0 C
"a C- ) Q0 (










U) 10 t
'O o 0\








I -N CP4








although this modified INDO method is capable of handling this

basis set videe supra), the amount of unpaired spin density

calculated on the protons is sufficiently small (Table IV)

that it can only be considered to be "noise." This is due to

the fact that the tolerance in this program cannot be set

sufficiently low to insure convergence, and thus calculate

meaningful densities for a total of one unpaired electron

(low-spin) or a total of three unpaired electrons (high-spin)

in the complex over ninety-nine basis orbitals. Consequently,

no meaningful importance should be attached to these calculated

densities.

Since INDO calculations did not produce meaningful

results, the common (but nonideal) approach of considering

the ligand to be a radical cation or anion was used even

for this simple complex. However, in this work the entire

valence set of the ligand was used rather than just the 7

symmetry basis functions as is frequently done (see reference

36 and references therein). The metal orbitals are completely

neglected in forming the molecular orbitals for the calcula-

tions; that is, the metal ion has been considered to be

either an electron source or sink. In addition, it is general-

ly accepted that ligands containing an a-diimine linkage

serve not only as a donors but as 7-acids as well (d7--pr

back bonding with the metal ion). That is, there may be M L

and L M spin transfer simultaneously. Obviously it is an

oversimplification to consider the ligand to be either a

cation or an anion under such circumstances. Furthermore,







the effects of one ligand on another in the complex has been

neglected. And the bond lengths of the neutral ligand

molecule were used (30). This certainly is not the ideal

approach, but there is little choice since there is no rea-

sonable way to estimate the magnitude of the change of bond

length.

Table V is a presentation of the INDO spin densities for

the ligand radicals and the corresponding spin densities

calculated from the contact shifts and the relative spin

density of each. An inspection of these densities leads to

the conclusion that, for all of the ligand radicals,the agree-

ment between the relative magnitudes of the calculated spin

densities and, in the case of the cations, the experimental

densities is, in general, fairly good. The BMI1 calculation

produces the best agreement in relative magnitude: prel(INDO)

= 0.166 and 1.00 for CCH3 protons and NCH3 protons, respec-

tively, and prel(exptl) = 0.236 and 1.00 for CCH3 protons and

NCH3 protons, respectively. In comparison, Scarlett et al.

(36) report good agreement for the pyridine N-oxide cation:

prel(INDO) = 1.05, 0.41 and 1.00 for 2-H, 3-H and 4-H,
respectively, and prel(exptl) = 0.81, 0.45 and 1.00 for

2-H, 3-H and 4-H, respectively. Only the GdH radical cation

calculations with sp2 hybridized amine nitrogen orbitals

agree in both sign and relative magnitude videe infra) with

the experimental densities. This may be taken as an indication

of L M charge transfer mechanism of spin delocalization.

In contrast, the calculated spin densities for both the BdH

radical cation and anion with sp2 hybridized amine nitrogen









































CQ)



C)
0



OO


4O
CC
OU








)O
-W H
C)CC



ScO





0t 0


LC
0
1
O C




U 4-4




,(-A
*r-
'* 4
(L T


4-J

















*H-
41






'-4
ct
u






0






I









O
4-J
ed















0C








C)




r-
0

I-



















U
-A
U


cn r-
coo

I N


coo
Cn r-










c10


0 -


-o










Cl








0.o
*-l 7H
Q-4 C









CO


rC'
Td r


oo
co
kfl
I






o0

- 0











1 -4











N





C4
NO


0 n
in kD
I r-


,-- co
4 00

NW






U


r
o N


^~\s


cri

CC


NC3





















c?0
00


0r4
C )


CC

00

0 H r


r--4
r--


)0
U- r-l

r--l 0 0






0

O0



,- co o

cU i C





*r-4

cil 0 rn

Q)
r4 r











-4 -4 -
\D
0






O O's












c-4

+
N


CM
C- 0
















N
cn
cCM








0o4




1r--4 C
-n o
emn















--4







Ci

/)



a a f
pqu


00

,--i
o4




















-.. 0
cq















N
Nl
































C14
00












0 coi

cq
r'N






0cl
0c0
00
























C)
1^~ LD


CM

0 0














o4
C0 r-

0 o
r-4








00













o4
0 r-













N4
























0-4 C
r-4


r-4


































0



O
r-H














c3


M o



.O-I
*r- )
'-c-c


Ni
U)

















Q)
;c>



O O N O C) o C)o
C)o CD ojo C) o7 CD








.--
I








CCO C Ll 00 Ll U0
>r u4

4C-C






Cd 4 CO C O CD
CC)


















00 No NO
-A 0 0 c
4( O. ,O .O






















b0 0
l 0 0 C N 0









U


-C 0 0 r 00 Xu0
uC d
Cd 0)
Cd O ,-0






:> ( o Cs o C'. o


r-,I

0




O, 0




Q I





-U/-
CdC)
,--I










u0 0







H+ L) m +j c o7
UU d co H Z]

















c0
>

















*r-
CO
r--i
























-4
0















r-1
X





























r-4
CO
LI






4-1
o








r4

















U
r.-
X










,-





l-a


r--








o


co
00







nI .O 00 C






-A 0N r-4







o
0-40C
Cn iO c00\







c-4 CM C








r-4








cNI o
0- r-


C) r-o o








N,- r N- n- oo
m co c7 m
c D Co CIN
CN N i-

I







IIN





0-
C


Zn4
WC fi < PQ
u ^ 1


> ,--1 co .."
0








u o



-40 0 0

0 r)-- 0
-I CM a' r1











4 o r-4 o D














,1-
Nr) N ) c-


-4fm







ift 0 %- --

N ONO
0 0 00





oC i-4 0 Ca








I IN







uZ-4
I m






r-41


co
00


--" 0 i o

S-1 00










cn 110 00 a'










-t 0 m m'
r-4 0 r-4 0

0 0 0










clM CrM r-

N -
I





S o

00 Dam-t
0 0 r-4 CM
0,00









0t Crm 0 0


on rn 0 CD





I I








'0 C'1 N
<3 M Q
0-1 '-
I~-0



vo m CM
a \1 0 p


r-4
4-i


cL
0C
a









































H
.,-


0

z













(-



- C




i-l


















4-t
a)




r4
,D
/-4




C)

0































C)
or
4-1
(U

-4




I


0


-u
U
-4
(C







(4
r>

-4




CD
i-u







(U
r-
CCs
U0


-4 0 o'


0 C- 0


r-4
4-1

































O
CL























-0



Ho


CCO
0 0


m -1 00 CV) ^D 00 mo
r--l r-l


4J














41
*r)
o


















41
C



















<)



.4
0














4-J
C:
Cl)


















H
0
cn










































C)
4-1












C)
aH
0













































.-l
0)
























CJ
0)


1)



42
aC
*i-



C)
N






C







*H

*r






*Hr

COl

0)



(0



tO


O *Hr
0 1-4 (U



S0 Cl ,




0 0 0 N



-H
4- j Cl)



co ,-1 1


s:: cd






71 0 eH
C l)















r 4
4 4 r- -(















0 U 0 C'

r--l r-4

0r- co
1 0)
C r- 42 N

















T- 0 0



Q)> l


0 C O0

4- Cu C




t Q) 0l
;j 0 CO


42 U U
Ct .)0 r
U4) l U
S- -

(IC t4 C
(a H






(n) (14 ) -4
) O











CC C 0
c -H 0
*r-4 r 1':





0 O

io c d 3
C



C Cn C
CJ r C 4c
ma- r-
01^ (


N CN CO r-
of o r-- L)
N LO CO r-i
I-







CN 0 L C

LO 0 C(n r-






o o o o






(D 0 CO CD
- co cr- _4













cr) c"I
r 2


o
C




0
rl-






0



o




0
N

CC
7e


0

Q)
42




4

a)





-0


r4
4-1
CC

4-u





cn
















*.r- C
o-
CC



UN








orbitals agree in sign and relative magnitude with the experi-

nentally calculated densities. Thus the charge transfer mech-

anism may be either L M and/or M L. Calculated spin

densities for the BII radical anion agree in both sign and

relative magnitude of spin densities with the experimental

densities. This may indicate M L charge transfer mechanism

of spin delocalization. Although the radical anion calculated

densities for 2,6-PdAdH with sp3 hybridized hydrazone nitrogen

agree with the relative magnitude of the experimental densities,

the signs of the densities do not agree, which means there is

no agreement at all videe infra).

The experimentally determined densities listed in Table

V, calculated using Equation (2), apply to the special case

of normal Curie behavior and are derived from (12,18)

AH = Ah z /N N (3)

where is the time-averaged value of the z component of

the electron spin. So, the problem of calculation reduces

to an evaluation of
= g g S(S + 1) H/3kT (4)
Golding (13) evaluated Sz> for complexes of octahedral site

symmetry. This evaluation included spin-orbit coupling and

magnetic field interactions


Z(2J + 1) Sz > e[-E/kT]
SSz> J (5)
zJ -(2J + 1) e[-/kT]
J


where E is expressed as a power series in H and










(Y Y) V

(6)
This reduces to


3kT (A + (7)

where C is the spin-orbit coupling constant of the transition

metal ion. Golding tabulated values for the parameters y, v,

g, A, and B required to evaluate <(S) for octahedrally coor-

dinated metal ions of dn configuration.

Under octahedral symmetry and by including spin-orbit

coupling, the T1 and E terms split into 2U' + E" + E'

and U' terms respectively. At ambient temperature both the

ground term U' E ) and first excited term E' ( j

are occupied (44). Therefore the summation of
Equation (5) must include these occupied terms. Substitution

into and evaluation of Equation (3) requires two hyperfine

coupling constants (12). Thus there are four unknowns AU,

AE,, E, and r and only one equation. By assuming the free

ion value of c and varying E to produce a best fit, Equation

(3) has been used to calculate coupling constants for spin-

state equilibrium systems (13). However under the noncubic

D3 symmetry of these cobalt(II) complexes the situation is

much more complex. In addition, the free ion value of C is

certainly not appropriate. To be sure, a value less than the

free ion value usually 75% of the free ion value should

be used to account for covalency effects (44). And, under









D3 more than two term levels may be occupied at ambient tem-

perature adding further complications. Furthermore, mixing

of terms by spin-orbit coupling must also be considered.

Although matrix elements for spin-orbit coupling under octa-

hedral symmetry have been calculated previously by Griffith

(14), those for D3 symmetry have not been published. Con-

sequently, it was not feasible to determine coupling constants

and thus spin densities from the observed contact shifts.

Fortunately relative contact shifts and therefore relative

spin densities can be determined using Equation (3) with-

out evaluating That is, there is no dependence on the

distribution over more than one term, as evidenced by the

close agreement of relative spin densities for high-spin

and low-spin values recorded in Table V.

The degeneracy of the t2g and eg orbitals is lifted

under D3 symmetry as shown in Figure 4 (33). The t2g T-

bonding orbitals transform as ag and e" under D ; the e

transforms as e. It is seen that a high-spin or a low-spin

configuration is possible for seven d electrons under D3;

thus, the Oh and D3 situations are qualitatively analogous.

That is, anomalous magnetic behavior can arise under either

symmetry. Considerations of the spin transfer mechanisms

represented in Figure 3 and the electron configurations

given in Figure 4 reveal videe infra): (i) If the electron

configuration is either high-spin or low-spin, (+) spin can

be transferred onto the ligand by L MI or 'I -+ L, respectively.

(ii) If the electron configuration is low-spin, (p) spin













-- e -. e
g

-- a1




-- t2g


is {(e)4(a) 2(e) 1





hs {(e)4l 1 )2
(al) (e)}


7T
e



Oh D3







-- -- e -------------- e
g


ST


- g t2


Is {(a 2(e)4(e) }





hs {(al 2 e)3 )2


7T
al


Figure 4. Possible Oh--D3 correlation diagrams and electron
configurations for a d7 ion under D3 symmetry









can be transferred onto the ligand by L M. (iii) If the

electron configuration is either high-spin or low-spin, both

(t) and (i) spin can be transferred onto the ligand by M L.

The ligand orbitals and metal ion orbitals can interact

and thereby transfer unpaired spin only if they are of

appropriate energy and of the same symmetry. The 7 MO's for

a symmetric bidentate ligand can be classified as (+) or (-)

according to their parity, Figure 5, which either remains

unchanged or is reversed upon rotation about the C2 symmetry

axis (33). The (-) MO's can interact only with the metal

e orbitals, while the (+) MO's can interact with either al or

e metal ion orbitals depending on the choice of C2 axis for

the metal ion orbitals. Table VI is a presentation of the

INDO highest bonding and lowest antibonding T MO's, their

symmetries, (+) or (-), and their energies. These are pre-

sented for the ligand radicals which afford the best fit

with experimental densities.

The reader should apply the general discussion above

and refer to Figures 3 and 4 and Table VI for the discussion

of transfer mechanisms which follows. The calculations of
2
INDO spin densities for the radical cation with sp hybridized

amine nitrogen orbitals gave the best fit with the GdH com-

plex. For this cation the unpaired spin is predicted to oc-

cupy i17, whose symmetry upon rotation about the C2 axis is

(+). Therefore it can interact through L M spin transfer

with e(d) or al(d) metal orbitals. Thus a low-spin configu-

ration should result in (+) spin and a high-spin configuration

should result in (t) spin in i'17.




















I -_' / 7

C2


N+)
N


NC
*N. K-


+)


-x



2


dz2 (+)


C2
dxy (-) or (+)
2


Figure 5. Illustration of (+) and (-) symmetries of selected
's and selected metal orbitals upon rotation
about the C2 axis









Table VI
Selected INDO Molecular Orbital Symmetries
and Energies of Ligand Radicals


Ligand Molecular Occupation Energy
Radical Orbitala Number (Hartrees)


GdH+1 (sp2)


BdH-1 (sp2)


BdH+1 (sp2)


BMI-1


- 0.627005

- 0.573568

- 0.066018


16 (-)

17 (+)
18 (-)



23 (+)
24 (-)


25


2'22
423

24

25


222

0'23

'24

J25


(a)b


(-)


(+)

(-)

(+)




(c)b
ab


- 0.072167

0.124016


0.513461


- 0.587260

- 0.552269

- 0.050940

0.058413


- 0.086492

- 0.083043


0.107814

0.496251


(+)


an MO's are classified (+) or (-) depending on whether they
remain unchanged or reverse sign upon rotation about their
C2 axis (32). The MO's are numbered in order of increasing
energy.

bThese are molecular orbitals and are not classified further.








The calculations of the INDO spin densities for the

radical cation and anion with sp2 hybridized amine nitrogen

orbitals give the best fit with the BdH complex. The anion

calculation predicts the unpaired spin to occupy 1i24, its

symmetry upon rotation about the C2 axis is (-). Therefore

it can interact through M L spin transfer only with the

al(d) metal orbital. Consequently, a low-spin configuration

should result in (+) spin and a high-spin configuration

should result in (t) spin in q924. In addition, it may well

be that 25 is energetically accessible. If it were acces-

sible it could accept (+) spin and/or (4) spin from a e (d)

orbital. The BdH cation calculation of INDO spin densities

predicts the unpaired spin to occupy 123' whose symmetry

upon rotation about the C2 axis is (+). Therefore, it can

interact through L M spin transfer with e(d) or a (d) metal

orbitals. Thus a low-spin configuration should result in

(+) spin and a high-spin configuration should result in (t)

spin in 23.

The radical anion calculation of INDO spin densities

gives the best fit with the BMI complex. The unpaired spin

is predicted to occupy P24, whose symmetry upon rotation

about the C2 axis is (-). Therefore it can interact through

M L spin transfer only with the al(d) metal orbitals. Thus

a low-spin configuration should result in (+) spin and a high-

spin configuration should result in (t) spin in r24'

A discussion of the 2,6-PdAdH complex is deferred,

because the complex is of D2d symmetry rather than D3 symmetry,

it will be discussed separately videe infra).









Because it also places opposing spins at the same posi-

tion in the ligand, spin polarization adds further ambiguity

to the interpretation of the contact shifts in these five-

member (or any odd member) chelate ring complexes. If, for

example, the spin is delocalized into a a orbital, this places

metal spin into s orbitals of all atoms. Also, the bulk of

the delocalized spin resides in the lone pair of the coor-

dinated nitrogen atom causing a "slight unpairing of all

filled a bonds" by spin polarization effects (23). This is

readily seen for GdH by structures such as those in Figure 6.

If (,) spin is considered to be placed in a MO at the

coordinated nitrogen, the other electrons associated with

that nitrogen atom will also align with (t) spin preferentially

(Hund's rule of maximum multiplicity). The electron associated

with the adjacent atom must be paired, (+), with the nitrogen

atom electron. All the electrons associated with the adjacent

atom must also align (+) (Hund's rule), and so on through

the molecule. Thus spin polarization places (+) spin and

(+) spin at each position simultaneously. These considerations,

when applied to the other ligands, lead to the same con-

clusions for each of the other ligands. Thus the ambiguity.

The foregoing discussion certainly indicates there are

many difficulties in reaching definitive conclusions about

the contact shifts of these compounds. However, considera-

tion of how the lack of agreement between INDO calculated

and experimentally determined spin densities relates to.low-

spin contributions may be more enlightening. In each instance






56








i- +Ht



C+ + C


H+ fH



















+C+ tC

Ht +H


N+ tN' WN4 tNN
4- 1 41

Ht ....M +H









Figure 6. Representation of spin polarization effects in a
fragment of a [M(GdH)++] complex









for which the signs and relative magnitudes of the calculated

densities agree with the experimentally determined densities

(Table V), the largest density is on the protons of the

amine or imine moiety; this density is always positive. The

smaller density is on the remaining protons of the ligand;

but the agreement in relative magnitude is only fair. This

may indicate contribution from both high-spin and low-spin

species. As can be deduced from the transfer mechanisms in

Figure 3 and the electronic configurations in Figure 4, con-

tribution from both configurations would place (t) spin

and (+) spin at the same proton. For example, if the elec-

tronic configuration is Is {(e) (a) 2(e) 1, the unpaired

spin occupies the e MO and will align parallel with the

external magnetic field. The transfer must be L -* M. Con-

sequently, (-) spin will remain on the ligand. By comparison,

if the electronic configuration is hs {(e) (a) (e) 2}, the

unpaired spins occupy the a1 and e and will also align

parallel with the magnetic field. The transfer mechanism

may be either L M (L e ) or M L (a -T IT). If it is

L M1, (+) spin will remain on the ligand. Consequently, if

there is low-spin contribution to the contact shift, the

effect of the opposing spins on the ligand would be a smaller

shift and therefore a smaller relative density than that

calculated.

The preceding syllogism can be applied to the contact

shifts of these cobalt(II) compounds. Susceptibility studies

yielding intermediate, room temperature, magnetic moments of









4.65 BM and 4.23 BM for the BdH and GdIH complexes, respec-

tively, are taken as an indication of relative contributions

from both high-spin and low-spin configurations. The low-

spin contribution is about the same in the GdH complex as in

the BdH complex and is presumably reflected in the similar

disagreement between the INDO calculated and the experi-

mentally determined relative magnitudes of spin density.

Consequently, it might be inferred that the effect of a con-

tact shift contribution attributable to the low-spin (2E)

configuration is observed in these isotropic shifts. By

comparison, the relative magnitudes of the experimentally

determined densities for the BMI complex are much larger -

not smaller than those calculated by the INDO method. The

intermediate, room temperature, magnetic moment of 3.30 BM

of the BMI complex also indicates contribution from both

high-spin and low-spin configurations, and indicates a greater

relative contribution from the low-spin configuration than

in the BdH and GdH complexes. It may be that in the BMI

complex the high-spin and low-spin contributions-reenforce

rather than oppose one another. This mutual reenforcement

of high-spin and low-spin contribution may also account for

the larger shift of the BMI methyl protons compared to the

BdH methyl protons.

It should be reemphasized at this point that with the

present stage of development of theory it is not possible








to distinguish with certainty between contributions from

the 2E term and spin polarization involved in these iso-

tropic shifts. As pointed out in the introduction contact

shifts of cobalt(II) complexes are previously known to arise

from the high-spin configuration ( T under Oh) only.

The 2,6-PdAdH complex was not considered in the pre-

ceding discussion because its symmetry is D2d rather than

D Under D2d symmetry 2E (Oh) + A + Bl(Dd) and T (Oh

- B2 + E(D2d) (6). As under D3 all terms contribute zero

orbital angular momentum (A, B, and E). Consequently, the

isotropic shifts are due entirely to contact interaction.

Therefore pseudocontact contributions should be negligible.

The possible d orbital correlation diagrams and electronic

configurations for D2d are shown in Figure 7. Based upon

the room temperature magnetic moment of 2.28 BM, a low-spin

configuration should apply. This value is effectively

within the range generally accepted for six-coordinate low-

spin cobalt(II) complexes (1). Nonetheless, a high-spin

configuration is also feasible under D2d, a condition which

can give rise to a degree of anomalous magnetic behavior as

is possible under octahedral symmetry.

None of the INDO spin density calculations agrees in

either sign or relative magnitude with the experimental den-

sities of the 2,6-PdAdH complex. Consequently no assertion

can be made as to whether the spin transfer mechanism is

M L and/or L M. The complete lack of agreement for the

2,6-PdAdH complex completes the trend discussed above for
















- -e
9


al + b


s ((e)4(b2) (al or bl) }



hs {(e)4(b2) l(al) 1(b)}


--e




D2d


- -e


al + b1


- e


-t2g.


Is {(b2)2(e)4(a or b1 I



hs {(b2) 2e) 3 (a)1 (b)


D2d


Figure 7. Possible 0h--D2 correlation diagrams and electron
configurations or a d7 ion under D2d symmetry


- b2


- 2g.








the BdH, GdH and BMI complexes. That is, increasing dis-

agreement between INDO calculated spin densities and experi-

mentally determined spin densities may reflect a greater

low-spin contribution in each complex. The low-spin con-

tribution is smaller in the BdH complex and the agreement

is better. By comparison, the low-spin contribution is

greater in the 2,6-PdAdH complex and there is no agreement

between the INDO calculated and experimentally determined

spin densities.

It appears that the low-spin configuration may indeed

contribute to the contact shifts of these cobalt(II) com-

plexes. It is also of interest to note that the hydrazone

amine and imine proton resonances are shifted downfield,

and are the largest shifts observed in each complex. There-

fore it is reasonable to conclude that they are dominated

by the same delocalization mechanism. Since delocalization

into a a 10 leads irrevocably to a downfield shift (23), it

is tempting to propose that delocalization into a a MO

dominates at these positions with an opposing 1 MO contribu-

tion which makes an upfield contribution to the contact shift.

This postulate is in keeping with that of Stoufer and co-

workers (41) that f back bonding plays an important role in

these anomalous cobalt(II) complexes.

Although attempts were made to prepare and investigate

the proton contact shifts of cobalt(II) complexes of all the

ligands listed in the introduction, these attempts did not

meet with success in some cases. Specifically, it was not

possible to prepare the complexes of the ligands in Figure 2.







The tris-glyoxalbismethylimine cobalt(II) complex decomposes

so rapidly it can not be isolated in pure form. The complexes

of the remaining ligands could not be prepared, probably be-

cause of steric and/or electronic effects (20). That is,

either intramolecular steric strain or the electron with-

drawing capabilities of the substituents was sufficiently

great to prevent coordination of the ligand to the metal ion.

In conclusion, although the modified INDO program used

was capable of handling the minimum basis set of the simple

tris-glyoxaldihydrazone cobalt(II) complex, the spin den-

sities calculated were so small as to be of little value.

Consequently the commonly used approach of treating each

ligand as a radical anion and/or cation was used for this

work. However the various ambiguities which arise from this

treatment (spin polarization, delocalization into odd number

chelate ring complexes, uncertainty as to relative energies

of ligand orbitals upon interaction with the metal ion)

require caution in making very definitive conclusions. There

is evidence, however, that T back bonding plays an important

role, and that there is contribution to the contact shifts

attributable to the low-spin, 2E(D3) derived from 2E (h)'

configuration.

This investigation of contact shifts of magnetically

anomalous cobalt(II) complexes has lead to essentially the

same conclusion reached by Bertini and Getteschi (3). "The

enthusiasm of researchers in the proton contact field has

recently attenuated owing to the lack of a simple model

correlating the observed shifts to the nature of the





63

increasing sophistication have been proposed and laborious

calculations have been performed which gained in some cases

partial success but never found general applicability. In

many instances however these partial models have been the

origin of lively controversies."















APPENDIX


Chemical shift assignments for the proton magnetic res-

onance spectra of the 2-pyridinalphenylimine ligands are

listed in Table VII. They are based on the assignments

made by Matsubayashi, Okunaka, and Tanaka, which were made

for substituted 2-pyridinalphenylimines, in which the 4-

phenyl substituents are OCH3, CH3, and C1 (31). The follow-

ing numbering system applies.



4
5 r 3 a
6 1' YN X

H X
a





Table VIII is a list of the chemical shifts of the tris-

(2-pyridinalphenylimine) complexes. The purpose in preparing

these complexes was to obtain and investigate contact shifts

of a homologous series of complexes which span both sides

of the cross-over point. Because this situation did not ob-

tain, Table IX, the compounds are of little value for this

purpose. Furthermore the symmetry of these complexes is no

higher than D3 (face isomer) and may be C1 (edge isomer).










Table VII
Proton Chemical Shifts of the Substituted
2-Pyridinalphenylimine Ligands


Position H NH2 CH3 Cl 2-PAH

H3 8.15 8.15 8.17 8.16 7.96

H4 7.83 7.85 7.85 7.86 7.73

H5 7.45 7.39 7.41 7.46 7.23

H6 8.66 8.67 8.69 8.70 8.63

H 8.56 8.62 8.61 8.58 8.13

H 7.30 7.25 7.23 7.31 a

H 7.30 6.71 7.23 7.38

H 7.30 4.22 2.34
x


Note: Chemical shifts are
at 30 C.


in ppm's relative to TMS


in CD3CN


aThere are no phenyl protons in the 2-PAH ligand. However,
the NH2 chemical shift in d6-DMSO was found to be 4.13 ppm
from TMS.










Table VIII
Proton Chemical Shifts of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II) Complexes


[Co(2-PAH)]3


















[Co(2-PAI) ]


7.1
7.6
15.7
51.2
61.7
64.2
81.7
107.8
128.8
163.7
174.7
182.7


53.4
42.0
32.3
- 29.8
- 1.9
3.9
5.4
7.3
9.5
14.1
15.1
15.4
50.7
53.3
58.2
67.4
68.7
72.2
92.3
144.7


[Co(2-PPI) ]


- 51.5
- 40.3
- 25.5
- 19.6
- 14.8
- 11.0
- 8.0
- 6.5
4.2
6.8
8.9
10.4
11.0
16.0
16.6
17.6
18.7
51.5
58.1
67.3
70.3
72.1
76.6
87.8
111.5
122.2










Table VIII continued



[Co(2-PCI)3 53.2 [Co(2-PMI) ] 51.1
40.4 38.6
28.8 28.1
21.1 20.7
6.3 6.2
8.3 8.2
11.1 10.0
15.1 10.6
16.1 11.6
17.9 16.0
19.4 16.4
51.8 16.9
59.3 17.7
69.6 51.5
72.7 57.1
78.1 58.0
89.5 64.4
112.0 70.8
125.0 74.5
77.5
97.5
115.7
Note: Chemical shifts are given in ppm from TMS in CD3CN
at 300 C.










Table IX
Effective Magnetic Moments of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II) Compounds


Compounds Oeff (BM) 'eff(soln) (B8)a

Co(2-PAH)3(BF4)2 b 4.61
Co(2-PPI) (BF4)2 4.68 4.60

Co(2-PAI) 3(BF4) 4.64 4.09
Co(2-PMI)3 (BF4)2 b 4.45

Co(2-PCI)3(BF4)2 4.42 c


aDetermined by the method of L6liger and Sheffold (27).

Undetermined

CCo(2-PCI)3(BF )2 was not sufficiently soluble to determine
Ueff in solution.






69

In addition, a mixture of the two isomers is obtained; there-

fore, two superimposed spectra resulted. Interpretation of

contact shifts of a mixture of these complexes of such low

symmetry would be even more difficult than for D3 complexes.

Consequently, neither assignments of the contact shifts nor

a discussion of their significance was attempted.















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BIOGRAPHICAL SKETCH


Joseph B. Vaughn, Jr, was born April 21, 1942, in Elfers,

Florida. He graduated from Paxon High School in Jacksonville,

Florida, in June, 1960. After a year of work and frugality,

he entered the University of Florida in 1961 and completed

his freshman year. After three more years of work, frugality,

and military service, he returned to the University of Florida

in September, 1964, and received the Bachelor of Science

degree with Honors in Chemistry in June, 1968.

Graduate study began in September, 1968, at the University

of Florida. From September, 1968, until June, 1973, he held

a teaching assistantship, and in September, 1973, no longer

eligible for an assistantship, be began work for the Pharmacy

Department at the University of Florida. From December, 1973,

until June, 1976, he held a medical technician position at

the V. A. Hospital in Gainesville, Florida. He also held an

Instructor's position, June, 1976, to September, 1976, with

Santa Fe Community College in Gainesville, Florida. Returning

to his graduate work at the University of Florida in September,

1976, he received the Ph. D. degree in August, 1977.











I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philos phy.




R. Carl Stoufer, Chairman
Associate Professor'of Ch istry







I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.




Wallace S. Brey, Jr.
Professor of Chemistry ]







I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.




Gerhard M. Schmid
Associate Professor of Chemistry











I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.




Richard D. Dresdner
Professor of Chemistry







I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.




Paul Urone
Professor of Enviromental
Engineering Sciences







This dissertation was submitted to the Graduate Faculty
of the Department of Chemistry in the College of Arts and
Sciences and to the Graduate Council, and was accepted as
partial fulfillment of the requirements for the degree of
Doctor of Philosophy.

August 1977


Dean, Graduate School





























































UNIVERSITY OF FLORIDA


3 1262 08553 5226




Full Text

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PROTON CONTACT SHIFTS OF SOME MAGNETICALLY
ANOMALOUS COBALT(II) COMPLEXES
By
JOSEPH BENJAMIN VAUGHN, JR.
A DISSERTATION PRESENTED TO THE GRADUATE COUNCIL
OF THE UNIVERSITY OF FLORIDA
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1977

TABLE OF CONTENTS
Page
LIST OF TABLES a±
LIST OF FIGURES iv
ABSTRACT v
INTRODUCTION 1
EXPERIMENTAL 14
Preparations ......... 14
Elemental Analysis ....... 24
Mass Spectrum Analysis ...... 24
Spectrometers ........ 24
RESULTS 26
DISCUSSION 34
APPENDIX 64
BIBLIOGRAPHY 70
BIOGRAPHICAL SKETCH 73
ii

LIST OF TABLES
Table
Page
I
Selected Infrared Absorptions and
Assignments for the Co(II) and Fe(II)
Compounds Studied .......
27
II
Magnetic Moments and g Values for the
Co(II) and Fe(II) Compounds Studied
30
III
Observed and Relative Isotropic Shifts
and Half-Height Peak Widths for Cobalt(II)
Compounds .........
32
IV
INDO Calculated and Experimental Proton
Spin Densities for [Co(GdH)^]
39
V
INDO Calculated and Experimental Proton
Spin Densities for Cationic and Anionic
Radicals .........
42
VI
Selected INDO Molecular Orbital Symmetries
and Energies of Ligand Radicals
53
VII
Proton Chemical Shifts of the Substituted
2-Pyridinalphenylimine Ligands
65
VIII
Proton Chemical Shifts of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II)
Complexes .........
66
IX
Effective Magnetic Moments of the Sub¬
stituted Tris(2-pyridinalphenylimine)
Cobalt(II) Compounds ......
68
iii

LIST OF FIGURES
Figure Page
1. Ligands whose cobalt(II) complexes were
successfully prepared ........ 10
2. Ligands whose cobalt(II) complexes could
not be prepared ......... 11
3. Possible spin transfer mechanisms . . . .37
4. Possible 0^--Do correlation diagrams and
electron configurations for a dV ion under
symmetry .... T ..... 50
5. Representation of ligand tt MO and metal ion d
orbital parity upon rotation about the C2 axes . . 52
6. Representation of spin polarization effects
in a fragment of a [M(GdH)^+] complex . . . .56
7. Possible 0v1--D2ci correlation diagrams and
electron configurations for a d7 ion under
symmetry 60
IV

Abstract of Dissertation Presented to the
Graduate Council of the University of Florida
in Partial Fulfillment of the Requirements
for the Degree of Doctor of Philosophy
PROTON CONTACT SHIFTS OF SOME
MAGNETICALLY ANOMALOUS COBALT(II) COMPLEXES
By
Joseph Benjamin Vaughn, Jr.
August 1977
Chairman: R. Carl Stoufer
Major Department: Chemistry
A series of cobalt(II) complexes known to lie near the
2 2 A
E ( G)-- T-, ( F) cross-over point was prepared. The first
o ^
reported proton contact shift investigation of such magnetic¬
ally anomalous cobalt(II) complexes was carried out in order
to obtain information about the metal--ligand bonding proper¬
ties. These complexes contain ligands with the a-diimine
linkage (glyoxaldihydrazone, GdH; biacetyldihydrazone, BdH;
biacetylbismethylimine, BMI; 2,6-pyridinediacetyldihydrazone,
2,6-PdAdlI), which provides relatively large ligand field
splitting. The symmetry of the complexes containing these
symmetric, bidentate ligands is D^ and that of the terdentate
2
ligand complex is D3tp Under D^ symmetry ^g(^) transforms
as 2E(D3)
and
Vv
transforms as '4A^(D3)
and
4E(D3).
Under
D9(3 symmetry
2V°h>
2
transforms as A-^(D9(3)
and
VD2d>
and
VD2d>
and
Vv
transforms as q'B9(D9(3)
and
4E(D2d)•
Because these terms are isotropic, the dipolar shift
v

contribution to the contact shifts is neglected. The magni¬
tudes of these shifts (100 ppm and less) are typical of those
observed for other complexes.
Spin densities were calculated using a modified INDO
program. The approach commonly used in such calculations is
to treat a single ligand as a radical cation or anion, using
only the "it" symmetry atomic orbitals in some instances, pre¬
sumably due to the limited capacity of the available programs.
In contrast, the program used for these calculations is capable
of handling the complete valence basis set (ninety-nine
_j |_
atomic orbitals) for [Co(GdH)^ ], the simplest complex in
the series. That is, metal penultimate d and ultimate £ and
£ orbitals and all the ligand valence orbitals were included.
But, meaningful results for even this simple complex were
not obtained, presumably because the number of unpaired spins
(three for high-spin complexes and one for low-spin complexes)
is sufficiently small that the spin densities calculated over
ninety-nine basis orbitals represent only "noise." Thus it
is concluded that such calculations do not yet have the degree
of sophistication necessary to provide meaningful results
++
even for such a simple complex as fCo^dH)^ ]• Consequently
each ligand was treated both as a radical anion and a radical
cation, as is commonly done, including the complete valence
basis set. Evaluation of these INDO calculated and experi¬
mentally determined spin densities leads to the following con¬
clusions .
(1) Because the best fit was obtained for the GdH4^,
vi

BdH+^ and BdH ^, BMI ^ , and no fit was obtained for 2,6-
PdAdH, it might be concluded that the spin transfer mechanism
is :
(a) L -*â–  M in the complex [CoCGdH)^].
(b) L -+ M and/or M -*â–  L in the complex [Co(BdH)^1"].
(c) H -â–º L in the complex [Co (BMI.
(d) Undetermined in the complex [Co (2,6-PdAdH)^] .
However, such conclusions are somewhat tenuous because of the
possibility of spin polarization, which adds further ambigu¬
ity by placing opposing spins simultaneously at the same
position in the ligand.
(2) More than one molecular orbital is involved in the
transfer mechanism in each complex, making definitive conclu¬
sions somewhat more difficult. Comparison of the INDO molec¬
ular orbital calculations and the experimental data provides
evidence that spin transfer onto the hydrazone amine protons
occurs via a o molecular orbital with an opposing -rr molecular
orbital contribution, in keeping with the usual assertion
that dTT--pir back bonding is involved in these complexes.
(3) There is increasing lack of agreement within the
series of ligands, between INDO calculated spin densities and
the experimentally determined spin densities, which corresponds
to increasing contribution of the low-spin configuration to
the ground state of the complex. This trend is taken as an
indication that there is a contribution to the contact shifts
attributable to the low-spin configuration, the first such
contribution reported for an octahedrally coordinated d^ ion.
Vll

INTRODUCTION
During the past century or so a great deal of research
has been directed toward the investigation of transition
metal coordination compounds. A significant portion of this
work has involved octahedral cobalt(II) compounds, the ma¬
jority of which are high-spin, y eff = 4’7 - 5'2 BM, having
a ^Tig(^F) ground term. There are relatively few octa¬
hedral cobalt(II) coordination compounds which are low-spin,
ye££ = 1.8 - 2.2 BM, having a 2E (2G) ground term. Since
O
1956 (40), there have been reported an increasing number of
magnetically anomalous cobalt(II) coordination compounds,
i.e., compounds which exhibit both room temperature magnetic
moments intermediate between high-spin and low-spin moments
and anomalous Curie-Weiss behavior.
Variable temperature magnetic studies led Stoufer and
coworkers (41) to postulate that these anomalous compounds
lie near the "cross-over" point and that Boltzmann distri¬
bution over high-spin and low-spin states is attained, i.e.,
AE (ZEg--^T-j^g) = kT. They noted that the majority of these
compounds contain ligands , with an ct-diimine linkage
(-N=C-C=N-), which provide relatively large ligand field
splittings due in part to their ir-acid capabilities. Research
interest in these compounds still exists as evidenced by
1

2
the recent work of Simmons and Wilson (39), who investigated
solution and solid state magnetic and spin lifetimes of some
six-coordinate bis(N-R-2,6-pyridinedicarboxaldimine) cobalt(II)
complexes. For theR=NH(CH^) derivative, they report solution
rates for 2E(ls) 4T(hs) of ^ = 3.2 x 106 s-1 and k_^ =
9.1 x 10° s with corresponding spin state lifetimes of r(zE)
= 3.1 x 10"7 s and t(4T) = 1.1 x 10"7 s.
Information about bonding properties in paramagnetic
transition metal complexes as inferred from unpaired spin
distribution is traditionally obtained by electron spin
resonance techniques, where the isotropic couplings are mani¬
fest as nuclear multiplet structure. The sign of the cou¬
pling constants cannot be determined directly from the ESR
experiment. However, the sign is obtained directly from
the NMR contact shift, and thus the sign of the spin density
is gotten directly. Because it is the sign of the spin den¬
sity which gives the most useful information about bonding,
the major advantage is with the NMR technique. Furthermore,
the spectral complexity is greater in the ESR experiment.
That is, for a system with n coupling constants, ESR spectra
will exhibit 2n lines; whereas NMR spectra will exhibit only
n lines. However, the two techniques are more complementary
than competitive, since it is rare that both can be usefully
employed for the same system. This is because of the re¬
quirements of very short (<. 10 ^ s) electron-spin relaxation
time, T-^ , for observing NMR spectra and long T-^'s (> 10 s)
for observing ESR spectra.

3
In their review article, Lewis and Morgan (26) make
a general distinction between species containing one un¬
paired electron (s = %) and those with more than one (s > 1).
For S = % systems, electron-spin relaxation is very ineffi¬
cient, yielding well-resolved ESR spectra. Exceptions can
occur in systems possessing low-lying excited states, which
reduce T^e considerably. For systems with S > 1 the presence
of orbital degeneracy in the ground state, low-lying excited
states and/or modulation in solution of zero field splitting
can produce rapid electron-spin relaxation and therefore
well-resolved NMR spectra.
Because T-^e is typically long for complexes in which
o
S = %, it has been predicted (42) that low-spin, E octa-
hedral cobalt(II) complexes will not yield observable NMR
contact shift spectra. And — to date no NMR data have been
reported for any low-spin, octahedral, d^ ion (42). This is
due, in part, to the fact that there are no proton-containing,
low-spin, octahedral cobalt(II) complexes known. It may be
2
that the E species present in anomalous cobalt(II) com-
O
plexes investigated in this work will have T-^e shortened
sufficiently to observe NMR spectra by the presence of the
low-lying T-^ excited state near the cross-over point (vide
infra).
Nuclear magnetic resonance spectra of paramagnetic
transition metal complexes typically exhibit very large
(200 ppm and larger for protons) contact shifts. Contact
shifts are measured relative to a diamagnetic standard,

4
usually the free ligand or, better yet, a homologous dia¬
magnetic metal complex. And, the shifts may be related to
the transfer of spin either to or from the ligand. The
ligand is usually a closed shell diamagnetic molecule, where¬
as the metal ion may possess some unpaired electrons in in¬
completely occupied d or f orbitals. If any of the metal
orbitals containing the unpaired electron(s) is involved in
M-L bonding, the unpaired electron(s) may be transferred to
a new molecular orbital (MO) of the complex. Consequently
the unpaired spin may be partially delocalized onto the
ligand. McConnell and Chestnut (28) showTed that for ligands
possessing nuclei with 1^0, this delocalized spin may
interact with these nuclei via the Fermi contact interaction
(8). For a complex possessing m spins, the Hamiltonian for
the Fermi interaction is given by ~R = A Sm,I. A is the
hyperfine coupling constant for the contact interaction
between the electron spin, S, and the nuclear spin, I, and
is given by (36)
A = 3 71 ge gN 6e 3N _1 p(r) I '(> (D
where p(r) is the spin density operator at the nucleus, ip is
the electronic wave function for the molecule, is the
time averaged value of the z component of the electron spin,
and the other symbols have their usual meanings. Thus Fermi
coupling constants arise only from spin in s_ orbitals, and
A is a linear function of that spin density. The variation
in coupling constants for the various nuclei may lead to

5
the determination of the distribution of unpaired spin density
in the ligand, and thus to the characterization of the MO
containing the unpaired spin. Consequently a better under¬
standing of electronic structure in transition metal complexes
may be achieved.
Both the sign and magnitude of A can be determined from
contact shift spectra. A is related to the contact shift
A v by (10,28)
fiv , _ A Se Be S(S + ^ (2)
V gN 3 kT
if the Curie law is obeyed. It is clear from equations (1)
and (2) that a downfield shift in resonance position corre¬
sponds to a positive value for A and therefore to positive
spin on the corresponding nucleus.
La Mar has outlined the three general areas of considera¬
tion for the elucidation of the observed coupling constants
in terms of the M-L bonding: "(i) The relationship between
the observed coupling constant for a given nucleus and the
probability (density) of finding the unpaired spin in an
atomic orbital at the nucleus of interest and/or at a neigh¬
boring nucleus, particularly as to how the coupling constant
reflects the type (a or tt) of ligand molecular orbital con¬
taining the spin; (ii) the evaluation of the mechanisms by which
the unpaired spin on the metal is transferred to the ligand;
and (iii) the qualitative and, where possible, the quantitative
relationships between the fraction of total spin transferred
and the M-L covalency" (23 :87).

6
There are several possible problems which may introduce
uncertainties into the interpretation of contact shift data.
Three major ones are: (i) Separation of the dipolar or
pseudocontact shifts from the measured isotropic shifts. These
dipolar shifts occur only in complexes which are magnetically
anisotropic, and are produced by the dipolar interaction
between the electron's magnetic moment, y, and the nuclear
magnetic moment h y^ I. The dipolar shift averages to zero
by tumbling in solution unless the electronic magnetic moment
is anisotropic. It can, under appropriate conditions, be
separated from the isotropic shift by (36)
Av
dip
D (3 cos ° ' b (gM - g1)
(3)
where D is a positive quantity whose value depends on the
temperature and various relaxation times (17,21,22,29) and the
other symbols have their usual meaning. Due to uncertainties
in D, most published calculations of absolute dipolar shifts
are open to question (7). However, if the g value is iso¬
tropic or nearly so (gj. - - 0) dipolar shifts can be neg¬
lected. (ii) Electron correlation effects can cause spin
density to be induced on L in the M-L fragment, even if the
metal orbital containing the spin is symmetry restricted
from interacting with any orbital on L. These effects can be
accounted for by use of an appropriate quantum mechanical
model for calculated spin densities. (iii) The possibility
of spin-polarization effects can also cause uncertainties in
determining the mechanism of spin delocalization. Again the

7
appropriate choice of quantum mechanical model can help over¬
come this problem.
The mechanism by which the unpaired spin is delocalized
onto the ligand is related to the symmetry of the MO containing
it. These MO's are predominantly metal d orbitals with smaller
contributions from ligand orbitals which may be filled tt or-
titals, empty it orbitals, a orbitals, or a mixture of all
three. Symmetry considerations sometimes restrict the possi¬
bilities, in complexes of high symmetry for instance. But —
in general there is no easy method available for sorting out
the various contributions. Quantitative evaluation of M-L
covalency is no simple task, and the present development of
the evaluation of spin-delocalization mechanisms does not
permit a definitive presentation (23). "Thus the theoretical
complexities of relating transferred spin density by any
single mechanism to covalency, plus the practical problems
of multiple delocalization mechanisms, suggest that caution
should be exercised in any quantitative interpretations"
(23:120).
The ideal procedure for theoretical analysis would be
a complete, many-electron calculation on the entire complex
with a concomitant breakdown of the isotropic shifts into
their component parts. This would be a formidable undertaking
even for a simple complex. The nonideal approach commonly
employed is to treat a single ligand molecule as a free radical
anion or free radical cation depending on the delocalization
mechanism involved. The spin densities are then calculated

8
from the MO's containing the unpaired electrons and a corre¬
lation is attempted between the calculated spin densities
and the relative spin densities gotten from the observed
contact shifts (vide infra). In some cases, experimental
results satisfactorily fit calculated spin densities assuming
only one ligand MO is involved in the delocalization. In
other cases, it appears more than one ligand MO is involved.
In any case, this extreme approximation leaves much to be
desired regarding any quantitative conclusions. One obvious
objection to this approach is that the metal orbital contri¬
butions are excluded in forming the MO's. The metal is con¬
sidered to be simply an electron source or sink for the
ligand. This obviously assumes the relative energies of the
ligand MO's are unaltered by interaction with the metal.
Certainly this situation is seldom achieved.
There are several theoretical methods available for
treating spin-density calculations in metal complexes. In
this work, the INDO method was chosen because it was readily
available and it is more realistic than the others for the
following reasons (23,34,36):
(i)It is an all valence electron method.
(ii)It is capable of calculating negative spin densities
arising from electron correlation since it uses unrestricted
wave functions. The extended Hückel method, for example, can
calculate only zero or positive densities.
(iii)It accounts for spin polarization by including one-
center exchange integrals.

9
On the negative side, INDO calculations have "black-box"
aspects which tend to hinder a qualitative understanding of
the resultant calculations.
To date no calculations reported have included the metal
ion orbitals and all the ligand valence orbitals in the basis
set. This is because the present theoretical development is
*
simply not capable of handling such a large basis set. However,
one complex investigated in this work [Co(GdH)^"] does have
a sufficiently small basis set (ninety-nine atomic orbitals)
to include the metal ion orbitals and all the ligand orbitals.
The purpose of the present work is to use NMR proton
contact shifts to investigate the metal-ligand (M-L) bonding
properties of a series of anomalous cobalt(II) compounds.
The bis(2,6-PdAdH) cobalt(II) complex and tris cobalt(II)
complexes of the other ligands in Figure 1 were prepared for
this purpose. Ligands with X = -CN and -N09 were also pre¬
pared, however, they did not coordinate to the metal ion. The
"tris" cobalt(II) complexes of GdH, BdH, BMI and the bis-
(2,6-PdAdH) cobalt(II) complex have been prepared previously
2
and are known to lie near the cross-over point with a E
v g
ground term (5,10,37,41,44). The 2-pyridinalphenylimine
complexes were prepared with the expectation that increasing
the TT-withdrawing capability of the substituent would give a
homologous series of complexes which would exhibit magnetic
moments over the range of high-spin and low-spin values,
the ultimate goal for this series being to relate contact
shift changes to the Hammett acidity function. To accomplish
this goal the ligands in Figure 2 were prepared or their prep¬
aration attempted.

10
H •
/
H
.c—cvS
h2n-n* 'n- nh.
H3c \ /CH
^C~CVN
H2N -N N-NH2
Glyoxaldihydrazone
(GdH)
Biacetyldihydrazone
(BdH)
h2cn CH~
H3C-N n-ch3
Biacetylbismethylimine
(BMI)
2,6-Pyridinediacetyl-
dihydrazone
(2,6-PdAdH)
TT
N-NH
2
2-Pyridinalhydrazone
(2-PAH)
X = H, CHo, NH0,
r
."N'
1 /
f ;
N
2-Pyridinalphenylimine
(2-PPI)
Figure 1. Ligands whose cobalt (II) complexes were success
fully prepared
Cl

11
Figure 2. Ligands whose cobalt(II) complexes could not
be prepared

12
X = H, OH, OCH3, N02
H H
X = H, Br, CN

13
The rationale underlying this choice of ligands was to
prepare a homologous series of ligands (containing the a-di-
imine linkage) whose cobalt (II) complexes should lie near
the cross-over point. The substituents, X, have known a-and
iT-acid effects. It is known that the ir-acid character of
the ligand is important in the formation of the octahedral,
low-spin cobalt(II) complexes. Therefore, it was expected
that a series of ligands with increasing n-acid effects would
produce a homologous series of complexes some of whose
members have a ^T-^ ground term and some of whose members
2
have a E ground term (i.e., span both high-spin and low-
O
spin sides of the cross-over point). And, it was expected
that the resultant 7r-acid capability could be related to the
change in proton contact shifts.

EXPERIMENTAL
In the preparations described below, melting points of
previously prepared compounds are all within experimental
error of those reported. The yields for all the compounds
were rather poor, and varied considerably upon repeating the
preparations. These poor yields were a result of washing or
recrystallizing the soluble product many times in order to
obtain a pure product. Consequently yields are not reported.
For several compounds, mass spectra parent peaks and their
relative intensities are reported in lieu of elemental analy¬
sis.
Preparations
Biacetyldihydrazone, BdH. BdH was prepared by the method
of Fisher (10). M. p., 159°C.
Anal , : Caled, for C^qN^: C, 42.11; H, 8.77; N, 49.12.
Found: C, 42.23; H, 8.91; N, 49.02.
Glyoxaldihydrazone, GdH. GdH was prepared by the method
of Fisher (9). However, better yields were obtained using
only 40 ml of hydrazine rather than 75 ml. M. p., 99°C.
Anal.: Caled, for C9HgN^: C, 27.91; H, 6.98; N, 65.12.
Found: C, 27.78; H, 7.11; N, 64.85.
14

15
2,6-Pyridinediacetyldihydrazone, 2,6-PdAdH. A solution
of 50 ml of methanol and 4.0 ml (0.12 mole) of hydrazine was
prepared and placed in an ice bath. Another solution of
8.1380 g (0.0498 mole) of 2,6-diacetylpyridine dissolved in
50 ml hot methanol was added dropwise with stirring. The
white powdery product was washed with cold methanol and dried
in vacuo over P^O^q. M. p. , 179°C.
Anal. : Caled, for CgH^N^ C, 56.54; H, 6.81; N, 36.65.
Found: C, 56.16; H, 6.98; N, 36.47.
2-Pyridinalhydrazone, 2-PAH. To 3.0 ml (0.09 mole) of
hydrazine in an ice bath, 7.2 ml (0.06 mole) of 2-pyridine-
caroxaldehyde (2-PCA) was added dropwise with stirring. The
golden yellow solution was "dried" in vacuo over P^O^q leav¬
ing a golden yellow oil. B. p., 181°C.
Anal. : Caled, for C^HyN^: Mol. wt., 121 d.
Found: Mol. wt.,121 d (46%).
2-Pyridinalphenylimine, 2-PPI. To a solution of 9.1 ml
(0.1 mole) aniline in 10 ml hexane was added dropwise with
stirring 10.7 ml (0.1 mole) of 2-PCA. The bright yellow
powdery precipitate isolated was recrystallized from hot
hexane and dried in vacuo over P^O-^q. 1% P- . 40°C.
Anal.: Caled, for Cp2H10^2: ^°1- wt• » 182 d.
Found: Mol. wt., 182 d (14%).

16
2-Pyridinal-p-rnethy lpheny 1 imine , 2-PMI. A solution of
10.7155 g (0.1 mole) of p-aminotoluene dissolved in 50 ml
hexane was prepared by heating on a steam bath, and 10.7 ml
(0.1 mole) of 2-PCA was added slowly with stirring. This
yellow solution was allowed to cool to ambient temperature.
The yellow needle-like crystals which precipitated were
recrystallized from hot hexane and dried in vacuo over P^Opg
M. p., 55°C.
Ana 1. : Caled, for Cp3H]_2N2: Mol. wt. , 196 d.
Found: Mol. wt., 196 d (69%).
2-Pyridinal-p-aminophenylimine , 2.-PAI. A solution of
10.8982 g (0.1 mole) of p-phenylenediamine dissolved in
100 ml hot deionized water was prepared. One gram of sodium
hydrosulfite was added to this dark brown solution, and some
black particles were removed by filtration. To this beige
solution, 10.7 ml (0.1 mole) of 2-PCA was added dropwise
with vigorous stirring. The gold powdery precipitate prod¬
uced was washed with hot deionized water and dried in vacuo
over P^O-^q. M. p., 154°C.
Anal. : Caled, for Mol. wt. , 197 d.
Found: Mol. wt., 197 d (100%).
2-Pyridinal-p-chiorophenylimine, 2-PCI. With gentle
heating and stirring, 10.7 ml (0.1 mole) of 2-PCA was added
to 12.7613 g (0.1 mole) of p-chloroaniline, and the solution
was allowed to cool to ambient temperature. The resultant

17
pale yellow powder was recrystallized from hot hexane and
dried in vacuo over P^O-^q . M. p. , 65°C.
Anal. : Caled, for C-^Hg^Cl: Mol. wt. , 217 d.
Found: Mol. wt., 217 d (41%).
2-Pyridinal-p-cyanophenylimine, 2-PCPI. While heating
7.0 ml (0.06 mole) of 2-PCA on a steam bath, 6.0353 g (0.05
mole) of 4-aminobenzonitrile was added slowly. The resultant
yellow solution v.7as heated for one hour, and then extracted
six times with 100 ml boiling hexane. The hexane extracts
were evaporated and the resultant yellow powdery product
was dried in vacuo over P^O-^q. M. p. , 109°C.
Anal. : Caled, for C]_2H9N3 : Mol. wt. , 207 d.
Found: Mol. wt. , 207 d (517.).
2-Pyri din a. 1-p-nitropheny limine , 2-PNI. While heating a
slurry of hot hexane and 13.8017 g (0.1 mole) of p-nitro-
aniline on a steam bath, 10.7 ml (0.1 mole) of 2-PCA was
added. The amount of solid appeared to increase and became
more flocculent. This mixture was heated subsequently for
one hour, washed four times with hexane, and dried in vacuo
over P^O-^q. M. p., 122°C.
Anal. : Caled, for Mol. wt. , 227 d.
Found: Mol. wt., 227 d (9%).
Tris(biacetyldihydrazone)cobalt(II) tetrafluoroborate,
[ Co (BdH) ^ ] (BF¿l) 9 . To a solution of 1.8247 g (0.01 mole) of

18
Co(NC>3)2 in 15 ml deionized water, 3.4175 g (0.03 mole) of
BdH was added with stirring. The resultant deep brown solu¬
tion was filtered and 20 g NaBF^ was added with stirring.
The light brown powder that separated was filtered, washed
with deionized water, and dried in vacuo over P40-^q. M. p.,
246°C dec.
Anal. : Caled, for [ Co (CZ)H1QN^) 3 ] (BF^) 2 : C, 25.06; H, 5.22;
N, 29.25.
Found: C, 25.23; H, 5.29; N, 29.36.
Tris(biacetyldihydrazone)iron(II) iodide dihyrate,
[Fe(BdH)I2•2H20. Excess BdH was added with stirring to a
solution of 3.9219 g (0.01 mole) of Fe(NH^)2(SO^)9•6H20 in
15 ml deoxygenated deionized water. This deep red-brown
solution was filtered, and 45 g Nal was added to the 45 ml
of solution. The red-brown powdery product was washed with
deionized water and dried in vacuo over P^O-^q- M. p. ,
205°C dec.
Anal.: Caled, for Fe(C4H1QN4)3I2â– 2H20: C, 20.91; H, 4.94;
N, 24.46.
Found: C, 20.61; H, 4.46; N, 24.20.
Tris(glyoxaldihydrazone)cobalt(II) iodide, [Co(GdH)^]I2.
All reagents, solvents and equipment were deoxygenated under
nitrogen for twenty-four hours, and the preparation was also
carried out under nitrogen. A solution of 3.6751 g (0.02
mole) of Co(N03)2 in 10 ml deionized water was made, and a

19
white slurry of 5.1116 g (0.06 mole) of GdH in 15 ml deionized
water was added with stirring. The resultant brown slurry
was allowed to stir thirty seconds and then filtered. To
the brown solution, a solution of 25 g of Nal in 25 ml de¬
ionized water was added with stirring. Thirty seconds later
a black powder was filtered off and dried in vacuo over P^O-^q.
M. p., 165°C dec.
Anal.: Caled, for Co(C2H6N4)3I2: C, 12.62; H, 3.15; N, 29.44.
Found: C, 12.29; H, 3.08; N, 28.71.
Tris(glyoxaldihydrazone)iron(II) iodide, [Fe(GdH)q]I?.
A solution of 1.3675 g (0.01 mole) of FeC^ in 25 ml deionized
deoxygenated water was made under nitrogen. To this solution
2.5817 g (0.03 mole) of GdH was added with stirring. The
resultant intense red-brown solution was filtered, and 25 g
of Nal was added with stirring. The red powdery product was
filtered, washed with deionized water, and dried in vacuo
over M. p., 202°C dec.
Anal.: Caled, for Fe(C2H6N4)3I2: C, 12.68; H, 3.17; N, 29.60.
Found: C, 12.79; H, 3.22; N, 29.80.
Tris(biacetylbismethylimine)cobalt(II) iodide,
[ Co (BMI) 3 19 . This complex was prepared following the method
of Clevenger (5) being extremely careful to deoxygenate all
equipment and reagents with nitrogen. M. p., 220°C dec.
Anal.: Caled, for Co(C6H12N3)3I2: C, 33.30; H, 5.55; N, 12.95.
Found: C, 33.13; H, 5.56; N, 12.80.

20
Tris(biacetylbismethylimine)iron(II) tetrafluoroborate,
[Fe(BMI)(BF4)2â–  This compound was prepared by a method
analogous to that used for preparing Co(BMI)^I7 except that
FeCl2 and Nal were used. M. p., 268°C dec.
Anal. : Caled, for Fe 3 (BF^) 9 : C, 38.28; H, 6.37;
N, 14.86.
Found: C, 38.32; H, 6.44; N, 14.91.
Bis(2,6-pyridinediacetyldihydrazone)cobalt(II) iodide,
[Co(2,6-PdAdH)9]I9• A solution of 5.3108 g (0.03 mole) of
Co(N03)2 in 50 ml of deionized water was prepared, and
12.1071 g (0.06 mole) of 2,6-PdAdH was added with stirring.
The resultant brown solution was filtered to remove excess
ligand, and 50 g of Nal was added with stirring. The black
powdery product was then removed by filtering, washed with
cold water, and dried in vacuo over P^O^q. M. p., 270°C.
Anal. : Caled, for Co (CpH^N^) : C, 31.09; H, 3.75;
N, 20.15.
Found: C, 31.25; H, 3.76; N, 20.23.
Bis(2,6-pyridinediacetyldihydrazone)iron(II) iodide di¬
hydrate , [Fe (2,6-PdAdH) 9 ] I9 ' 211^0 • This compound was prepared
by a method analogous to that used for preparing
[Co(2,6-PdAdH)9]I9 except that Fe(NK^)9(SO^)2’6H2O was used.
M. p., 300°C.
Anal.: Caled, for Fe(CgH13N5)^2‘2H9°: C, 29.81; H, 4.13;
N, 19.36.
Found: C, 30.30; H, 4.10; N, 19.46.

21
Tris(2-pyridinalhydrazone)cobalt(II) iodide,
[Co(2-PAH)2 Io 1• A solution of 3.0 ml (0.06 mole) of hydra¬
zine in 10 ml of deionized water was prepared in an ice bath,
and 7.2 ml (0.06 mole) of 2-PCA was added dropwise with
stirring. To this straw yellow solution, 3.6592 g (0.02
mole) of CoiNO^^ dissolved in 10 ml deionized water was
added dropwise with stirring. The resultant deep brown
solution was then heated with stirring to 100°C, allowed to
cool to ambient temperature, and placed in a refrigerator
for 6 hours. A beaker containing the solution was then
scratched with a freshly broken glass rod and allowed to
sit 30 minutes. The resultant tan powder was washed with
cold deionized water and dried in vacuo over P^O-^q . M. p.,
220°C.
Anal. : Caled, for Co (Cgl^N-j) 312 : C, 31.96; H, 3.11; N, 18.65.
Found: C, 31.26; H, 3.11; N, 18.09.
Tris (2-pyridinalphenyl inline) cobalt (I I) tetra fluorobo rate ,
[Co(2-PPI)3](BF/)2. A solution of 1.8175 g (0.01 mole) of
Co(N03)9 dissolved in 30 ml of 50% ethanol and 50% deionized
water was prepared. To this solution, 5.7518 g (0.03 mole)
of 2-PPI was added with stirring. After adding 100 ml de¬
ionized water, this red-brown mixture was allowed to stir
for 24 hours. The aqueous layer was separated from a red oil,
and 6 g of NaBF^ was added with stirring to the aqueous mix¬
ture. The resultant red-brown oily precipitate and aqueous
mixture was heated with stirring, and a beige powder was

22
produced. It was washed with deionized water and dried in
vacuo over P^O^q. M. p., 280°C dec.
Anal. : Caled, for Co (C12H1C)N2) 3 (BF4) 2 : C, 55.49; H, 3.85;
N, 10.79.
Found: C, 55.33; H, 3.92; N, 10.88.
Tris(2-pyridinal-p-methylphenylimine)cobalt(II) tetra-
fluorobórate, [Co(2-PMI)3 ](BF^)2. This compound was prepared
by a method analogous to that used for [Co ( 2-PPI) 3 ] (BF^) 2 .
M. p., 283°C.
Anal.: Caled, for Co(C13H12N2)3(BF4)2: C, 57.03; H, 4.39;
N, 10.24.
Found: C, 56.76; H, 4.45; N, 10.33.
Tris(2-pyridinal-p-aminophenylimine)cobalt(II) tetra-
f luoroborate, [Co(2-PAI)3] (BF<,)2- All reagents, solvents,
and equipment were deoxygenated and all procedures were carried
out under nitrogen in order to obtain this compound. A solu¬
tion of 1.8271 g (0.01 mole) of Co(N03)2 in 25 ml deionized
water was prepared, and 6.1028 g (0.03 mole) of 2-PAI was
added and the mixture was allowed to stir for 15 minutes.
This red-brown mixture was then filtered, and 5 g of NaBF^
was added with stirring to the aqueous solution. The result¬
ant red-orange powdery solid was washed with deionized water
and dried in vacuo over P^O^q- P- » 170°C dec.
Anal. : Caled, for Co (C^2H-^-^N3) 3 (BF^) 2 : C, 52.46; H, 4.01;
N, 15.30.
Found: C, 52.18; H, 4.11; N, 15.22.

23
Tris(2-pyridinal-p-ch1orophenylimine)cobalt(II) tetra-
fluoroborate, [Co(2-PCI)3](BF7)q. To a solution of 1.8294 g
(0.01 mole) of CoCNO^)^ in 15 ml of deionized water, 6.5026 g
(0.03 mole) of 2-PCI was added and the mixture was allowed
to stir for 24 hours. The resultant orange powder was dis¬
solved in 50 ml of hot ethanol, and 10 g of NaBF^ was added
with stirring. This red-brown solution was allowed to sit
at ambient temperature for 72 hours. A red-brown solid
consisting of needle-like crystals (soluble in dimethylsulf-
oxide, DMSO) and round globs (insoluble in DMSO) was obtained.
Enough DMSO was added to just dissolve the crystals, and
the round globs were removed by filtering. The orange DMSO
solution was warmed to about 75°C, and deionized water added
until an orange powder just began to precipitate. This solu¬
tion was placed in a refrigerator for 6 hours. The orange
product was dried in vacuo over P/40]_g- M. p. , 245°C.
Anal.: Caled, for Co(C12HgN2Cl)3(BF^)2: C, 48.98; H, 3.07;
N, 9.53.
Found: C, 48.72; H, 3.16; N, 9.39.
The cobalt(II) complexes of 2-PCPI and 2-PNI could not
be prepared. Repeated attempts to prepare these complexes
in deionized water, ethanol, methanol, tetrahydrofuran,
carbon tetrachloride, chloroform, and dimethylsulfoxide met
with failure. The cobalt complexes of GdH, BdH, BMI, and
2,6-PdAdH are not stable in air. However, the remaining
cobalt complexes and the iron complexes are considerably
more stable.

24
Elemental Analysis
All elemental analyses were performed by Atlantic Micro¬
lab, Inc., Atlanta, Georgia.
Mass Spectrum Analysis
Parent peaks for several ligands were obtained, in lieu
of elemental analysis, using a voltage of 70 eV, on an AEI
MS 30 mass spectrometer equipped with a DS 30 data system.
Spectrometers
A Varian Associates XL-100 NMR spectrometer was used to
obtain the NMR spectra.
Infrared spectra were obtained on a Perkin-Elmer 137
sodium chloride spectrophotometer, and were calibrated with
polystyrene. Pressed potassium bromide samples were used to
obtain infrared spectra.
Electron paramagnetic resonance spectra were obtained
at 80 K using a Varian Associates E-3 EPR spectrometer,
V-4502-14.
The magnetic susceptibilities were determined by the
Gouy method. The equipment has been described previously
(5). The magnetic field was calibrated using mercury(II)
tetrathiocyanatocobaltate(II) (25). The maximum field strength
at a current of 4 A was 6750 ± 80 gauss.

25
Solution susceptibilities were determined using a
Varian Associates A-60 NMR spectrometer, and the method of
Loliger and Scheffold (27). Spectral grade methanol was
used to prepare solutions of 2 x 10 ^ mole/ml.
Theoretical spin density calculations were obtained from
a INDO/1 program modified to include calculations on molecules
containing atoms from the first transition series. The ori¬
ginal INDO model is not changed from that given by Pople and
Beveridge (34). This modified program was written and devel¬
oped by Michael Zerner and Alan Bacon at the Chemistry Depart¬
ment, University of Guelph, Ontario.

RESULTS
The positions of the infrared absorptions and their
corresponding assignments for characterization of the compounds
prepared for this investigation are listed in Table I. These
assignments are based on the ranges given in Bassler and
Silverstein (2). In addition, similar v(C=N) assignments
were made by Fisher and Stoufer (11) for Co(GdH)3Br2, Co(BdH)3I2>
Fe(GdH)3I2, and Fe(BdH)3I2.
Average room temperature magnetic moments, Beff, cal¬
culated from susceptibility data, and g values calculated
from ESR data are reported in Table II. These are average
moments determined at field strengths of 5649 Hz and 6750 Hz
to insure there was no field strength dependence and thus
no ferromagnetic or antiferromagnetic interactions were
present. Pascals constants for diamagnetic corrections to
the measured susceptibilities were obtained from Lewis and
Wilkins (25) and from Selwood (38). It would have been
preferable to determine Pe££ in solution for all the compounds,
however, only Co(2,6-PdAdH)2I2 was amenable to this technique,
since the other cobalt(II) compounds decompose too rapidly
in solution. Solution measurements for these would yield
values which would be the average for each compound
and its decomposition products. As reported previously (5,10)
26

27
Table I
Selected Infrared Absorptions and Assignments for
the Co(II) and Fe(II) Compounds Studied
Compounds
Co(GdH)3I2b
Co(BdH)3(BF4)2b
Co(2,6-P dAdH)019
Absorption (cm b)
Assignment3
3200 s
vs(N-H)
3100s(sh)
uas(N-H)
3000 s
vs(C-H)
1590 m(sh)
v) ( C=N)
1570 s
v(C=N)
1410 in
6(C-H)
3510 s
vs(N-H)
3420 s
vas(M-H)
3220 s
vas(C-H)
3000 w
vs(C-H)
1650 m
6(N-H)
1610 s
v(C=N)
1440 m
{as(C-H)
1375 m
3400 s
vs(N-H)
3300 s
vas(N-H>
3180 s
vas(ring C-H)
3100 s
vg(ring C-H)
3000 m(sh)
vas(methyl C-H)
2850 m
vs(methyl C-H)

28
Compounds
Co(2,6-PdAdH)2I2
Co(BMI)3I9
Fe(GdH)3I2b
Fe(BdH)3I2-2H2Ob
Table I - continued
Absorption (cm b)
1620 m
1580 s
1440 m
1375 s
800 m
2900 w(sh)
2850 w
1590 m
1420 m
1375 m
3225 s
3100 s
3000s
1580 s
1540 s
1416 w
3350 m
3250 m
3150 m
Assignment^1
ó(N-H)
v(C=N)
áas ('rnetbyl C_H)
(methyl C-H)
6(ring C-H)
vas
vg(C-H)
v(C=N)
6as^C-H)
5S(C-H)
vs(N-H)
vas(N-H)
vs(C-H)
v(C=N)
v(C=N)
Sas(C'H)
vs(N-H)
vas(N-H)
"as
vg(C-H)
2900 w

29
Table I - continued
Compounds
Fe(BdH)3I2'2H20
Fe(2,6-PdAdH)2I2
Fe(BMI)3(BF4)2
Absorption (cm )
1620 m
1590 m
1440 m
1375 m
3400 m
3300 m
3180 m
3100 m
3000 w(sh)
2850 w
1620 m
1590 m
1440 m
1375 m
800 in
2950 w
2850 m
1610 w
1430 s
1375 (s)
3.
Assignment
6(N-H)
v(C=N)
6as(C-H)
6S(C-H)
vs(N-H)
vas(N-H)
vas(rin§ C~H)
VS(ring C-H)
va£, (methyl C-H)
vg(methyl C-H)
6(N-H)
v(C=N)
ó (methyl C-H)
6s(methyl C-H)
6(ring C-H)
yas(C'H)
vs(C-H)
v(C=N)
6as(C-H>
6 s(C-H)
Based on ranges given in Bassler and Silverstein (2).
^Similar v(C=N) assignments were made by Fisher and Stoufer (11).

30
Table II
Magnetic Moments and g Values for the
Co (II) and Fe(II) Compounds Studied
Compounds3
4eff (3°° K- in BM)
g (80 K)
Co(GdH)3I9
4.23
2.15
Co(BdH)3I?
4.65b
2.06
Co(2,6-PdAdH)9I2
2.28b,c
2.09
Co(BMI)3I?
3.30d
2.13
Fe(GdH)3I9e
0.00
-
Fe(BdH)3I2-2H90e
0.61
-
Fe(2,6-PdAdH)9I9e
0.57
-
Fe(BMI)3(BF4)2e
0.80
-
Pascals constants for diamagnetic corrections to the measured
susceptibilities were obtained from Lewis and Wilkins (25)
and Selwood (38).
bThe reported value is 4.17 BM (40).
determined by NMR in solution (27).
dThe reported value is 2.91 BM (41).
0
No EPR signal was observed.

31
these anomalous moments for all the cobalt(II) compounds
span the range between low-spin, 1.7 BM, and high-spin, 5.2
BM, values. It is noteworthy, however, that the 2,6-PdAdH
complex has a sufficiently low, 2.28 BM, that it may be
considered to be a low-spin complex. The observed nonzero
moments for the "diamagnetic" iron(II) compounds are a result
of temperature independent paramagnetism.
The EPR. spectra for these cobalt(II) compounds consisted
of a very broad signal centered near a g value of 4 attributable
to the slightly populated E^(^Tr> ) excited term, and, superim-
posed on it, a relatively narrow signal centered upon a g value
in the range 2.06 - 2.15, very close to the spin-only value of
2
2.00229 and the isotropic value of 2 expected for the E level.
The iron(II) compounds exhibited no ESR signal — as expected.
The EPR data presented above and the NMR data which follow
are correlated in the discussion (vide infra). The isotropic
shifts, Av, for the cobalt(II) compounds are reported in
Table III together with the relative shifts and the corre¬
sponding half-height widths. These relative isotropic shifts,
Avrel’ are normalized with respect to the largest observed
shift, which in each case happens to be the proton-con¬
taining moiety bound to the hydrazone or imine nitrogen, As
stated previously the GdH, BdH, and BMI complexes decompose
rather quickly in solution. Therefore, for each it was
necessary to make all instrumental adjustments and run a
trial spectrum on one sample and run the final spectrum
quickly on a freshly prepared sample. Within ten minutes
resonances attributable to the decomposition products
began to appear. Thus it was possible to make unambiguous

32
Table III
Observed and Relative Isotropic Shifts
and Half-Height Peak Widths for Cobalt(II) Compounds
Ligand
Av (ppm)
Avrel (P?m)a
Width (Hz)
GdH
C-H
2.29
- 0.022
60
N-H
- 103.66
1.00
500
BdH
C-H
- 2.72
0.031
30
N-H
- 86.69
1.00
150
BMI
C-C-H
- 10.14
0.236
60
N-C-H
- 42.90
1.00
150
2,6-PdAdH
C-H
8.97
- 0.144
110
N-H
- 61.91
1.00
750
ha
- 12.09
0.195
60
hb
- 6.06
0.098
80
Relative isotropic shifts are normalized with respect
to the largest observed shift.

33
assignments based on these observations, the relative neak
intensities, and the fact that all the protons bound to the
hydrazone or imine group would be expected to shift the same
direction. Furthermore, they exhibit the same relative
magnitude (i.e., largest shifts observed in each spectrum)
and all are the broadest peaks in the spectra, which broaden¬
ing is due to their proximity to the nitrogen quadrupole.
There was no decomposition problem involved in the
assignments for 2,6-PdAdH. However, the and Hg isotropic
shifts as assigned by relative intensities are reversed
compared to the expected shifts (vide infra). Consequently,
the assignments for and Hg are based on the following
syllogism. Because the electron-proton dipolar relaxation
mechanism often dominates the line widths in paramagnetic
complexes, the relative line widths for nonequivalent protons
will be determined primarily by their relative values of
R 6, where R is the distance from the central metal ion to
the proton (4,24). Since RT7 > Ru , .
F \ Hr (see the structures m
B
the introduction) , R,,^ < R ^ and the line width for H. should
hA HB A
be less than that for H,-,.
D

DISCUSSION
Theoretically it is predicted that proton contact shift
spectra for complexes with s = % will not be observable
because T^e is usually too long (42). And — as noted in the
introduction — no spectra have been reported for low-spin,
octahedrallv coordinated d^ ions. As inferred from the
introduction, the purpose of this investigation was, in part,
to determine if proton contact shift spectra are observable
for anomalous cobalt(II) complexes near the cross-over point,
and to determine if the spectra are attributable to the E ( G)
term. If so, this would mean T-^£ is sufficiently lowered by
low-lying excited states (42).
There are several approaches which should be considered
in arriving at an interpretation of the isotropic shifts of
these complexes. (i) The site symmetry of these complexes
may be taken to be octahedral. Accordingly, if the concen¬
tration of the ^T^(^F) species is small, and if the ^Eg--^T-^
interconversion provides a mechanism whereby T-^e is sufficient-
2
ly lowered, then a contact shift should arise from the E
g
term. In this case there would be no spin-orbit coupling
9
effects from "E and, because E and A terms are isotropic,
5
there would be no dipolar shift complications to consider (15).
(ii) Under octahedral symmetry the contact shift arises from
34

35
the term. This is the origin of all reported cobalt(II)
contact shift spectra (42). In this case there are spin-orbit
coupling and dipolar shift complications. (iii) Both (i)
and (ii) contribute to the observed spectra. In this case,
2 A-
if the E -- T-, interconversion time is large on the NMR time
o
scale, two superimposed spectra should result. This was not
observed for these complexes (Table III). (iv) If (i) and
2 A
(ii) contribute and the E -- T, interconversion time is
o
small on the NMR time scale, time-averaged spectra should
result. In this case spin-orbit coupling (44) and dipolar
or pseudocontact shifts from the ^T-^ term are involved. But,
approaches (i) - (iv) are valid explanations only under octa¬
hedral symmetry.
Strictly speaking six-coordinate complexes containing
symmetric bidentate ligands are inherently of symmetry;
consequently approaches (i) - (iv) are not valid considerations.
Under symmetry ^Eg(O^) -* ^Eg(D^) and ^T-^(O^) -* ^A(D^) +
^E(D^). Since all the term levels in symmetry are isotropic
(A and E), the observed isotropic shifts should be due en¬
tirely to contact interaction (15,29,43). Therefore pseudo¬
contact contributions either should be absent or negligible.
The nonideal approach commonly used to interpret contact
shift data is to treat a single ligand molecule as a free
radical cation or anion depending on the delocalization mech¬
anism involved. That is, the delocalization of unpaired spin
is considered to result from L -> M or M -> L charge transfer
(23,36). In the discussion which follows, (-t) and (l)

36
represent an electron whose magnetic moment is aligned par¬
allel and antiparallel, respectively, to the external magnetic
field.
Figure 3 illustrates the spin transfer mechanisms possible
(23). The ligands contained in these complexes are all dia¬
magnetic and thus the ligand MO's, 4>L, either are doubly
occupied or are unoccupied. L -> M charge transfer is
illustrated in (a) - (c). If the metal ion has a singly
occupied d orbital of a symmetry (a) is the only transfer
possible. The spin is transferred from a ligand a orbital
(^) to the occupied metal d( and paired with the spin in the
d^ orbital. Since the metal ion unpaired spin will be par¬
allel (t) to the external field, this leaves net (t) spin
on the ligand. The metal ion can contain a singly occupied
d71 orbital with a low-lying da orbital, (b) and (c). Un¬
paired spin ( + ) and (;) can be transferred into the da from
the However, correlation favors transfer of (+) spin
(Hund's rule of maximum multiplicity) leaving net (4-) spin
in (32). M ->• L charge transfer is illustrated in Figure
3 (d) - (f) . For a diamagnetic ligand a vacant is re_
quired. Vacant o MO's are too high in energy to be im-
portant, so only vacant tt MO's need be considered. If the
7T
metal d orbital is singly occupied, (d) is the only trans¬
fer possible and net (+) spin is delocalized onto the ligand.
However, if the unpaired spin occupies a d° orbital, and
the d' orbitals are doubly occupied, (e) and (f) spin trans¬
fer is possible. In this case, correlation favors transfer
of (I) spin from metal to ligand.

37
da
du
t
t + , a
±
(a)
Cb)
(c)
v
+ , 7T " ' _
d
L
'L
<í>
(d)
(e)
(f)
Figure 3. Possible spin transfer mechanisms
r1 =3

38
A theoretical analysis of the [Co(GDH)^ ] spectrum was
attempted based on a complete, many-electron calculation for
the entire complex. Such a calculation was performed with
O
the complex considered to be (i) high-spin with sp and with
3
sp hybridized amine nitrogen orbitals and (ii) low-
2 . 3
spin with sp" and with sp hybridized amine nitrogen orbitals.
These calculations were accomplished using a modified INDO
program. Since X-ray crystallography data were not available,
the bond lengths used in these calculations are the average
bond lengths given by March (30). This program is capable
of handling the minimum basis set (ninety-nine atomic orbitals)
of the complex, including metal penultimate d and ultimate s
and £ orbitals and all of the ligand valence orbitals.
The results of the calculations together with the ex¬
perimental results are listed in Table IV. An inspection of
the data will reveal that neither the absolute magnitudes
nor the signs of the theoretical spin densities agree with
the experimentally determined values and signs. Indeed, even
the relative magnitudes do not agree. That is, although the
hydrazone protons are found to have the largest shifts and
therefore the largest spin densities, the INDO calculations
predict the protons attached to the carbons to exhibit the
largest shift. One factor causing this disagreement may be
that these experimentally determined densities were calculated
using Equation (2) which is applicable to complexes having
normal Curie behavior, a condition which does not exist for
this complex (vide infra). Another factor may be that,

INDO
Calculated and
Table IV
Experimental Proton
Spin Densities for
[Co(GdH)
++]
3 J
GdH
position
pINDO
°exptl
low spin
high
spin
low spin
high
spin
caled relative
(xlO-6)
caled
(xlO-6)
relative
caled relative
(xlO-6)
caled relative
(xl0_ 6)
sp2
Hydrazone*
CH
- 2108 1.00
- 4698
1.00
- 58 0.0222
- 35
0.0224
NH
- 306 0.145
- 1105
0.235
2607 1.00
1563
1.00
sp3
Hydrazone
CH
- 2589 1.00
- 5036
1.00
- 58 0.0222
- 35
0.0224
NH
493 0.190
425
0.0943
2607 1.00
1563
1.00
INDO calculations were made for this complex with the hydrazone nitrogen considered to be
sp2 and sp3 hybridized. Robinson and Busch (35) provide evidence that sp^ hybridization
is preferred.
\o

40
although this modified INDO method is capable of handling this
basis set (vide supra), the amount of unpaired spin density
calculated on the protons is sufficiently small (Table IV)
that it can only be considered to be "noise." This is due to
the fact that the tolerance in this program cannot be set
sufficiently low to insure convergence, and thus calculate
meaningful densities for a total of one unpaired electron
(low-spin) or a total of three unpaired electrons (high-spin)
in the complex over ninety-nine basis orbitals. Consequently,
no meaningful importance should be attached to these calculated
densities.
Since INDO calculations did not produce meaningful
results, the common (but nonideal) approach of considering
the ligand to be a radical cation or anion was used even
for this simple complex. However, in this work the entire
valence set of the ligand was used rather than just the it
symmetry basis functions as is frequently done (see reference
36 and references therein). The metal orbitals are completely
neglected in forming the molecular orbitals for the calcula¬
tions; that is, the metal ion has been considered to be
either an electron source or sink. In addition, it is general¬
ly accepted that ligands containing an a-diimine linkage
serve not only as a donors but as iT-acids as well (dir — ptr
back bonding with the metal ion) . That is , there may be M -> L
and L -> M spin transfer simultaneously. Obviously it is an
oversimplification to consider the ligand to be either a
cation or an anion under such circumstances. Furthermore,

41
the effects of one ligand on another in the complex has been
neglected. And — the bond lengths of the neutral ligand
molecule were used (30). This certainly is not the ideal
approach, but there is little choice since there is no rea¬
sonable way to estimate the magnitude of the change of bond
length.
Table V is a presentation of the INDO spin densities for
the ligand radicals and the corresponding spin densities
calculated from the contact shifts and the relative spin
density of each. An inspection of these densities leads to
the conclusion that, for all of the ligand radicals,the agree¬
ment between the relative magnitudes of the calculated spin
densities and, in the case of the cations, the experimental
densities is, in general, fairly good. The BMI+^ calculation
produces the best agreement in relative magnitude: pre^(IND0)
= 0.166 and 1.00 for CCH^ protons and NCH^ protons, respec¬
tively, and pre-^(exptl) = 0.236 and 1.00 for CCH^ protons and
NCH^ protons, respectively. In comparison, Scarlett et al.
(36) report good agreement for the pyridine N-oxide cation:
Prei(INDO) = - 1.05, 0.41 and - 1.00 for 2-H, 3-H and 4-H,
respectively, and pre^(exptl) = - 0.81, 0.45 and - 1.00 for
2-H, 3-H and 4-H, respectively. Only the GdH radical cation
calculations with sp hybridized amine nitrogen orbitals
agree in both sign and relative magnitude (vide infra) with
the experimental densities. This may be taken as an indication
of L -> M charge transfer mechanism of spin delocalization.
In contrast, the calculated spin densities for both the BdH
2
radical cation and anion with sp hybridized amine nitrogen

Table V
INDO Calculated and Experimental Proton Spin Densities
for Cationic and Anionic Radicals
Radical Ligand PIND0
and Position
Pexptl
low
spin
high spin
caled (xlO 6) relative3 calcd^(xlO"
Ó \ -i , • 3.
) relative
calcd^(xlO relative
GdH 1
CH
- 2491
1.00
- 83
0.0318
- 50
0.0320
NH2 (sp2)C
- 1183
0.475
2607
1.00
1563
1.00
GdH"1
CH
1070
0.312
- 83
0.0318
- 50
0.0320
NH2 (sp3)
3425
1.00
2607
1.00
1563
1.00
GdH+1
CH
- 1579
0.118
- 83
0.0318
- 50
0.0320
NH2 (sp2)
13417
1.00
2607
1.00
1563
1.00

Table V
continued
Radical Ligand pTKmn
and Position iINUU
pexpt1
low
spin
high spin
GdH+1
caled (xlO ^)
relative3
calcd^(xlO
6) relative3
calcd^(xlO
6) relative3
CH
1917
1.00
- 83
0.0318
- 50
0.0320
NH2 (sp3)
1215
0.634
2607
1.00
1563
1.00
BdH” 1
CH3
410
0.331
98
0.0426
59
0.0427
NH2 (sp2)
1239
1.00
2302
1.00
1381
1.00
BdH"1
ch3
1467
1.00
98
0.0426
59
0.0427
NH2 (sp3)
1295
0.882
2302
1.00
1381
1.00
BdH+1
CH3
NH2 (sp2)
1272
10950
0.116
1.00
98
2302
0.0426
1.00
59
1381
0.0427
1.00

Table V - continued
Radical Ligand pTNnn
and Position
Pexptl
low spin high spin
caled (xlO 6) relative3 calcd^CxlO relative3 calcd^(xlO relative3
CH3 ,
604
1.00
98
0.0426
59
0.0427
NH2 (sp3)
- 578
0.956
2302
1.00
1381
1.00
Bill' 1
cch3
332
0.0570
257
0.236
155
0.237
nch3
5822
1.00
1089
1.00
654
1.00
BMI+1
cch3
- 836
0.166
257
0.236
155
0.237
nch3
5049
1.00
1089
1.00
654
1.00

Table V - continued
Radical Ligand
and Position
pINDO
Pexptl
low
spin
high spin
caled (xlO
6) relative3
calcd^
(xlO-
6) relative3 calcd^(xl0
6) relative3
2,6-PdAdH-1
CH3 ,
- 3347
0.125
- 232
0.145
139
0.145
NH2 (spZ)
- 26872
1.00
1602
1.00
961
1.00
ha
- 20935
0.779
313
0.195
188
0.196
hb
12983
0.483
157
0.0980
94
0.0978
2,6-PdAdH"1
CH3
NH2 (sp3)
- 4303
- 17091
0.251
1.00
- 232
1602
0.145
1.00
139
961
0.145
1.00
ha
- 4770
0.279
313
0.195
188
0.196
hb
189
0.011
157
0.0980
94
0.0978
2,6-PdAdH+1
CH3
NH2 (sp2)
- 3604
- 43423
0.0830
1.00
- 232
1602
0.145
1.00
139
961
0.145
1.00
ha
- 8300
0.191
313
0.195
188
0.196
«B
10680
0.246
157
0.0980
94
0.0978

Table V - continued
Radical Ligand
and Position
pIND0
pe
xptl
low spin
high spin
caled
(xlO 6) relative3
i jb
caled
(xlO 6) relative
a ,b
caled i
(xlO 6) relative3
2,6-PdAdH+1
CH3 .
4564
0.532
- 232
0.145
- 139
0.145
NH2 (spJ)
8571
1.00
1602
1.00
961
1.00
ha
3398
0.396
313
0.195
188
0.196
hb
6210
0.725
157
0.0980
94
0.0978
Relative spin densities are normalized with respect to the largest shift or spin density
b
Spin densities are calculated from experimentally observed contact shifts using Equation
(2) and Ay = Kp /2s, where A^ is the hyperfine coupling constant, p^. is the unpaired spin
density at the hydrogen nucleus, 2s is the total number of unpaired spins, and K is the
hydrogen coupling constant for a whole unpaired spin (1420 MHz, 507 G)(16,19,22,23).
'Calculations were made
hybridized.
assuming, as indicated, the hydrazone group to be sp“ or sp'

47
orbitals agree in sign and relative magnitude with the experi¬
mentally calculated densities. Thus the charge transfer mech¬
anism may be either L + M and/or M -*â–  L. Calculated snin
densities for the BMI radical anion agree in both sign and
relative magnitude of spin densities with the experimental
densities. This may indicate M ** L charge transfer mechanism
of spin delocalization. Although the radical anion calculated
3
densities for 2,6-PdAdH with sp hybridized hydrazone nitrogen
agree with the relative magnitude of the experimental densities,
the signs of the densities do not agree, which means there is
no agreement at all (vide infra).
The experimentally determined densities listed in Table
V, calculated using Equation (2), apply to the special case
of normal Curie behavior and are derived from (12,18)
AH - Ah / g^j Bjq (3)
where <^S \ is the time-averaged value of the z component of
the electron spin. So, the problem of calculation reduces
to an evaluation of <\Sz^>. For normal Curie behavior (12,18)
Golding (13) evaluated \Sz^> for complexes of octahedral site
symmetry. This evaluation included spin-orbit coupling and
magnetic field interactions
I(2J + 1) I (2d +
J
1)
e[-E/kT]
(5)
where E is expressed as a power series in H and

48
(7)
where r is the spin-orbit coupling constant of the transition
metal ion. Golding tabulated values for the parameters y, v,
g, A, and B required to evaluate for octahedrally coor¬
dinated metal ions of dn configuration.
Under octahedral symmetry and by including spin-orbit
/ o
coupling, the x, and ZE terms split into 2U' + E" + E'
§
and U' terms respectively. At ambient temperature both the
2 4
ground term U' ( E ) and first excited term E' ('r-^)j.
are occupied (44). Therefore the summation of (S )T in
Equation (5) must include these occupied terms. Substitution
into and evaluation of Equation (3) requires two hyperfine
coupling constants (12). Thus there are four unknowns A^,
Ag,, E, and £ and only one equation. By assuming the free
ion value of £ and varying E to produce a best fit, Equation
(3) has been used to calculate coupling constants for spin-
state equilibrium systems (13). However under the noncubic
symmetry of these cobalt(II) complexes the situation is
much more complex. In addition, the free ion value of £ is
certainly not appropriate. To be sure, a value less than the
free ion value — usually 757, of the free ion value — should
be used to account for covalency effects (44). And, under

49
more than two term levels may be occupied at ambient tem¬
perature adding further complications. Furthermore, mixing
of terms by spin-orbit coupling must also be considered.
Although matrix elements for spin-orbit coupling under octa¬
hedral symmetry have been calculated previously by Griffith
(14), those for symmetry have not been published. Con¬
sequently, it was not feasible to determine coupling constants
and thus spin densities from the observed contact shifts.
Fortunately relative contact shifts — and therefore relative
spin densities — can be determined using Equation (3) with¬
out evaluating . That is, there is no dependence on the
distribution over more than one term, as evidenced by the
close agreement of relative spin densities for high-spin
and low-spin values recorded in Table V.
The degeneracy of the t2g and e^ orbitals is lifted
under symmetry as shown in Figure 4 (33). The t2g r-
bonding orbitals transform as al and eTr under ; the e°
-L O
transforms as ec. It is seen that a high-spin or a low-spin
configuration is possible for seven d electrons under ;
thus, the 0^ and situations are qualitatively analogous.
That is, anomalous magnetic behavior can arise under either
symmetry. Considerations of the spin transfer mechanisms
represented in Figure 3 and the electron configurations
given in Figure 4 reveal (vide infra): (i) If the electron
configuration is either high-spin or low-spin, (+) spin can
be transferred onto the ligand by L -*â–  M or M -> L, respectively,
(ii) If the electron configuration is low-spin, (0 spin

50
Is { (e)‘
hs {(e)
°h
D
3
—• — e
a
Ot
D,
Is {(a1;
2°-
CD .
hs {(a^'
Figure 4. Possible 0^--D^ correlation diagrams
configurations for a ion under
(a1)2(e)1}
(a1)1(e)2}
2(e)4(e)1}
2(e)3(e)2}
and electron
symmetry

51
can be transferred onto the ligand by L -* M. (iii) If the
electron configuration is either high-spin or low-spin, both
(i) and (-t ) spin can be transferred onto the ligand by M L.
The ligand orbitals and metal ion orbitals can interact
and thereby transfer unpaired spin only if they are of
appropriate energy and of the same symmetry. The tt MO' s for
a symmetric bidentate ligand can be classified as (+) or (-)
according to their parity, Figure 5, which either remains
unchanged or is reversed upon rotation about the C2 symmetry
axis (33). The (-) MO's can interact only with the metal
e orbitals, while the (+) MO's can interact with either a-^ or
e metal ion orbitals depending on the choice of C0 axis for
the metal ion orbitals. Table VI is a presentation of the
INDO highest bonding and lowest antibonding tt MO's, their
symmetries, (+) or (-), and their energies. These are pre¬
sented for the ligand radicals which afford the best fit
with experimental densities.
The reader should apply the general discussion above
and refer to Figures 3 and 4 and Table VI for the discussion
of transfer mechanisms which follows. The calculations of
2
INDO spin densities for the radical cation with sp hybridized
amine nitrogen orbitals gave the best fit with the GdH com¬
plex. For this cation the unpaired spin is predicted to oc¬
cupy i¡whose symmetry upon rotation about the C2 axis is
(+) . Therefore it can interact through L -+ M spin transfer
with e(d) or a-^ (d) metal orbitals. Thus a low-spin configu¬
ration should result in (4) spin and a high-spin configuration
should result in (1) spin in •

52
«•
Figure 5
Illustration of (+) and (-) symmetries of selected
<}>£' s and selected metal orbitals upon rotation
about the C2 axis

53
Table VI
Selected INDO Molecular Orbital Symmetries
and Energies of Ligand Radicals
Ligand
Radical
Molecular
Orbital3
Occupation
Number
Energy
(Hartrees)
GdH+1 (sp2)
*16
2
- 0.627005
h7 (+)
1
- 0.573568
i—1
03
1
0
- 0.066018
BdH"1 (sp2)
^23 (+)
2
- 0.072167
^24
1
0.124016
b25 (a)b
0
0.513461
+1 o
BdH (sp*-)
i'22 i-)
2
- 0.587260
4* 2 3 (+)
1
- 0.552269
^24 ^
0
- 0.050940
^25 (+)
0
0.058413
BMl”1
\\>22 (a)b
2
- 0.086492
, vb
P23 (a)
2
- 0.083043
'^24 ^~^
1
0.107814
^25 ^
0
0.496251
a 7t MO' s are classified (+) or (-) depending on whether they
remain unchanged or reverse sign upon rotation about their
Cg axis (32). The MO's are numbered in order of increasing
energy.
^These
are
a molecular orbitals
and are not classified further.

54
The calculations of the INDO spin densities for the
2
radical cation and anion with sp hybridized amine nitrogen
orbitals give the best fit with the BdH complex. The anion
calculation predicts the unpaired spin to occupy 1^24» its
symmetry upon rotation about the C? axis is (-). Therefore
it can interact through M -> L spin transfer only with the
a-^ (d) metal orbital. Consequently, a low-spin configuration
should result in (1) spin and a high-spin configuration
should result in (t) spin in 24 • In addition, it may well
be that is energetically accessible. If it were acces¬
sible it could accept (+) spin and/or (1) spin from a ea(d)
orbital. The BdH cation calculation of INDO spin densities
predicts the unpaired spin to occupy ^23’ whose symmetry
upon rotation about the C2 axis is (+). Therefore, it can
interact through L -> M spin transfer with e(d) or a^Cd) metal
orbitals. Thus a low-spin configuration should result in
(;) spin and a high-spin configuration should result in (t)
spin in ^23-
The radical anion calculation of INDO spin densities
gives the best fit with the BMI complex. The unpaired spin
is predicted to occupy ^24 > whose symmetry upon rotation
about the C2 axis is (-). Therefore it can interact through
M -> L spin transfer only with the a-^(d) metal orbitals. Thus
a low-spin configuration should result in (+) spin and a high-
spin configuration should result in (+) spin in 1^24-
A discussion of the 2,6-PdAdH complex is deferred,
because the complex is of D0(., symmetry rather than D^ symmetry,
it will be discussed separately (vide infra).

55
Because it also places opposing spins at the same posi¬
tion in the ligand, spin polarization adds further ambiguity
to the interpretation of the contact shifts in these five-
member (or any odd member) chelate ring complexes. If, for
example, the spin is delocalized into a a orbital, this places
metal spin into s_ orbitals of all atoms. Also, the bulk of
the delocalized spin resides in the lone pair of the coor¬
dinated nitrogen atom causing a "slight unpairing of all
filled o bonds" by spin polarization effects (23). This is
readily seen for GaH by structures such as those in Figure 6.
If (i) spin is considered to be placed in a MO at the
coordinated nitrogen, the other electrons associated with
that nitrogen atom will also align with (t) spin preferentially
(Hund's rule of maximum multiplicity). The electron associated
with the adjacent atom must be paired, (i), with the nitrogen
atom electron. All the electrons associated with the adjacent
atom must also align (+) (Hund's rule), and so on through
the molecule. Thus spin polarization places (+) spin and
(;) spin at each position simultaneously. These considerations,
when applied to the other ligands, lead to the same con¬
clusions for each of the other ligands. Thus the ambiguity.
The foregoing discussion certainly indicates there are
many difficulties in reaching definitive conclusions about
the contact shifts of these compounds. However, considera¬
tion of how the lack of agreement between INDO calculated
and experimentally determined spin densities relates to low-
spin contributions may be more enlightening. In each instance

56
Hi
Hi
Ht
Ht
Figure 6.
Hi
tH
c;
fNt 1N1
T 4
tH
|N
I
I
tH
H t
4H
"Cl
tC
t
t
tNf
t
_tM
IH
|H
Representation of spin polarization effects in a
fragment of a [M(GdH)^+] complex

57
for which the signs and relative magnitudes of the calculated
densities agree with the experimentally determined densities
(Table V), the largest density is on the protons of the
amine or imine moiety; this density is always positive. The
smaller density is on the remaining protons of the ligand;
but the agreement in relative magnitude is only fair. This
may indicate contribution from both high-spin and low-spin
species. As can be deduced from the transfer mechanisms in
Figure 3 and the electronic configurations in Figure 4, con¬
tribution from both configurations would place (f) spin
and (4-) spin at the same proton. For example, if the elec-
tronic configuration is Is {(e)^(a^)^(e)1}, the unpaired
spin occupies the e° MO and will align parallel with the
external magnetic field. The transfer must be L + M. Con¬
sequently, (;) spin will remain on the ligand. By comparison,
if the electronic configuration is hs {(e)4(a-^) (e; }, the
unpaired spins occupy the a1^ and ea and will also align
parallel with the magnetic field. The transfer mechanism
may be either L + M (<¡>£ ea) or M ->- L (a^ -> tf>£) . If it is
L -+ M, ( + ) spin will remain on the ligand. Consequently, if
there is low-spin contribution to the contact shift, the
effect of the opposing spins on the ligand would be a smaller
shift and therefore a smaller relative density than that
calculated.
The preceding syllogism can be applied to the contact
shifts of these cobalt(II) compounds. Susceptibility studies
yielding intermediate, room temperature, magnetic moments of

58
4.65 BM and 4.23 BM for the BdH and GdH complexes, respec¬
tively, are taken as an indication of relative contributions
from both high-spin and low-spin configurations. The low-
spin contribution is about the same in the GdH complex as in
the BdH complex and is presumably reflected in the similar
disagreement between the INDO calculated and the experi¬
mentally determined relative magnitudes of spin density.
Consequently, it might be inferred that the effect of a con-
O
tact shift contribution attributable to the low-sp in ( E)
configuration is observed in these isotropic shifts. By
comparison, the relative magnitudes of the experimentally
determined densities for the BMI complex are much larger —
not smaller — than those calculated by the INDO method. The
intermediate, room temperature, magnetic moment of 3.30 BM
of the BMI complex also indicates contribution from both
high-spin and low-spin configurations, and indicates a greater
relative contribution from the low-spin configuration than
in the BdH and GdH complexes. It may be that in the BMI
complex the high-spin and low-spin contribuí ions"reenforce
rather than oppose one another. This mutual reenforcement
of high-spin and low-spin contribution may also account for
the larger shift of the BMI methyl protons compared to the
BdH methyl protons.
It should be reemphasized at this point that with the
present stage of development of theory it is not possible

59
to distinguish - with certainty - between contributions from
2
the E term and spin polarization involved in these iso¬
tropic shifts. As pointed out in the introduction contact
shifts of cobalt(II) complexes are previously known to arise
from the high-spin configuration (^T under 0^) only.
The 2,6-PdAdH complex was not considered in the pre¬
ceding discussion because its symmetry is D2d rather than
D^. Under D2d symmetry 2Eg(0h) -* A-^ 4- B^D^) and ^Tq(0h)
+ B2 + E(D2d) (6). As under all terms contribute zero
orbital angular momentum (A, B, and E). Consequently, the
isotropic shifts are due entirely to contact interaction.
Therefore pseudocontact contributions should be negligible.
The possible d orbital correlation diagrams and electronic
configurations for D2d are shown in Figure 7. Based upon
the room temperature magnetic moment of 2.28 BM, a low-spin
configuration should apply. This value is effectively
within the range generally accepted for six-coordinate low-
spin cobalt(II) complexes (1). Nonetheless, a high-spin
configuration is also feasible under D2d, a condition which
can give rise to a degree of anomalous magnetic behavior as
is possible under octahedral symmetry.
None of the INDO spin density calculations agrees in
either sign or relative magnitude with the experimental den¬
sities of the 2,6-PdAdH complex. Consequently no assertion
can be made as to whether the spin transfer mechanism is
M ->â–  L and/or L -+ M. The complete lack of agreement for the
2,6-PdAdH complex completes the trend discussed above for

60
'g
al + bl
Is { (e)4 (b2) 2 (a^ or b-^)1}
b,
â– 2g-.
hs {(e)4(b2)1(a1)1(b1)1}
0
h
°2d
al + bl
Is { (b9) 4 (e) 4 (a-^ or b-^)b}
hs {(b2)2(e)3(a1)1(b1)1}
0
h
Figure 7. Possible 0^--D9^ correlation diagrams and electron
configurations“tor a d7 ion under symmetry

61
the BdH, GdH and BMI complexes. That is, increasing dis¬
agreement between INDO calculated spin densities and experi¬
mentally determined spin densities may reflect a greater
low-spin contribution in each complex. The low-spin con¬
tribution is smaller in the BdH complex and the agreement
is better. By comparison, the low-spin contribution is
greater in the 2,6-PdAdH complex and there is no agreement
between the INDO calculated and experimentally determined
spin densities.
It appears that the low-spin configuration may indeed
contribute to the contact shifts of these cobalt(II) com¬
plexes. It is also of interest to note that the hydrazone
amine and imine proton resonances are shifted downfield,
and are the largest shifts observed in each complex. There¬
fore it is reasonable to conclude that they are dominated
by the same delocalization mechanism. Since delocalization
into a a MO leads irrevocably to a downfield shift (23), it
is tempting to propose that delocalization into a a MO
dominates at these positions with an opposing tt -MO contribu¬
tion which makes an upfield contribution to the contact shift.
This postulate is in keeping with that of Stoufer and co¬
workers (41) that tt back bonding plays an important role in
these anomalous cobalt(II) complexes.
Although attempts were made to prepare and investigate
the proton contact shifts of cobalt(II) complexes of all the
ligands listed in the introduction, these attempts did not
meet with success in some cases. Specifically, it was not
possible to prepare the complexes of the ligands in Figure 2.

62
The tris-glyoxalbismethylimine cobalt(II) complex decomposes
so rapidly it can not be isolated in pure form. The complexes
of the remaining ligands could not be prepared, probably be¬
cause of steric and/or electronic effects (20). That is,
either intramolecular steric strain or the electron with¬
drawing capabilities of the substituents was sufficiently
great to prevent coordination of the ligand to the metal ion.
In conclusion, although the modified INDO program used
was capable of handling the minimum basis set of the simple
tris-glyoxaldihydrazone cobalt(II) complex, the spin den¬
sities calculated were so small as to be of little value.
Consequently the commonly used approach of treating each
ligand as a radical anion and/or cation was used for this
work. However the various ambiguities which arise from this
treatment (spin polarization, delocalization into odd number
chelate ring complexes, uncertainty as to relative energies
of ligand orbitals upon interaction with the metal ion)
require caution in making very definitive conclusions. There
is evidence, however, that it back bonding plays an important
role, and that there is contribution to the contact shifts
2 2
attributable to the low-spin, E(D^) derived from Eg(O^),
configuration.
This investigation of contact shifts of magnetically
anomalous cobalt(II) complexes has lead to essentially the
same conclusion reached by Bertini and Getteschi (3). "The
enthusiasm of researchers in the proton contact field has
recently attenuated owing to the lack of a simple model
correlating the observed shifts to the nature of the

increasing sophistication have been proposed and laborious
calculations have been performed which gained in some case
partial success but never found general applicability. In
many instances however these partial models have been the
origin of lively controversies."

APPENDIX
Chemical shift assignments for the proton magnetic res¬
onance spectra of the 2-pyridinalphenylimine ligands are
listed in Table VII. They are based on the assignments
made by Matsubayashi, Okunaka, and Tanaka, which were made
for substituted 2-pyridinalphenylimines, in which the 4-
phenyl substituents are OCH^ , CH^, and Cl (31). The follow¬
ing numbering system applies.
4
a
Table VIII is a list of the chemical shifts of the tris-
(2-pyridinalphenylimine) complexes. The purpose in preparing
these complexes was to obtain and investigate contact shifts
of a homologous series of complexes 'which span both sides
of the cross-over point. Because this situation did not ob¬
tain, Table IX, the compounds are of little value for this
purpose. Furthermore the symmetry of these complexes is no
higher than (face isomer) and may be C-^ (edge isomer) .
64

65
Table VII
Proton Chemical Shifts of the Substituted
2-Pyridinalphenylimine Ligands
Position H
nh2
ch3
Cl
2-PAH
h3
8.15
8.15
8.17
8.16
7.96
«4
7.83
7.85
7.85
7.86
7.73
H5
7.45
7.39
7.41
7.46
7.23
H6
8.66
8.67
8.69
8.70
8.63
H
a
8.56
8.62
8.61
8.58
8.13
H
7.30
7.25
7.23
7.31
a
a
H6
7.30
6.71
7.23
7.38
-
H
X
7.30
4.22
2.34
-
-
Note :
Chemical shifts
at 30° C.
are in
ppm's relative to TMS
in CD3CN
£
There are no phenyl protons in the 2-PAH ligand. However,
the NHg chemical shift in dg-DMSO was found to be 4.13 ppm
from TMS.

66
Table VIII
Proton Chemical Shifts of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II) Complexes
[Co(2-PAH)^ ]
[Co(2-PAI)^ ]
7
. 1
[Co(2-PPI)^+]
- 51
. 5
7
. 6
- 40.
. 3
15.
. 7
- 25.
.5
51.
. 2
- 19.
. 6
61.
. 7
- 14.
. 8
64.
. 2
- 11.
, 0
81
. 7
- 8.
, 0
107.
. 8
- 6.
. 5
128.
. 8
4.
. 2
163.
, 7
6.
,8
174.
, 7
8.
9
182.
7
10.
4
11.
0
53.
4
16.
, 0
42.
0
16.
6
32.
3
17.
6
29.
8
18.
7
- 1.
9
51.
5
3.
9
58.
1
5.
4
67.
3
7.
3
70.
3
9.
5
72.
1
14.
1
76.
6
15.
1
87.
8
15.
4
Ill.
5
50.
7
122.
2
53.
3
58.
2
67.
4
68.
7
72.
2
92.
8
144.
7

67
[Co (2
Note :
Table VIII - continued
53
. 2
[Co (2-PMI)^"]
- 51
. 1
â–  40
.4
- 38
. 6
28
.8
- 28
. 1
21
. 1
- 20
.7
6
. 3
6
.2
8.
. 3
8
. 2
11.
. 1
10,
. 0
15.
. 1
10.
. 6
16.
, 1
11.
. 6
17.
9
16.
. 0
19.
4
16.
.4
51.
8
16.
. 9
59.
3
17.
7
69.
6
51.
5
72.
7
57 .
1
78.
1
58.
0
89.
5
64.
4
112.
0
70.
8
125.
0
74.
5
77 .
5
97.
5
115.
7
Chemical shifts are given in ppm from TMS in CD^CN
at 30° C. J

68
Table IX
Effective Magnetic Moments of the Substituted
Tris(2-pyridinalphenylimine) Cobalt(II) Compounds
Compounds
Meff (EM)
^eff(soln)
Co(2-PAH)3(BF^)£
b
4.61
Co(2-PPI)3(BF4)2
4.68
4.60
Co(2-PAI)3(BF4)2
4.64
4.09
Co(2-PMI)3(BF4)2
b
4.45
Co(2-PCI)3(BF4)2
4.42
c
Determined by the method of Loliger and Sheffold (27).
^Undetermined
Q
Co(2-PCI)^(BF^)9 was not sufficiently soluble to determine
Pe££ in solution.

69
In addition, a mixture of the two isomers is obtained; there
fore, two superimposed spectra resulted. Interpretation of
contact shifts of a mixture of these complexes of such low
symmetry would be even more difficult than for complexes.
Consequently, neither assignments of the contact shifts nor
a discussion of their significance was attempted.

BIBLIOGRAPHY
1. E. K. Barefield and D. K. Busch, Quart. Rev., 22,
457 (1968). ~
2. G. C. Bassler and Robert M. Silverstein, Spectrometric
Identification of Organic Compounds (John Wiley & Sons,
Inc., New York, 1968), pp. 64 - 78.
3. I. Bertini and D. Gatteschi, Inorg. Chem., 12, 2740
(1973).
4. N. Bloembereen and L. 0. Morgan, J. Chem. Phys., 34,
842 (1961)."
5. E. A. Clevenger, M. S. Thesis, The University of Florida,
Gainesville, 1961.
6. F. A. Cotton, Chemical Applications of Group Theory
(Interscience Publishers, New York, 1967), p. 285.
7. D. R. Eaton, NMR of Paramagnetic Molecules (G. N. La Mar,
W. DeW. Horrocks, Jr., and R. H. Holm, Eds., Academic
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BIOGRAPHICAL SKETCH
Joseph B. Vaughn, Jr, was born April 21, 1942, in Elfers,
Florida. He graduated from Paxon High School in Jacksonville,
Florida, in June, 1960. After a year of work and frugality,
he entered the University of Florida in 1961 and completed
his freshman year. After three more years of work, frugality,
and military service, he returned to the University of Florida
in September, 1964, and received the Bachelor of Science
degree with Honors in Chemistry in June, 1968.
Graduate study began in September, 1968, at the University
of Florida. From September, 1968, until June, 1973, he held
a teaching assistantship, and in September, 1973, no longer
eligible for an assistantship, be began work for the Pharmacy
Department at the University of Florida. From December, 1973,
until June, 1976, he held a medical technician position at
the V. A. Hospital in Gainesville, Florida. He also held an
Instructor's position, June, 1976, to September, 1976, with
Santa Fe Community College in Gainesville, Florida. Returning
to his graduate work at the University of Florida in September,
1976, he received the Ph. D. degree in August, 1977.
73

I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
R. Carl1' Stoufer, Chairman
Associate Professor' of Chárííistry
I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
Wallace S.
Professor
rr
Brey, Jr.
of Chemistry .
I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
-v^wvl I V i t1
Gerhard M. Schmid
Associate Professor of Chemistry

I certify that I have read this study and that in my
opinion it conforms to acceptable standards of scholarly
presentation and is fully adequate, in scope and quality,
as a dissertation for the degree of Doctor of Philosophy.
Richard D. Dresdner
Professor of Chemis-try
I certify that
opinion it conforms
presentation and is
as a dissertation fo
I have read this study and that in my
to acceptable standards of scholarly
fully adequate, in scope and quality,
r the degree of Doctor of Philosophy.
fdocJ Um'
Paul Urone
Professor of Environmental
Engineering Sciences
This dissertation was submitted to the Graduate Faculty
of the Department of Chemistry in the College of Arts and
Sciences and to the Graduate Council, and was accepted as
partial fulfillment of the requirements for the degree of
Doctor of Philosophy.
August 1977
Dean, Graduate School

UNIVERSITY OF FLORIDA
3 1262 08553 5226

UNIVERSITY OF FLORIDA
3 1262 08553 5226



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