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The preparation and properties of pyrolyzed polyacrylonitrile catalyst materials

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Title:
The preparation and properties of pyrolyzed polyacrylonitrile catalyst materials
Creator:
Clark, Jeffrey Lee, 1954-
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[s.n.]
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Language:
English
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ix, 161 leaves : ill. ; 28 cm.

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Subjects / Keywords:
Carbon ( jstor )
Catalysts ( jstor )
Catalytic activity ( jstor )
Chlorides ( jstor )
Ethanol ( jstor )
Hydrogen ( jstor )
Nitrogen ( jstor )
Pyrolysis ( jstor )
Surface areas ( jstor )
Temperature control ( jstor )
Acrylonitrile ( lcsh )
Catalysts ( lcsh )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Pyrolysis ( lcsh )
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bibliography ( marcgt )
non-fiction ( marcgt )

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Thesis:
Thesis (Ph. D.)--University of Florida, 1987.
Bibliography:
Bibliography: leaves 156-160.
General Note:
Typescript.
General Note:
Vita.
Statement of Responsibility:
by Jeffrey Lee Clark.

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Full Text


THE PREPARATION AND PROPERTIES OF PYROLYZED
POLYACRYLONITRILE CATALYST MATERIALS
By
JEFFREY LEE CLARK
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1987


DEDICATION
To IC, who almost never lost the faith.


ACKNOWLEDGEMENTS
There are many people who facilitated this work through
actual assistance as well as thoughtful discussion and
encouragement, and although there are too many to mention,
their contribution is not overlooked. I am especially
appreciative to the glass, machine and electronics shops for
their invaluable assistance in constructing and repairing
the equipment used in this study. I would also like to
thank Dr. Willie Hendrickson of the 3M corporation, Ann
Livesey of the U.S. Army, and Ngai Wong of the University of
Florida for their generous assistance in this work. In
addition, I am grateful to the University of Florida, the
U.S. Army, and Geo-Centers Inc. for financial support during
this endeavor. Finally, I am deeply indebted to Dr. Russell
S. Drago for the freedom and support he provided me during
this study.
in


TABLE OF CONTENTS
Page
ACKNOWLEDGEMENTS iii
LIST OF TABLES v
LIST OF FIGURES vi
ABSTRACT viii
INTRODUCTION 1
BACKGROUND 4
EXPERIMENTAL 2 8
Materials 28
Preparations 29
Methods 3 4
RESULTS AND DISCUSSION 50
Pyrolysis Studies 50
Basicity Studies 76
Catalytic Results 86
SUMMARY AND FUTURE RECOMMENDATIONS 132
APPENDIX: XPS SPECTRA 138
REFERENCES 156
BIOGRAPHICAL SKETCH 161
IV


LIST OF TABLES
Page
1. Elemental Analysis Results 52
2. Surface Area Results 74
3. CEES Products 98
4. XPS Results; Elemental Ratios 108
5. XPS Results; Ionization Energies 109
6. Effects of Adsorbed Metal Salts; Series 1 114
7. Effects of Adsorbed Metal Salts; Series 2 117
8. Miscellaneous Catalyst Results 119
9. Glass Spheres Results 122
10. Elemental Analysis Results For Suported
Catalysts 125
11. Catalytic Results For Supported Catalysts 127
12. GC-MS Results for Discolored Products 130
v


LIST OF FIGURES
Page
1. Proposed Pyrolysis Reaction 6
2. Imine-Nitrone Structure 7
3. Proposed Catalytic Mechanism 11
4. Dehydrogenation of Acridine 13
5. Possible Dehydrogenation Mechanisms 14
6. Proposed Structure and Pyrolysis Product for
Polycyanoacetylene 16
7. Initial Pyrolysis Apparatus 3 6
8. Modified Pyrolysis Apparatus 38
9. Initial Catalytic Evaluation System 41
10. Oven Design 43
11. Improved Catalytic Evaluation System 45
12. Syringe Design 46
13. Representitive Temperature-Time Profiles 51
14. Isothermal Thermogravimetric Analysis 54
15. Thermogravimetric Analysis, 2.5C/minute 56
16. Thermogravimetric Analysis, 5.0C/minute 57
17. Thermogravimetric Analysis, 10.0C/minute 58
18. Composite of Temperature Programmed
Thermogravimetric Analysis 59
19. Thermogravimetric Analysis, 0.5C/minute 61
20. Thermogravimetric Analysis and Differential
Thermal Analysis 62
vi


21. Temperature Programmed Differential Scanning
Calorimetry 65
22. Differential Scanning Calorimetry, 0.5C/minute 66
23. Isothermal Differential Scanning Calorimetry 68
24. Plot of Residual Nitrile Content Versus Area
of Exotherm 7 0
25. Naptheridine Binding Modes 79
26. Electron Paramagnetic Resonance Spectra of PPAN
and Co(DMGH)2 81
27. Electron Paramagnetic Resonance Spectra of PPAN
and Copper (II) Chloride 83
28. Polymerization Mechanisms 94
29. Structures of Nerve Gases and
Dimethyl-methylphosphonate 103
30. Selectivities of Metal Doped Catalysts 133
Vll


Abstract of Dissertation Presented to
the Graduate School of the University of Florida
in Partial Fulfillment of the Requirements for
the Degree of Doctor of Philosophy
PREPARATION AND PROPERTIES OF PYROLYZED
POLYACRYLONITRILE CATALYST
MATERIALS
By
Jeffrey Lee Clark
December, 1987
Chairman: Russell S. Drago
Major Department: Chemistry
The preparation and properties of pyrolyzed
polyacrylonitrile (PPAN) catalyst materials was
investigated. Specifically, the effects of variations in
the pyrolysis conditions (heating rate, maximum temperature,
atmosphere, etc.) were examined using elemental analysis,
Differential Scanning Calorimetry and Thermogravimetric
Analysis. In addition, the effects of metal additives and
oxide supports on the catalytic activity of PPAN catalysts
were studied using the dehydrogenation of ethanol as a test
reaction. It was found that small amounts of metal
additives were capable of profoundly affecting both the
activity and selectivity of PPAN catalysts. Supporting PPAN
on oxide supports served to increase the surface area by
Vlll


greater than an order of magnitude. However, the catalytic
activity seemed to be more a reflection of the catalytic
activity of the support used because blank runs using only
the support resulted in very similar selectivity and
activity. In addition, PPAN catalysts were found to
decompose chloroethyl ethylsulfide at temperatures as low as
200C under both aerobic and anaerobic conditions.
IX


INTRODUCTION
In the last twenty years, due to the increased
consumption of manmade materials and energy, the area of
catalysis has blossomed into a science in its own right.
Due to its more general application to spectroscopic
techniques, homogeneous catalysis has traditionally been
better understood than heterogeneous catalysis, which is
less amenable to spectroscopic techniques, making study
more difficult. In fact, it has been frequently stated by
industrial catalyst chemists that heterogeneous catalyst
preparation is more an art than a science. However, recent
advances in surface analytical techniques have enabled
chemists to more fully understand the reactions occurring on
the surfaces of heterogeneous catalysts. At present,
though, the high relative cost of these techniques is
frequently prohibitive in many academic environments.
Metals and their complexes and salts have traditionally
been the catalysts of choice industrially. Indeed, one of
the first catalytic chemical plants (excluding biological
"plants" of course) was built by Germany during World War II
1


2
and utilized a catalyst composed of mixed metal oxides to
convert CO and H2 (syngas) into liquid fuel using the
Fischer Tropsch Process.1 Both the petroleum and polymer
industries initially led research efforts in catalysis. With
the advent of automobile exhaust emission controls, however,
extensive research was dedicated to developing effective
emission control catalysts.2 In many cases, the catalysts
employed in industrial operations consist of small metal
crystallites or complexes supported on the surface of an
inorganic oxide (such as silica gel (Si02), titanium dioxide
(Ti02) and alumina (A1203)), whose primary function is to
uniformly disperse the active catalyst species. This
provides greater efficiency per unit of metal species by
increasing the catalytic surface area available for chemical
reaction. Due to the high cost of precious and strategic
metals, interest has been directed towards less expensive
alternatives to precious metal catalysts. One area of
investigation has involved the use of heat stable organic
materials. In particular, organic pyropolymers have been
examined as possible catalyst candidates due to their
paramagnetism, semiconducting tendencies and heat stability.
One such polymer, pyrolyzed polyacrylonitrile (PPAN), has
been studied extensively as a catalyst for several
reactions. At present there are very few published reports
on the effects of metal additives or of supporting PPAN
catalysts on inorganic oxide supports. This study involves


3
a systematic investigation into the effects of variations
the preparation of PPAN catalyst materials.
in


BACKGROUND
It was reported in 1958 by Burlant and Parsons3 that
upon treatment with thermal or ultraviolet radiation,
polyacrylonitrile (PAN) undergoes a chemical reaction
leading to discoloration and the appearance of an Electron
Paramagnetic Resonance (EPR) signal. Since that time,
research has been carried out on this process for three
general reasons: (1) textile chemists are concerned with
eliminating this phenomenon, which causes premature decay of
PAN containing fabrics; (2) catalyst chemists have been
interested in exploiting the catalytic properties inherent
in pyrolyzed polyacrylonitrile; and (3) materials scientists
have been interested in this process since it is the first
step in the conversion of acrylonitrile to carbon fibers,
which are an increasingly used component of composite
structural materials. Pyrolyzed polyacrylonitrile films
have even been proposed as a cost-effective amorphous
semiconducting material to replace silicon crystals in solar
4
collectors.
The earliest studies on the pyrolysis of
polacrylonitrile were infrared studies in which it was
observed that the carbon-nitrogen triple bond stretch
disappeared, the bands in the spectrum broadened
4


5
considerably, and a large broad absorption grew in the
region where C=N, C=0 and C=C would appear. However, the
spectrum was too featureless to permit specific assignments.
On the basis of this evidence, it was proposed that the
reactions occurring during pyrolysis and polymerization were
as illustrated in Figure l.5'6'7 a recent spectroscopic
study on the vacuum pyrolysis of PAN8 concluded that an
intermediate phase existed at pyrolysis temperatures between
200 and 260C. This material was said to be an intrinsic
semiconductor with an extensively delocalized n electron
system as depicted in Figure lb. They also suggested that
the reaction which occurs in this temperature range could
involve bonding and conjugation between adjacent chains
without interruption of the carbon backbone, due to the
atactic nature of the PAN starting material. This reaction
has also been proposed to occur during the alkaline
degradation of PAN.9 It has also been proposed that
partially oxidized species such as the imine-nitrone
copolymer illustrated in Figure 2 could be present and
contribute to the optical properties of PPAN (not pyrolyzed
under vacuum), based on studies done on synthetic model
compounds which have similar absorptions in the UV-visible
region.
PPAN exhibits a strong, fairly narrow EPR signal at
about g=2, and this has been used as a probe of the
pyrolysis reaction. The g value increases with increasing


6
nC H2=C H
i
C=N
POLYMERIZATION
a.
CH CH CH
\\ \\ %
N N N N
1
PYROLYSIS
CHo CHp CH2 c H2
CH CH CH
C c c c
7 \ / \ / \ / \
N N N N
I
PYROLYSIS
CH CH CH C^H
c.
\
/Cx\ /C. C.\
/ % / \ / % / \
N N N N
Figure 1: Proposed Pyrolysis Reaction


7
O
t
Figure 2: Imine-Nitrone Structure


8
thermal treatment time, becoming essentially constant after
about five hours.12 It was learned that the concentration
of unpaired spins per gram (as determined by an EPR
technique) was strongly dependent upon the pyrolysis
conditions as well as the temperature of measurements. In
addition, the number of unpaired spins per gram was
decreased by the presence of air during the measurement.
The important parameters of the pyrolysis seemed to be the
rate and duration of heating in addition to the atmosphere
in which the pyrolysis was being carried out. For example,
a sample pyrolyzed in an ammonia atmosphere contained more
unpaired spins per gram than an identical sample pyrolyzed
in air or nitrogen.13
There have also been studies to determine what volatile
products are formed during the pyrolysis reaction and it was
found that ammonia, hydrogen cyanide, acetonitrile,
acrylonitrile monomer, propionitrile, methacrylonitrile,
isobutylacrylonitrile and vinyl acrylonitrile were formed
during the reactions.14 It was also found that the relative
amount of volatiles produced generally increased with
temperature in the range studied (300-800C).
Unfortunately, the temperature-time profile of these
experiments consisted of very fast rise times (<1 sec) to
the desired temperature followed by maintenance at that
temperature for 10-20 seconds after which time the products
were analyzed by Gas Chromatography (GC). The temperature-


9
time profile refers to a plot of the temperature versus
time, which enables one to make better comparisons of the
thermal history of each sample. It has been known for some
time that there is a strong exothermic reaction which occurs
between 200 and 300C which is capable of generating enough
heat in a bulk sample to cause ignition of the sample15
concomitant with drastic weight losses resulting in a
material which has little catalytic activity. (Several
times during our investigations at the University of Florida
these runaway reactions were observed due to the
catastrophic failure of a temperature controller.) Since
the active catalytic species is presumed to be the product
of the reaction in Figure 1, it is logical to choose the
pyrolysis conditions designed to minimize weight losses
since the cyclization-dehydrogenation reactions depicted in
Figure 1 would result in a theoretical weight loss of less
than 5%.
There have been many reports describing the catalytic
activity of pyrolyzed polyacrylonitrile. PPAN has been
shown to be capable of isomerizing alkenes and
dehydrogenating alcohols,16 decomposing formic acid and
nitric oxide,17'18 dehydrogenating ethylbenzene19 and
cumene,20 and oxidizing ethylene to ethylene oxide.21
Manassen and coworkers, in a series of publications,
described the dehydrogenation of alkenes and of alcohols and
the isomerization of alkenes using PPAN catalysts. It was


10
found that unlike conventional dehydrogenation catalysts, no
gaseous hydrogen was evolved. Instead, the hydrogen was
thought to be either physically adsorbed or chemically bound
to the surface of the catalyst, and while prolonged heating
at elevated temperatures in a nitrogen atmosphere did not
restore the catalytic activity, a short treatment in air at
temperatures as low as 140C completely restored the
activity with the concomitant production of water. The
dehydrogenation of cyclohexene to benzene was also observed
to follow a different pathway using PPAN catalysts in that
no disproportionation occurred as with commercial
dehydrogenation catalysts. When cyclohexene was passed over
various types of charcoals or graphite at elevated
temperatures, significant amounts of cyclohexane and other
products were formed while PPAN catalysts produced benzene
exclusively. In addition, as with the alcohols tested, no
gaseous hydrogen was produced and prolonged heat treatment
under a nitrogen atmosphere failed to restore the catalytic
activity while a short treatment in air was found to
completely restore the catalytic activity. It was also
found that dehydrogenation reactions tended to deactivate
these catalysts quicker than isomerization reactions, again
suggesting that the hydrogenation of the catalyst surface is
responsible for the deactivation of the catalyst. These
authors proposed the reaction illustrated in Figure 3 to be
the mechanism for the catalytic activity of PPAN catalysts.


11
H
Figure 3: Proposed Catalytic Mechanism


12
It has been demonstrated by Braude and coworkers22 that
dihydroquinoline can by hydrogenated by acridine while the
reverse reaction does not occur, as shown in Figure 4.
On the basis of these observations it was proposed that
higher annelation of a system of condensed heterocyclic
aromatic rings will result in poor H-donating properties.
It is also known that annularly condensed aromatic molecules
become less and less stable with an increasing number of
rings while the hydrogenated compounds gain in stability,23
which may help to explain why a compound containing the
proposed structure (Figure 1) of PPAN catalysts could be
both a good hydrogen acceptor and a poor donor.
Furthermore, model compounds were synthesized which were
incapable of forming hydroaromatic structures while
retaining aromaticity and were found to be catalytically
inactive.24 In addition, the transformation of a
hydroaromatic structure to an aromatic structure by air
oxidation has been shown to occur with acridine25 and model
compounds of condensed pyridine rings.26
Manassen and coworkers used the reaction of
5-ethyl-5-methyl-l,3-cyclo-hexadiene over PPAN catalysts in
an attempt to elucidate the form of the hydrogen transfer.
The scheme in Figure 5 illustrates the possible products
based on hydride ion abstraction, hydrogen atom abstraction
and proton abstraction. A comparison of the product
distributions obtained by passing


13
Figure 4: Dehydrogenation of Acridine


14
(a)
Hydri de
ion abstraction
5-ethyl-5-methyl -
1.3-cyc1onexadiene
abs tract i o*
(b) Hydrogen atom
5-ethyl-5-methv1 -
1,3-cyc1ohexadi ene
+ J
ethyl radical
(c) Proton abstraction
5-ethy1-5-methy 1 -
1,3-cyclohexadiene
methyl
carbanion
Figure 5: Possible Dehydrogenation Mechanisms


15
5-ethyl-5-methyl-l,3-cyclohexadiene over various catalysts
led the authors to conclude that the behavior of PPAN
catalysts is somewhat intermediate between that of
commercial dehydrogenation catalysts, which abstract H-atoms
and acidic alumina, which is a proton abstractor. Pyrolyzed
polycyanoacetylene, the structure of which is illustrated in
Figure 6, was found to exhibit only H-atom abstraction,
leading to the proposal that the hydride transfer aspect of
PPAN catalysis originates from another structural element
than the proposed structure of the active catalyst, since
both PPAN and pyrolyzed polycyanoacetylene should
theoretically have the same structure. They found support
for this hypothesis by the observation that treating the
PPAN with dimethyl sulfoxide can greatly enhance the
percentage of ortho-ethyl-toluene formed.
It is of interest to note that different preparations of
PPAN catalysts can have very different catalytic activities.
For example, Manassen and coworkers used the following
general preparative method for synthesizing their catalysts.
Polyacrylonitrile was spread out in a thin layer in a draft
oven, slowly heated to 230C and kept at this temperature
for 12 hours, during which time the color went from white to
brown via yellow. The brown powder was pelleted at 8000
psi, crushed, sieved and then calcined at either 350 or
450C for 30 minutes, during which time the color went from
brown to black. It was found that catalysts calcined at


16
Figure 6: Proposed Structure and Pyrolysis Product for
Polycyanoacetylene


17
450C underwent decomposition, presumably yielding large
crosslinked structures, which were apparently somewhat
acidic resulting in a high activity for double bond shifts
as well as ethyl shifts. (It was the material calcined at
450C which was used for the 5-ethyl,5-methyl-l,3-
cyclohexadiene experiment.) PPAN catalysts which were
calcined at 350C showed much less activity for double bond
or ethyl shifts, which illustrates the importance of the
pyrolysis conditions on the activity of these catalysts.
These results are in agreement with electrical measurements
which have shown that drastic changes occur above 350C.26
Another theory to account for the catalytic properties
of organic pyropolymers correlated the number of unpaired
spins per gram (as as measured by an EPR technique) to the
catalytic activity for the decomposition of nitrous
oxides.28'29 This theory includes a scaling factor called
the intrinsic activity of the free spins, which presupposes
that all of the free spins are not active in catalysis. The
assumption is that structural rearrangements preceding
graphitization bring about exchange interactions between the
free spins rendering some of them catalytically inactive. A
correlation was developed between the width of the EPR line
and the relaxation time, T1 (as determined by the saturation
technique) and the intrinsic catalytic activity of the free
spins.30 This concept seems somewhat related to the
electronic theory of catalysis on semiconductors, described


18
by Vol1kenstein31 and Hauffe.32 In the electronic
mechanism, the semiconductor catalyst acts as an electron
reservoir by either donating electrons to or accepting
electrons from the substrate in guestion. This mechanism is
in contrast to that proposed for alumina, in which the
reaction is suggested to occur on Bronsted acid or proton
donating sites through the formation of an adsorbed
carbonium ion. Cutlip and Peters15 examined the kinetics of
dehydration of t-butyl alcohol over PPAN catalysts at 240 to
280C and applied both single site and dual site models to
describe the kinetics. After a considerable amount of
mathematical manipulation, these authors concluded that it
was impossible, based on their data, to uneguivocally
determine the reaction mechanism since a statistical best
fit could not be obtained for any of the models chosen.15
Since the electronic theory of catalysis greatly increases
the possible number of rate limiting steps in the mechanism
(dissociation of adsorbed molecules, transfer of electrons
between adsorbed species and the catalyst), it was only
possible to gualitatively evaluate this mechanism, and it
was proposed that the rate-limiting step would probably be
related to the concentration of free electrons in the
conduction band of the catalyst. This could possibly be
related to what Gallard-Nechtschein and coworkers described
as the intrinsic catalytic activity of the free spins and
its relationship to the catalytic activity of these


19
catalysts,30 thus raising the possibility that there are
unpaired electrons in addition to conduction band electrons.
One problem with reviewing the literature pertaining to
PPAN catalysis is that one is consistently faced with the
prospect of comparing apples to oranges in the sense that
many of the studies reported preparing their catalysts by
different methods, making it impossible to be certain
whether or not the preparations consisted of the same
chemical (or semiconducting, for that matter) species. In
many cases, the catalysts were only characterized by their
method of preparation and their catalytic activity.
Manassen and coworkers24 (as discussed earlier) noted a
difference in catalytic activity, as well as in elemental
analyses (C, H, N), by calcining in nitrogen for 30 minutes
at 350C instead of 450C. Another pertinent observation is
the fact that the sum total of the carbon, hydrogen and
nitrogen analyses is usually between 75 and 90%, indicating
that from 10 to 25% of the final catalyst is composed of an
element other than carbon, hydrogen or nitrogen. The most
likely candidate is oxygen since many of the catalyst
preparation methods involve pyrolysis in air for some
period, and none have sought to rigorously exclude oxygen
from the preparations. This gives further evidence that
structures such as the imine-nitrone depicted in Figure 2
may constitute a considerable proportion of PPAN catalysts,


20
as well as possibly being responsible for the Bronsted type
acidity.
A recent spectroscopic study done on silver backed
polyacrylonitrile films also implicated structures such as
the conjugated imine in Figure 1 as products in the UV
degradation of PAN films using light in the 250-400 nm
region.33 (It has been known for some time that UV light
and strong bases bring about similar structural changes as
heat treatment, producing dark colored, intrinsically
paramagnetic solids.) In the UV degradation study, changes
in the silver backed PAN films were monitored using Fourier
transform infrared reflection absorbance (FTIR-RA) under
both oxidative and non-oxidative conditions. Their results
indicated that under oxidative conditions, oxygen-containing
species such as alcohols, carboxylic acids, hydroperoxides
or ethers could be present. Since UV light is capable of
generating ozone, ozonolysis products are also theoretically
possible, although the presence of ozone was not detected.
In addition, the appearance of an N-H stretch and a C=N
stretch, as well as the loss of the carbon nitrogen triple
bond stretch, suggested that cyclization had occurred. The
film became discolored and developed an EPR signal; however,
when the film was redissolved in dimethyl sulfoxide, the EPR
signal disappeared while the color remained. This indicated
that discoloration is a necessary but not sufficient
condition for the presence of paramagnetism. Although


21
analogies can be drawn between the UV and thermal
degradation of PAN, it must be emphasized that there are
distinct differences since thermal degradation generally
results in an insoluble material. For example, only one
hour of heating at 160C results in a solid of which only
20% can be extracted with dimethylformamide, the solvent of
choice for PAN. This may be attributed to more extensive
crosslinking occurring in the thermal process than in the
photochemical reaction, suggesting that the reactions in the
thermal process are considerably more complex. An
additional complication is the fact that all
polyacrylonitrile is not identical since the polymerization
conditions and polymerization catalysts used will have a
strong effect on the composition of the final product due to
defects and inevitable chain ends which may or may not
consist of polymerization catalyst residues. In studies of
the production of carbon fibers it has been learned that
pre-oxidation (prolonged isothermal heating at low to
moderate temperatures (150-250C)) minimizes the exothermic
reaction and yields a flameproof material which has a
concentration of approximately 10% oxygen by weight.34
Although most of the catalytic studies of PPAN-based
catalysts had chosen simple, industrially unimportant
reactions in an attempt to establish a relationship between
the catalytic activity and the chemical or electronic
structure of the material, there have been a few studies of


22
industrially important reactions. Degannes and Ruthven18
investigated the oxidative dehydrogenation of ethylbenzene
to styrene (at atmospheric pressure and from 180 to 280C)
and found that the reaction was zero order in oxygen and
approximately second order in ethylbenzene. The zero order
dependence on oxygen suggests that the reaction rate is
controlled by the rate at which ethylbenzene can be adsorbed
and dehydrogenated. Although most of these experiments were
for kinetic purposes carried out at low conversions under
differential conditions, a limited series of integral
experiments showed that conversions greater than 80% at
325C could be achieved with no significant by-products.
This is in contrast to commercially available
dehydrogenation catalysts which do not operate oxidatively,
resulting in an endothermic process requiring temperatures
in excess of 500C to achieve conversions of 50% to styrene,
with benzene and toluene being significant side
products.35,36,37
There have been very few studies in the literature
dealing with two potentially important areas: doping PPAN
catalysts with metals and supporting PPAN catalysts on
inorganic oxide supports. In one study, acrylonitrile was
polymerized using 2,2'-azobis[2-methyl propionitrile (AIBN)
in the presence of silica gel and the resulting material was
pyrolyzed at several different temperatures.20 Metals were
then added to these materials by slurrying with ethanol


23
solutions of the corresponding metal chlorides followed by
filtration and washing with ethanol. X-ray photoemission
spectroscopy (XPS) data were cited to propose copper binding
to nitrogen atoms in the PPAN since it was noted that only
the Cu 2p 3/2 and the N Is binding energies shifted upon
addition of CuCl to the silica supported PPAN samples. The
reactions studied were the oxidations of cumene and
ethylbenzene at 100C and under 1 atm. of oxygen. For
cumene, the total conversion was 63% with the selectivity
being 63% cumyl alcohol and 28% acetophenone; for
ethylbenzene, the total conversion was 21% with a
selectivity of 87% acetophenone and 13% 1-phenylethanol.20
These results for ethylbenzene are quite different from the
vapor phase results obtained by Degannes and Ruthven19 who
observed styrene as the only product. This apparent
contradiction is not too surprising when one considers the
differences in the catalyst preparation and catalytic
reaction conditions. The silica-supported PPAN catalyst
contained of course Si02 and was also different in that it
was pyrolyzed for 12 hrs. at 190C. The previous work by
Degannes and Ruthven had employed a catalyst which had
slowly been heated to 230C in air, was maintained at that
temperature for 16 hrs., and then was calcined at 400C in
an atmosphere of nitrogen for 4 hrs. In addition, one study
passed gaseous ethyl- benzene over the catalyst from 180 to


24
325C, while the other reaction occurred in the liquid phase
(or at the liquid-solid interface) at 100C.
The data available in the open literature concerning the
catalytic activity of PPAN are sketchy at best, which is
rather surprising since these few studies suggest that PPAN-
based materials could be promising catalysts for
dehydrogenation as well as oxidation and dehydration.
Perhaps one reason why there are so few reports of PPAN
catalysts is because of the formidable problems associated
with studying a material without being able to employ
common solution techniques such as nuclear magnetic
resonance spectroscopy, UV-vis spectroscopy and infrared
spectroscopy.
It was the goal of this investigation to increase the
growing body of knowledge concerning catalysis using PPAN-
based materials in three general areas: the effects of the
pyrolysis conditions on the catalytic activity of PPAN
catalysts; the effects of metal additives on the catalytic
activity of PPAN catalysts; and the effects of supporting
PPAN catalysts on oxide supports such as silica (Si02),
titania (Ti02) and alumina (A1203). As in many of the
earlier studies, a simple model reaction was chosen for ease
in handling the analysis. In this case, ethanol was chosen
as the substrate and was shown to be capable of undergoing
dehydrogenations as well as dehydrations.


25
The effects of variations in the pyrolysis conditions
were investigated by several methods. Differential scanning
calorimetry (DSC) and thermogravimetric analysis (TGA) were
carried out on PAN samples both isothermally and in the
temperature programmed mode in an attempt to learn more
about the destructive exotherm and concomitant weight loss
characteristic of the pyrolysis reaction. In addition,
several different atmospheres were used in the pyrolysis
reaction in an attempt to learn whether the pyrolysis
atmosphere affects either the composition or catalytic
activity of these materials. Finally, the effect of
variations in the pyrolysis reaction heating rate on the
composition and catalytic activity of the resulting
materials was investigated.
In the hopes of obtaining useful information about the
effects of metal additives on the catalytic activity of PPAN
catalysts, three general approaches were taken: the
intrinsic basicity of PPAN was studied by titrating with
dilute acid, paramagnetic metal species were deposited on
the surface of PPAN in an attempt to learn more about the
metal environment using electron paramagnetic resonance
spectroscopy (EPR), and the effects of various metal
additives on the catalytic activity of PPAN were examined in
the dehydrogenation of ethanol.
The investigation of the effects of supporting PPAN on
oxide supports was carried out by preparing samples of PPAN


26
supported on silica gel and alumina and then comparing the
catalytic activities of these materials in the
fore-mentioned ethanol reaction with unsupported PPAN
catalysts as well as pure silica gel and alumina. The
surface areas of these materials were also measured in order
to make activity comparisons on a surface area basis. As
stated previously, the goal of this research was to expand
the general knowledge about PPAN based catalytic materials
in the hope that a greater understanding of the
structure-reactivity relationships could be obtained so that
in the future it may be possible for chemists to engineer
low-cost organic catalysts which are tailored to optimize a
specific reaction. Although the work of one graduate
student is not sufficient to achieve these lofty goals, it
is possible to address some apparent inconsistencies in the
literature regarding the selectivity of PPAN catalysts in
the reaction of ethylbenzene and the significance of the
support interactions, if any. The report of silica
supported PPAN complexes by Bai and co-workers20 was very
brief and many important experiments were either not carried
out or not reported. For example, the activities of the
metal-silica-PPAN catalysts were ranked as follows: Cu I >
Cu II > Co II > Mn II. Unfortunately, no comparison was
made between doped and undoped catalyst preparations. In
addition, no blank runs using only silica were attempted,
making it impossible to determine what role (if any) the


27
silica support plays in the catalytic reaction. Although
ethanol was used as the substrate in the present study for
ease in handling, it was felt that the results obtained
would be applicable to the ethylbenzene reaction since both
involved primarily dehydrogenation reactions.
In addition, a limited number of experiments were
carried out to determine the feasibility of using PPAN
catalysts to decompose dimethyl methylphosphonate (DMMP) and
chloroethyl-ethylsulfide (CEES), two compounds of military
interest as simulants for chemical warfare agents.
Preliminary experiments were also carried out to determine
whether PPAN catalysts are photochemically active or active
towards syngas conversion.


EXPERIMENTAL
Materials
N,N-Dimethylformamide was reagent grade, purchased from
Aldrich and used without further purification.
Acrylonitrile was reagent grade, purchased from Aldrich and
used without further purification. AIBN (2,2'-azobis[2-
methyl propionitrile]) was purchased from Eastman Chemicals
and used without further purification. Polyacrylonitrile
was reagent grade, purchased from Aldrich and used without
further purification. Alumina (neutral) was purchased from
Fischer and used as supplied. Titanium dioxide (anatase,
Ti02) was supplied by Baeyer and used without further
purification. Metal complexes were all reagent grade and
used as supplied. Zirconium basic carbonate and zirconium
basic acetate were generously donated by Mr. Brady Crom and
Dr. Tom Wilson, of Zirtech Inc., Gainesville, Florida, and
were used without further purification. Vacuum distilled
chloroethyl ethylsulfide (CEES) was generously provided by
Dr. Yu-Chu Yang of the U. S. Army. Silica gel (Si02), Grade
62 with a mesh size of 60-200, was provided by Davison and
used without further purification. Ethylbenzene was reagent
28


29
grade, purchased from Aldrich and distilled before use to
remove traces of toluene and benzene. Dimethyl
methylphosphonate (DMMP) was reagent grade, supplied by the
U.S. Army and used without further purification. Silver (I)
trifluoromethanesulfonate (AgCF3S3) was purchased from
Aldrich and used without further purification. Bis(2,2'-
bipyridine) ruthenium (II) chloride (Ru(bipy)2cl2) was
generously provided by Dr. E. Stine and used without further
purification. Glass spheres (8-58 /m, Standard Reference
Material 1003a) were obtained from the National Bureau of
Standards and used without further purification.
Preparations
Polymerization of Acrylonitrile
Acrylonitrile (100 g were added to 1500 ml of distilled
water in a flask fitted with a reflux condenser and
maintained at 60C in a silicone oil bath under an inert
atmosphere. AIBN (1.00 g) was added to this mixture with
vigorous stirring. The mixture was allowed to stir for 1.0
hr after which another gram of AIBN was added. The
suspension was allowed to stir overnight and then the white
polyacrylonitrile powder was recovered by filtration and
washed with copious amounts of acetone. The resulting
polymer was dried at 50C in a vacuum oven. It was found


30
that vigorous stirring using an overhead stirrer was
required during the reaction to prevent agglomeration of the
polymer into large lumps. In addition, to prevent loss of
the monomer the cooling water was passed through a saltwater
ice bath, using a copper heat exchange coil, before entering
the condenser. Due to its greater density, argon was found
to be superior to nitrogen in preventing the loss of the
highly volatile monomer.
Diphenvlglvoximato Cobalt (II) (Co(DPGH)2)
This synthesis was carried out under argon in
schlenckware apparatus using the method of Tovrog.38
Methanol was dried by distillation from calcium hydride and
stored over 4 A molecular sieves. Dry methanol (200 ml) dry
methanol was added to a 1 L three-neck round bottom flask
and the oxygen was removed by bubbling argon through the
system. Cobalt (II) acetate (12.45 grams) and
diphenylglyoxime (24.00 grams) were added to the reaction
flask and the mixture was allowed to stir for several hours
until the cobalt (II) acetate dissolved. After several more
hours of stirring, the brown suspension was filtered through
a schlenckware frit yielding a brown solid.
Analyzed: Carbon 59.23%, hydrogen 4.11%, nitrogen
9.85%
Calculated for Co(DPGH)2'2H20: Carbon 58.99%, hydrogen
4.57%, nitrogen 9.77%


31
PAN on Silica
Two preparations of silica gel supported PAN were
formulated with loadings of approximately 7 and 17% PAN by
weight. The appropriate amount (2.00 grams or 5.00 grams)
of PAN was dissolved in 250 ml of N,N-dimethylformamide.
Silica gel (30 g) was added to the solution and the
suspension was allowed to stir on low heat (90C) for 18
hrs. The suspension was then rotary evaporated to dryness
and subsequently pyrolyzed. While the 7% PAN on silica gel
became an off-white free flowing solid after rotary
evaporation, the 17% PAN on silica gel became caked up and
required grinding in a mortar and pestle before pyrolysis.
PAN on alumina
This sample was prepared as a 7% loading of PAN on
alumina. A solution of 3.66 grams of PAN and 250 ml
dimethylformamide was added to 50.00 grams of alumina and
the suspension was allowed to stir overnight at 90C. The
dimethylformamide was removed by rotary evaporation and as
in the case of the 17% PAN on silica gel, the resulting
solid was caked up and very hard. The material was
pyrolyzed after grinding in a mortar and pestle.


32
PAN on Titanium Dioxide
The same general procedure was followed to prepare a
solid which was approximately 15% PAN. PAN (3.00 g) was
dissolved in 300 ml of N,N-dimethylformamide to which was
added 20 g of anatase (Ti02) and the resulting suspension
was stirred for 18 hrs. at 90C. The dimethylformamide was
removed by rotary evaporation and the caked up material was
ground in a mortar and pestle and subsequently pyrolyzed.
Metal Incorporation for Catalyst Studies
For the purposes of comparing the effects of various
metal additives on the catalytic activity of PPAN catalysts,
a different method of metal incorporation was used in an
attempt to have equimolar amounts of metal species per gram
of PPAN. Two series of catalysts were prepared using 3.45 x
10-4 moles of metal complex per gram of PPAN. Within each
series of catalysts, the PPAN starting material was all from
the same batch in order to eliminate differences due to
pyrolysis conditions. A weighed amount of metal complex was
dissolved in absolute ethanol and stirred with the
appropriate weight of PPAN to obtain a concentration of 3.45
x 10-4 moles metal/gram of PPAN. This suspension was
stirred overnight without heat, rotary evaporated and then
dried in a vacuum oven at room temperature.


33
Copper-Lithium Catalyst Preparation
PAN (20.00 g) was dissolved in 200 ml of
N,-N-dimethylformamide (DMF) with stirring and then 1.43 g
of LiCl was added. The solution was allowed to stir on low
heat for 18 hrs after which time the DMF was removed by
rotary evaporation. (Before rotary evaporation several
films were prepared by filling a small petrie dish about one
half full of the solution and drying in a vacuum
desiccator.) The resulting hard amber colored chunk of
plastic was ground up in a Waring Blender and then mixed
with 0.40 g of copper powder. The mixture was allowed to
sit with occasional stirring for about two weeks. During
this time, the color of the polymer slowly turned from amber
to green and the copper metal ceased to be observable. This
material was subsequently pyrolyzed as will be described.
Another batch of catalyst was prepared in which the copper
metal was added directly to the PAN-DMF-LiCl solution and
allowed to stir for about two weeks until all of the copper
was dissolved. The dark brown solution was rotary
evaporated to yield a dark brown plastic which was
subsequently ground up and pyrolyzed. A lithium chloride
catalyst was prepared in the same manner as the copper-
lithium catalyst except that copper was not added in this
case.


34
Ruthenium Catalyst
A solution of Ru3(CO)was prepared by dissolving 0.20
g of Ru3(CO)22 in 200 ml of toluene in a 500 ml erlenmeyer
flask. Powdered PAN (20.00 g, Aldrich) was added to this
solution which was allowed to stir on low heat for two days.
During this time the suspension became green looking. The
material was filtered using a glass frit resulting in a
light green polymer powder. The filtrate was bright orange
and virtually indistinguishable from the original solution.
Upon drying, the polymer powder became somewhat off white in
color. The FT-IR diffuse reflectance spectrum of this
powder indicated that it contained Ru3(CO)12; however the
color of the filtrate indicated that most of the Ru3(CO)12
remained in solution. This material was subseguently
pyrolyzed as will be described.
Methods
Pyrolysis Reaction
The pyrolysis reaction was carried out using a variety
of methods for two general reasons. First, as more was
learned about the pyrolysis reaction, modifications were
made to the apparatus in order to obtain more homogeneous
products. Second, the means of temperature control was
upgraded as funding allowed. The pyrolysis tubes were
constructed of 1.00 in. diameter pyrex tubing with a glass


35
frit at one end and a 24/40 ground glass joint at the other.
The ends were then tapered to about 1/4 inch diameter to
accommodate stopcocks and fittings for tygon hose
connections. Since highly poisonous gases were produced
during these reactions, all operations were carried out in a
fume hood. Using tygon tubing, the mineral oil bubblers
were attached to both ends of the pyrolysis tube to monitor
the carrier gas flow as well as for leak detection. A
schematic diagram of the apparatus is shown in Figure 7.
Initially a Fisher mechanical temperature controller was
used to monitor and control the temperature of the system
which was contained in a commercially available tube furnace
manufactured by Lindberg. A mercury-glass thermometer was
also used as a back-up for monitoring the temperature. It
was subsequently discovered that this temperature controller
had a fluctuation of approximately + 50C at any given
set-point. This proved to be unacceptable since it had been
demonstrated in other studies (and born out in this
investigation) that accurate temperature control in the
200-300C region of the reaction is crucial in controlling
the strong exothermic reaction which occurs in this
temperature region. Failure to control this exothermic
reaction results in a runaway reaction causing the PAN to
char very quickly with a dramatic loss in weight. To
circumvent this problem, an Omega CN-300 digital temperature
controller was purchased. Unfortunately, electrical


36
Figure 7:
Initial Pyrolysis Apparatus


37
problems were encountered in that the relay responsible for
turning the oven off and on had a propensity to stick in the
on position causing the oven to heat to 800-900C, resulting
in damage to the tube furnace and the pyrolysis tube. In
theory, this problem should not have occurred since the tube
furnace was rated to draw only 6 amperes while the
temperature controller was rated for a 10 ampere load. This
problem (which was in all probability caused by power surges
in the line) was alleviated by including a second relay, or
slave relay, rated for 25 amperes in the circuit. In this
configuration, the relay in the temperature controller
served only to switch the slave relay, which directly
controlled the oven. This served to reduce the current
passing through the temperature controller's relay, thus
increasing its lifetime. During the course of these
investigations, several other modifications were made to the
system. Initially, the pyrolysis tube and tube furnace were
positioned horizontally; however this proved to be a problem
since the pyrolysis reaction produced noxious liquids which
tended to settle towards the bottom of the tube and
impregnate the product. It was found that by placing the
apparatus in a vertical configuration, these liquids were
able to drain out of the pyrolysis tube into the tygon
tubing. Figures 7 and 8 schematically illustrate two
different versions of the pyrolysis apparatus positioned
vertically. Glass wool was packed into the ends of the tube


38
TEMPERATURE
Figure 8:
Modified Pyrolysis Apparatus


39
furnace to prevent heat loss. In addition, it was found
that positioning the control thermocouple outside the
pyrolysis tube resulted in a systematically higher
temperature reading. Therefore the thermocouple was
inserted directly into the PAN near the center of the tube
furnace as shown in Figure 8.
One problem with the Omega CN-300 temperature controller
was that it was manually controlled, which made the exact
duplication of the temperature-time profile rather
difficult. In addition, the temperature was raised
incrementally making the temperature-time profile a step
function. Since one of the goals of this investigation was
to compare catalysts produced using different pyrolysis
conditions, it was imperative that the temperature-time
profiles be highly reproducible. Towards the end of this
study, an Omega CN-2000 programmable temperature controller
was purchased which enabled exact duplication of a given
temperature-time profile. In addition, this controller was
capable of increasing the temperature continuously, thus
eliminating the need to raise the temperature incrementally.
During the course of the pyrolysis experiments, several
aspects of the pyrolysis conditions were varied to determine
the effect on the catalytic activity as well as the
elemental composition. All elemental analyses were
performed by the University of Florida Micro-analytical
Services operated by Melvyn Courtney. One line of


40
investigation involved using the same temperature-time
profile and using ammonia, carbon monoxide, air or nitrogen
as the carrier gas. In addition, variations in both the
rate of increase in temperature and the duration of heating
were examined with respect to the effect that these factors
have on the catalytic activity and elemental composition.
Catalytic Studies
As in the case of the pyrolysis apparatus, the
catalytic evaluation system was modified and upgraded
throughout the course of these investigations. What was
needed was a system capable of operating at atmospheric
pressure in the temperature range from 100 to 300C. In
addition, a high degree of reproducibility in the reaction
conditions was desirable to enable accurate comparison of
the catalytic activities of different catalyst formulations.
The initial screening studies employed a catalyst evaluation
system as depicted in Figure 9. A carrier gas (usually
nitrogen or air, both unpurified) flowed through a bubbler
containing neat substrate and the resulting gas stream was
assumed to be saturated in substrate. The reactor tubes
were approximately 1 cm I.D. with a sintered glass frit in
the center to support the catalyst. The heating system
consisted of a homemade tube furnace controlled by a
variable AC power supply and monitored with a mercury
thermometer located near the center of the tube furnace


41
Figure 9:
Initial Catalytic Evaluation System


42
outside the reactor tube. The tube furnaces were
constructed as in Figure 10 using pyrex or vycor tubing and
heavy gauge nichrome wire. The following substrates were
used in this system to determine if PPAN catalysts had any
activity towards them: CEES, DMMP, ethanol, methanol, CO/H2
(syngas), ethylbenzene, norbornene and 2-propanol. In the
case of syngas, the gases were introduced with a system of
two bubblers which enabled one to vary the ratio of carbon
monoxide and hydrogen. The system contained enough volume
for adequate mixing and the resulting mixture was passed
through an additional bubbler (corresponding to the
substrate bubbler in Figure 9) which was filled with mineral
oil and used to monitor the gas flow. The mineral oil
bubbler at the end of the system was used to monitor the
effluent flow as well as to aid in the detection of leaks in
the system. In the reactions involving CEES and DMMP,
Clorox bleach was used in the last bubbler to hydrolyze the
potentially harmful starting materials and reaction
products.
Although a system such as that illustrated in Figure 9
was useful for determining whether a particular catalyst had
any catalytic activity towards a certain substrate, there
were several inherent problems in the system which prevented
accurate catalyst activity comparisons. The two main
problems with this system were the inability to accurately
control and monitor the temperature and rate of influent


43
GLASS
Figure 10
Oven Design


44
supply. In addition, the amount of substrate that was
supplied by the bubbler was so small that it generally took
more than a week of continuous running to pass 1 ml of
substrate through the system. Figure 11 is a schematic
drawing of a catalytic evaluation system which was designed
to eliminate these difficulties. To overcome the
temperature control problem, a thermocouple well was built
into the reactor tube allowing accurate measurement of the
temperature near the catalyst bed. The K thermocouple was
connected to a digital, time-proportioning control, solid
state temperature controller (Omega CN 300) which enabled
the temperature to remain at + 2 C of the set point. It
was found that the temperature as measured by the
thermocouple in the thermocouple well was consistently
5-10C cooler than the temperature measured by the
thermometer outside the reactor tube. In addition, a
mechanical syringe was designed and constructed as
illustrated in Figure 12 in order to both stabilize and
increase the feed rate. Generally a 5 ml Hamilton Gas-tight
syringe was employed and the mechanical syringe, operating
with a 10 RPM motor, delivered 4.6 ml in approximately 7
hrs, making the Weight Hourly Space Velocity (WHSV) equal to
approximately 0.6 hour--'-. The WHSV denotes the ratio of the
mass flow rate of feed to the mass of the catalyst used as
defined by the following equation: WHSV= pV/W where p is
the density of the feed, W is the weight of the catalyst and


45
GAS IN -HZ
MEG HANICAL
SYRINGE
TUBE FURNACE
THERMOCOUPLE CO
/
EXTERNAL -1
RELAY
TEMPERATURE
CONTROLLER
DRY ICE/acetone
BATH
MINERALOI L
BUBBLER
SEPTUM
TYGON TUBE
LASS BEADS
CATALYST
FRI T
^*GAS OUT
MINERALOI L
BUBBLER
LIQUID
TRAP
Figure 11: Improved Catalytic Evaluation System


46
Figure 12: Syringe Design


47
V is the characteristic volumetric flow rate of the fluid.
The space time is defined as the reciprocal of the space
velocity. This corresponds to a very fast flow rate, almost
as high as used in many effecient, industrial processes.
Therefore the conversions obtained using this system (Figure
11) tended to be much lower than in the earlier catalytic
set-up illustrated in Figure 9.
Basicity Studies
Two types of basicity studies were attempted in this
investigation: the direct titration with dilute HC1 of
slurries of PPAN in distilled water, and the determination
of metal complex uptake from slurries of PPAN in agueous
solutions of several metal complexes using UV-vis or EPR
spectroscopy. In the direct titration method, a weighed
sample of PPAN was slurried in a 100 ml erlenmeyer flask
containing about 50 ml of distilled water (Millipore
Nanopure Water System) and titrated with a 0.01 molar
solution of HC1, using a Fisher pH meter to monitor the pH.
Since the titrations were often lengthy, the buret tip and
pH electrode were sealed to the flask using parafilm to
prevent evaporation. In the metal binding studies,
solutions of metal complexes were slurried with PPAN
preparations for about 24 hours and then washed and dried
under nitrogen. The dried PPAN was subseguently examined
in the EPR for evidence of metal binding. In another set of


48
experiments, a solution with a known absorbance was slurried
with PPAN and then filtered and washed. The filtrate and
washings were combined and reduced by blowing nitrogen to
the original volume of solution. The absorbance of this
solution was then checked by UV-vis and compared to the
absorbance of the original solution.
Photolysis Reactions
A Hanovia medium pressure mercury lamp was used as the
light source for photolysis reactions. Suspensions of PPAN
in various substrates were placed in quartz or pyrex
reaction vessels and stirred using a magnetic stirrer.
Photolysis products were analyzed by gas chromatography.
Thermal Analysis
Differential Scanning Calorimetry (DSC) and
Thermogravimetric Analysis (TGA) were performed by Ann
Livesey of the U.S. Army on a Dupont 9900 DSC and a Kahn
TGA. These analyses were carried out using commercially
available PAN supplied by Aldrich Chemical Company. In
addition, TGA and DSC results were obtained at the
University of Florida using a Perkin Elmer Series 7 Thermal
Analysis System.


49
Surface Area Measurements
A Micromeritics Digisorb 2600 was used to measure the
surface areas of many of the catalysts. In addition,
surface areas were also obtained courtesy of Dr. Willie
Hendrickson, 3M Corporation. The measurements were based on
the BET method. On analyses performed on the Digisorb 2600,
the samples were degassed for 12 hours at 90C, after which
time the samples were weighed. By weighing after degassing,
the contribution due to adsorbed water was minimized. The
sample weights were entered into the computer and the
analysis, calculations and report printout were performed
automatically by the instrument.
Gas Chromatography-Mass Spectrometry Analysis
GC-MS Analysis of the catalytic reaction products from
several substrates was performed by Dr. Dennis Rohrbaugh of
the U.S. Army.


RESULTS AND DISCUSSION
Pyrolysis Studies
One area of investigation in this study was the effects
that different pyrolysis conditions had upon the chemical
and catalytic properties of PPAN. Variations were made in
the temperature-time profile of the pyrolysis reaction as
well as the atmosphere under which the pyrolysis was carried
out in order to determine what effects these factors had
upon the elemental analysis and catalytic activity of these
preparations. Table 1 contains the elemental analysis
results obtained by varying the atmosphere under which the
pyrolysis was carried out. The last two entries differ from
the other samples in that they were pyrolyzed using a
programmable temperature controller. In addition, the
temperature programs for these two samples were identical;
therefore, the only difference between these two
preparations was the atmosphere under which the pyrolysis
was carried out. Figure 13 gives some representative
temperature-time profiles for some of the preparations
listed in Table 1. It should be noted that with the
exception of sample number 5, all of the manually controlled
pyrolysis reactions employed approximately the same
50


51
time (days)
Figure 13: Representitive Temperature-Time Profiles


52
Table 1: Elemental Analysis Results
Sample C% H% N% C/N Total Pyrolysis Maximum
ID Atmosphere Temp.(C)
Number
1*
52.41
2.63
24.44
2.14
79.48
air
300
2*
69.25
3.00
20.95
3.31
93.20
nitrogen
450
3*
59.47
2.40
23.21
2.56
85.08
air-250
N2-400
400
4*
74.11
0.87
17.19
4.31
92.17
nitrogen
800
5*
68.73
3.73
21.40
3.21
93.86
nitrogen
400(fast)
6*
70.25
4.00
21.82
3.23
96.07
ammonia
400
7*
69.99
3.15
22.38
3.13
95.47
ammonia
400
8*
69.76
2.70
21.05
3.31
93.51
CO
400
g**
69.39
3.78
21.59
3.21
94.76
nitrogen
400
10**
65.02
3.19
22.64
2.87
90.85
air
400
*manual temperature controller
**programmable temperature controller


53
temperature-time profile within the limits of experimental
error. Although at first glance one sees no apparent
relationships between the different preparations, several
general conclusions can be drawn. Reactions carried out in
an air atmosphere tend to have lower total percentages of
carbon, hydrogen and nitrogen than reactions carried out
under atmospheres of nitrogen, ammonia or carbon monoxide.
Presumably, this is due to increased oxygen incorporation
into the products of air pyrolyzed samples. In addition,
air pyrolyzed samples tend to have a lower percentage of
carbon than samples pyrolyzed in nitrogen, carbon monoxide
or ammonia. This effect appears more dramatic when one
looks at the C:N ratios of less than 3:1 for air pyrolyzed
samples while samples pyrolyzed in other atmospheres have
C:N ratios of greater than 3:1. Another general observation
is that all other factors being equal (temperature-time
profile, atmosphere), higher pyrolysis temperatures seem to
result in lower percentages of hydrogen in the final
product. These results are in general agreement with
reported results.3
Thermooravimetric Analysis Results
The TGA results also indicated that the pyrolysis
conditions are very important in determining the composition
of the final PPAN product. Figure 14 shows the effects of
isothermally heating at 280, 290 and 300C on the TGA


WEIGHT (%)
54
Samle: POLYACRYLONITRILE
Size: 10.32 mg
Method: RAMP/ISOTR 2B0 FOR SO
Comment: PT PAN / N2 100ML/MIN
TGA
105
65 _
60 r i i i t : : r t 1 i
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70
Time (min)
Figure 14: Isothermal Thermogravimetric Analysis


55
thermograms of PAN. The method used was to ramp the
temperature rapidly (20C/minute) to the given temperature
and then monitor the weight loss as a function of time. It
can be seen that the onset times for weight loss were
shorter and that the total weight loss was greater for
higher isothermal temperatures. Figures 15, 16, and 17 are
the results of temperature programmed TGA scans for
temperature ramping at 2.5C/minute, 5.0C/minute and
10C/minute, respectively. The typical sample size was
approximately 10.0 mg and the analyses were carried out in a
platinum boat under nitrogen flowing at 100 ml/minute.
These scans indicate that the weight remains essentially
constant until at least 280C; however, Figure 18, which is
a composite of Figures 15, 16, and 17, exhibits somewhat
anomalous behavior. Specifically, the scan corresponding to
a program rate of 5C/minute seems to exhibit rather
uncharacteristic behavior. For the purposes of catalyst
preparation, 500C is probably the maximum temperature of
interest and at that temperature one would expect the 5
C/minute scan to fall between the scans corresponding to 10
C/minute and 2.5C/minute. Instead, the scan at 5C/min
exhibits a smaller weight loss than either of the other two
scans. It is most likely that either an instrument
malfunction or an experimental error was the cause of this
discrepancy. If one looks at Figure 16, it can be seen from
the derivative curve that this sample is gaining weight from


56
Samle: POLYACRYLONITRILE
Size: 13.12 mg
Method: PAN 900 8 2.5
Comment: PT PAN / N2 100ML/MIN
TGA
Temgerature (*C)
Figure 15: Thermogravimetric Analysis, 2.5C/minute


57
Samle: POLYACRYLONITRILE
Size: 10.05 mg
Method: PAN 900 § 5
Comment: PT PAN N2 100 ML/MIN
TGA
L 16
3 >
c_
a)
1-5
Temperature (*C)
Figure 16: Thermogravimetric Analysis, 5.0C/minute


58
Sample:
Size:
Method:
Comment:
no
100-
qn -
30 -
c.
Cl
60-
50-
1
40 -)
J
30 -l
0
POLYACRYLONITRILE
10.02 mg
PAN 900 0 10
PT PAN / N2 100ML/MIN
TGA
- 2E
- 19£
CJ)
- 14>
c.
IS

- 10
- 5
- 2
100
200 300
400 5Qo"
600
700 800
T- -2
900
Temperature (*C)
General VI.OJ OuPont 9900
Figure 17: Thermogravimetric Analysis, 10.0C/minute


59
Sample: POLYACRYLONITRILE
Size: 10.02 mg
Method: PAN 900 @ 10
Comment: PT PAN / N2 100ML/MIN
TGA
File: A: PAN. 13
Ooerator: ABL
Run date: 01/13/S5 13:42
Figure 18
Composite of Temperature Programmed
Thermogravimetric Analysis


60
about 500C to about 700C. Since this is not apparent in
Figures 15 and 17, it is doubtful that the increase in
weight is due to nitrogen incorporation from the atmosphere.
Most likely, the atmosphere in the TGA chamber was
contaminated with oxygen, which was responsible for the
increase in weight. Unfortunately, it was impossible to
repeat these experiments to determine the cause of the
error. The data obtained before 325C appear to be
reliable, however, and several important conclusions can be
drawn about the pyrolysis reaction.
As earlier studies have indicated, the reaction begins
at higher temperatures when faster heating rates are used.
This has been attributed to the induction period for the
exothermic polymerization of the nitrile groups39> 40< 41
which can reach explosive speeds resulting in the
destruction of the polymer chain.42 Figure 19 contains the
TGA (obtained at the University of Florida) corresponding to
a temperature program rate of 0.5C/minute and the results
qualitatively agree with the scans done at faster heating
rates. The onset temperature for weight loss is
considerably less than in the scans shown in Figure 18. In
addition, the percentage weight loss is also less than for
the other runs; however, since this run was carried out
using a different instrument, the absolute percentages
should not be taken too literally since the instruments may
not have been identically calibrated. Figure 20, from some


61
Sanle Weight* 0.742 mg
poiyacryionltrile
PERKIN-ELMER
7 Senes Thermal Analysis System
Figure 19: Thermogravimetric Analysis, 0.5C/minute


sample weight
62
Figure 20:
Thermogravimetric Analysis and
Differential Thermal Analysis


63
work by Grassie and McGuchan, shows Differential Thermal
Analysis (DTA) curves for polyacrylonitrile heated for
10C/minute in air, nitrogen and vacuum.43 It can be seen
from Figure 20 that air pyrolysis results in greater weight
loss than nitrogen pyrolysis, while the greater weight loss
for the vacuum pyrolysis can be attributed to the ease of
volatilizing high-boiling fractions. In addition, it can be
seen that the onset temperature for weight loss is lowest
for vacuum and lower for air than for nitrogen. This would
seem to suggest that there are processes occurring in the
air atmosphere at lower temperatures than in the nitrogen
atmosphere. Most likely, the lower temperature reactions
are due to oxygen incorporation reactions not occurring
during the nitrogen pyrolysis. The Differential Thermal
Analysis (DTA) curve in Figure 20 is also interesting in
that it is tilted due to the fact that the sample is giving
off so much heat that the system becomes hotter than the
programmed rate of temperature rise (the DTA instrument
measures the difference in temperature between a reference
and a sample when both are heated under identical
conditions). It should be noted that these results have
severe implications for the pyrolysis of bulk samples of PAN
due to the exothermicity of the nitrile polymerization
reaction. The destructive reaction could occur at
relatively low temperatures if the heat of polymerization is
not dissipated rapidly enough to prevent the interior of the


64
polymer from reaching the critical temperature.42 Since the
catalyst preparation method used in this study employed a
pyrex tube containing about 15 grams of PAN, heat
dissipation was potentially a serious problem.
Differential Scanning Calorimetry Results
The DSC results also suggested that there is some sort
of an induction period associated with the pyrolysis
reaction. Figure 21 shows the effects of heating rate on
the exotherms. As in the temperature programmed TGA
analysis, the onset temperatures increase with the heating
rate. If there were no induction period, one would expect
the onset temperature to be lower with increasing heating
rate. A comparison between the onset temperatures for the
DSC and TGA results of temperature programming at 2.5, 5.0,
and 10C per minute demonstrates that there is fairly good
correlation between the onset temperatures for weight loss
and the onset temperatures for the exotherm, implying that
the reaction resulting in weight loss is an exothermic
reaction. The shapes of the exotherms in Figure 21 are
quite similar but the curves differ in the onset
temperatures and the amount of heat given off. Figure 22
contains the DSC results for a heating rate of 0.5C/minute,
and again the results qualitatively agree with those for
faster heating rates in that the onset temperature for heat
loss is less while the amount of heat loss seems to be


MUAL/SEC
65
187 207 227 2<7 267 287 307 317
TEMPERATURE (C)
SC
86/B1/08 TIME: 14:03
Figure 21:
Temperature Programmed Differential
Scanning Calorimetry


Heat Flow (mW)
66
PERKIN-ELMER
7 Series Thermal Analysis System
Temperature (C)
Figure 22
Differential Scanning Calorimetry, 0.5C/minute


67
reduced with respect to the results obtained at faster
heating rates. Again, direct comparisons should not be
taken too literally due to the fact that two different
instruments were used to obtain these results. In general,
it seems that faster heating rates result in a greater
amount of heat being given off, correlating with the TGA
results which indicated that more weight was lost with
faster heating rates. It is also interesting to note that
there seems to be an endothermic process occurring at higher
temperatures than the exothermic process. In the exothermic
process, the amount of heat absorbed also seems to slowly
increase with increasing temperature, although this could be
an instrumental artifact. Figure 23 shows the effects of
isothermally heating at 240, 245, 250, 255, and 260C. The
samples were heated at the rate of 1C per minute until
reaching the desired temperature and then maintained at that
temperature for a period of time. The isothermal DSC
results qualitatively agree with the isothermal TGA results
in that lower isothermal temperatures result in longer onset
times. Also, the amount of heat given off increases with
increasing temperature, in agreement with TGA results which
showed that higher isothermal temperatures resulted in
greater weight loss. In addition, the shape of the exotherm
varied with isothermal temperature in that higher isothermal
temperatures resulted in larger, sharper exotherms. As in
the temperature programmed DSC runs, there seems to be an


68
247 257 267 277 287 297 307 317 327
TEMPERATURE CO
DSC
Figure 23: Isothermal Differential Scanning Calorimetry


69
endothermic process occurring after the exothermic process
and the amount of heat absorbed increases with increasing
isothermal temperature. Again it is not known at this time
whether this is an instrumental artifact. From the thermal
analysis results, several conclusions about the pyrolysis
reactions can be drawn. The reactions occurring during the
pyrolysis are dominated by a strongly exothermic reaction
concomitant with a dramatic loss in weight, preceded by an
induction period. The onset times for weight loss and heat
loss are strongly dependent upon the rate of heating and the
absolute temperature. In addition, there appear to be some
inconsistencies in treating this pyrolysis reaction as
merely a nitrile polymerization reaction since the dramatic
weight losses observed in this study suggest that a
considerable amount of chain destruction is taking place.
Earlier studies on a commercial acrylic fiber containing
methyl acrylate and acrylonitrile found a high correlation
between the nitrile content and the heat evolved during the
pyrolysis.44 The amount of unreacted nitrile groups was
determined by infrared analysis before and after pyrolysis
(Figure 24). The results of these analyses suggest that the
exothermic reaction is associated with the disappearance of
the nitrile groups, although it seems like an
oversimplification to attribute this reaction to a simple
nitrile polymerization reaction in the case of a
polyacrylonitrile homopolymer. In a copolymer containing


Nitrile Content, % (2240 cm
70
'00 r
i I
/ X-
ow ^
SC
.O'
' o
o o
60
I
c jr*'
4, y > G.
40 ^ o 5
> =
o.
ZC c
I I
40 50 SC 100 120 140 160
Residual Exotherm (cal/gram)
Figure 24: Plot of Residual Nitrile Content Versus
Area of Exotherm


71
relatively unreactive polymer subunits interspersed with the
nitrile containing subunits, it might be possible to view
the pyrolysis as a simple nitrile polymerization reaction.
If one calculates the theoretical weight loss associated
with the reaction in Figure 1, the proposed catalyst
formation reaction, the result is about 5% weight loss
principally due to hydrogen loss during cyclization and
aromatization. This does not account for the fate of
inevitable chain ends, defects and polymerization catalyst
residues which are present to various extents depending upon
individual sample preparation techniques. The weight losses
obtained for isothermal runs at 280, 290, and 300C were
about 25, 30, and 35%, respectively (Figure 18), after 1
hour of heating at these temperatures. These percentages
seem far greater than what one would expect from the nitrile
polymerization reactions, even with the contributions from
the reactions undergone by the defects and impurities
previously mentioned. Furthermore, the nitrile
polymerization reaction should theoretically produce no
nitrogen containing volatiles; however, these were observed
in this as well as in all other studies. Although it is
conceivable that the nitrogen could have originated from
atmospheric sources, this is unlikely since the nitrile
moiety is far more reactive than the dinitrogen molecule.
In conclusion, while these and other results indicate that
the disappearance of the nitrile functional group is


72
intimately related to the exothermic reaction(s) occurring
during the pyrolysis, it is unlikely that a simple
intramolecular nitrile polymerization-aromatization is the
only reaction occurring to a significant extent. The
polymers used in this and other studies were atactic
polymers, making the likelihood of extensive intramolecular
cyclization more remote than in an isotactic polymer. In
reality one must view the nitrile groups as being randomly
oriented about the polymer backbone. This conformation
would increase the probability of reactions between adjacent
chains resulting in extensive crosslinking. The observation
that polyacrylonitrile becomes harder and more brittle after
heating supports this interpretation. Although the thermal
analysis studies, which employed about 10 mg of sample, may
not be directly applicable to the bulk pyrolysis of 15 to 20
grams of material, the results indicate the there may be a
serious design flaw in the pyrolysis apparatus used in these
studies in that the dissipation of the heat produced in the
exothermic reaction is crucial in preventing the destructive
runaway reaction. The apparatus employed about 20 grams of
PAN in a 1 inch diameter tube and had typical gas flow rates
of less than 100 ml/minute. The combination of rather
densely packed material and relatively slow gas flow rate
may have resulted in a heat dissipation problem. Although
the pyrolysis apparatus probably was not producing the best
catalysts possible, it proved to be adequate for the


73
purposes of this study. Some catalytic material was
prepared in our controlled temperature pyrolysis since
little or no activity was observed for materials heated too
quickly.
Surface Area Results
The results of the BET surface area determinations are
contained in Table 2. Since these results were determined
using different instruments and conditions (degas time,
degas temperature, etc.), these factors should be taken into
account when comparing the data. Several conclusions may be
drawn from these results. In general, the surface areas of
the pure PPAN materials are quite low, on the order of 0-9
square meters per gram. These results are considerably
lower than those reported by other workers who have obtained
values of 19.0-19.2 m2/gram,19 17.7 m2/gram15 and 18
m2/gram.16 These discrepancies can in part be explained
when one observes that the surface area of the PPAN'CuLi
catalyst increases after passing 15 ml of ethanol over two
grams of catalyst. In addition, it was observed that the
products collected from the catalytic reaction frequently
were amber colored, indicating that some of the waxes and
oils produced in the pyrolysis reaction had impregnated the


Table 2: Surface Area Results
Sample
Max
BET Surface
C
H
N
Comments
Temp(C)
Area (M2/g)
(%)
(%)
(%)
PPAN on
Silica3
400
260 (272)c
high load
PPAN on
Silica3
400
243 (263)c
2.55
0.34
0.66
low load
Silica3

(332)c
0.20
0.32
0.43
PAN

(286)c
67.12
5.82
26.17
PPAN3
400
5 (6)c
65.02
3.19
22.64
air pyr.
PPAN3
400
7 (9) c
69.39
3.78
21.59
nit. pyr
Alumina3

153c
0.03
0.24
0.00
PPAN on
alumina3
400
134 (13 4)c
3.55
0.52
0.73
nit. pyr.
PPANb
400
>lc
59.47
2.40
23.21
air pyr.
PPANb
450
4C
69.25
3.00
20.95
nit. pyr.
PPAN Cu-Lib
500
>lc
60.20
4.07
18.26
freshd
PPAN Cu-Lib

lc



spente
aThese samples were all pyrolyzed using an identical program
on an Omega CN 2000 temperature controller.
bThese samples were pyrolyzed using an Omega CN 300 manual
temperature controller.
cThese surface areas were generously provided by Dr. W.
Hendrickson of the 3M corporation while the other numbers were
obtained using a Micromeritics Digisorb 2600 surface area
analyzer.
dAfter pyrolysis.
eAfter passing 15 ml of ethanol over 2.00 g of this material
at 300C.


75
catalyst, possibly filling up some of the pores and reducing
the surface area. The passage of ethanol through the
material may serve to wash out these residues and increase
the surface area. Granted, the surface areas of these two
samples are very small and possibly within the error limits
for this instrument. However, both samples were run on the
same instrument and although the absolute difference in the
"spent" and "fresh" catalyst samples is relatively small,
the surface areas differ by a factor of five. Degannes and
Ruthven19 reported that the surface areas of fresh and used
catalysts remained identical within the limits of
experimental error. The studies which reported surface
areas in the neighborhood of 20 m2/gram all used a similar
pyrolysis method in that they spread the PAN thinly on trays
and carried out the reaction in a draft oven. This
pyrolysis method is probably superior to the method used in
this investigation for two reasons. First, there is a
greater capacity for heat dissipation in a draft oven;
second, with the PAN spread in a thin layer is less of a
propensity for the pores on the surface to become clogged
with low molecular weight residues from the pyrolysis
reaction, thus lowering the surface area.
The surface area results for the silica and alumina
supported PPAN catalyst preparations were approximately 250
and 130 m2/gram, respectively, indicating that the surface
areas of these supports were not significantly altered by


76
the addition of up to 17% PAN by weight. These surface
areas were more than an order of magnitude greater than the
largest of the unsupported PPAN catalyst preparations, and
these results will be discussed later with respect to
catalytic activity comparisons between different
preparations.
Basicity Studies
Titration Results
Several attempts were made to determine the number of
basic sites on the surface of PPAN catalysts by directly
titrating with dilute HCl. The PPAN samples were either
used as is after the pyrolysis reaction or slurried with
concentrated sodium hydroxide, filtered and then washed with
distilled water. For a number of reasons it was impossible
to obtain quantitative data from this titration method. For
some unknown reason, the pH meter would not stabilize on a
reading in suspensions of PPAN materials. The meter
readings would continually randomly jump around when
immersed in PPAN suspensions but would stabilize immediately
upon being immersed in a solution not containing PPAN. In
addition, the meter would drift up as much as 4-5 pH units
over a period of several hours, making the time at which the
reading was taken very important. Throughout the titration,
the calibration of the pH meter was checked and found to be


77
accurate. Therefore, the instability in the readings was
caused by the PPAN itself, suggesting that some unknown
reaction was occurring at the electrode and/or on the
surface of PPAN. These results are in general agreement
with the only other titration study which demonstrated that
long-time exposure to dilute acid will result in some acid
uptake, indicating a generally basic nature.36 Although, as
stated previously, no quantitative results could be obtained
from this titration study, it is possible to make some
general observations about the basic properties of PPAN
materials. Although PPAN is essentially insoluble in all
acids, bases and solvents, it is by no means inert, as
evidenced by the instability of the pH electrode and the
slow uptake of acid. In addition, PPAN samples pyrolyzed in
air seem capable of absorbing more acid than samples
pyrolyzed in a nitrogen atmosphere.
Metal Binding Studies
As in the case of the titration studies, the results of
the metal binding studies were somewhat inconclusive in that
no direct evidence of metal binding was ever established.
The general method used was to add an EPR active metal
complex to PPAN by adsorption from solution and then look
for nitrogen hyperfine in the EPR. It was assumed that if
the polypyridine type structure in Figure 1 is the basic
site, the metal species would bind to the nitrogen and the


78
unpaired electron on the metal ion would interact with the
nuclear spin of the nitrogen (3/2), resulting in a nuclear
hyperfine interaction. Copper and cobalt were chosen as the
metal species since nitrogen hyperfine is known to be
readily observable for both at liquid nitrogen temperatures.
There have recently been several studies on the use of 1,8-
napthyridine based ligands in the formation of metal
complexes and these studies indicate that the napthyridine
moiety can accommodate several bonding modes, as shown in
Figure 25. Unlike bipyridines, the nitrogens in 1,8-
napthyridines are rather close together, resulting in a
significantly reduced "bite angle", thus making
napthyridines much poorer candidates for bidentate ligands.
In addition, its steric bulk can make the napthyridine
moiety a poor candidate for a monodentate ligand in certain
instances. There are numerous examples in the literature of
napthyridine based ligands forming complexes with many of
the metals in the periodic table, including the lanthanides
and the rare earth metals.45-50 The three general bonding
modes for these are depicted in Figure 25. These materials
are frequently highly colored due to the presence of a metal
to ligand charge transfer band, and have also been known to
exhibit fluxional behavior.51'52
i


79
Figure 25: Naptheridine Binding Modes


80
Diphenvlglvoximato cobalt (II) (Co(DPGH)2
This compound was prepared and stored in a nitrogen
atmosphere. The EPR samples were prepared in a glove box
using a 50:50 toluene:methylene chloride mixture which had
been exposed to three freeze-pump-thaw cycles to remove
dissolved oxygen. A saturated solution of Co(DPGH)2 was
slurried with a small amount of PPAN, the mixture was
filtered and EPR samples were prepared from the filtrate and
the solid material. The EPR spectra of solid samples of
Co(DPGH)2 and PPAN after being exposed to Co(DPGH)2 in
solution are given in Figure 26. The PPAN sample was run at
a higher sensitivity to detect any small changes in the
spectrum upon addition of the cobalt complex. The spectra
demonstrate that PPAN has little or no effect upon the EPR
spectra for Co(DPGH)2, other than the superimposition of the
narrow signal at about g=2 which is characteristic of PPAN.
Both the lineshape and g value of the cobalt complex remain
unchanged upon addition of PPAN. It should be noted that
solutions (50:50 toluene:methylene chloride mixture) of
Co(DMGH)2 and its pyridine adduct were prepared and the EPR
spectra were consistent with earlier published results which
showed the typical eight line cobalt spectrum with five line
nitrogen hyperfine in the case of the pyridine adduct
(Co(DPGH)2'2Pyridine).38 These results suggest that the
electronic structure of the cobalt is not perturbed upon
addition of Co(DPGH)2, implying that the Co(DPGH)2 present


81
3300 GAUSS
Electron Paramagnetic Resonance Spectra of
PPAN and Co(DMGH)2
Figure 26:


82
in the sample is merely physically adsorbed and not
chemically bound to the surface. Since pyridine binds to
this complex, the results show that PPAN is a poorer base
than pyridine, although at this time it is not possible to
discern whether the chemical structure, electronic structure
or steric bulk is responsible for PPAN's poor binding
ability. Since Co(DPGH)2 is a square planar complex with a
nearly planar ligand, the steric requirements for axial
adduct formation are not particularly stringent. This
suggests that the steric bulk of the PPAN is probably not
the over-riding cause for the poor binding ability of PPAN
materials.
Cupric chloride (CuCl2)
Cupric chloride (0.20 g) was added to a strongly basic
suspension of PPAN in distilled water and the resulting
mixture was allowed to stir for several hours, after which
time the mixture was filtered and the PPAN*CuC12 washed with
copious amounts of distilled water. The EPR spectra of this
material as well as that of solid CuCl2 are presented in
Figure 27. The EPR of PPAN'CuC12 consists of the usual PPAN
signal at about g=2, with another signal appearing at
slightly greater than g=2 as well as a much smaller but
discernable signal at about g=2.5. After washing with 0.1
molar HCl and copious amounts of distilled water the signal
remained unchanged. Since a strongly basic solution could


83
Figure 27: Electron Paramagnetic Resonance Spectra of
PPAN and Copper (II) Chloride


84
result in the formation of Cu(OH)2 on the surface of the
PPAN, it was hoped that washing with hydrochloric acid would
result in the protonation of any remaining basic sites and
the formation of chloride salts on the surface of the PPAN.
A comparison of the EPR spectra of solid cupric chloride and
PPAN*CuC12 reveals that the PPAN*CuC12 spectrum is not a
simple addition of the spectra for CuCl2(s) and PPAN as in
the case of Co(DMGH)2. Although the characteristic PPAN
signal appears unchanged, the copper signal in PPANCuC12 is
markedly different from that of CuCl2(s), implying that
there is some sort of electronic interaction between the
PPAN and copper's unpaired electron, unlike the case of
Co(DMGH)2. Although the g values and lineshapes of the
copper signals are quite different, there is no evidence of
a nitrogen hyperfine interaction. Bai and coworkers20
prepared silica gel supported PPAN materials and studied
their physical and catalytic properties. Their work
presented XPS results for CuCl, PPAN on silica gel and PPAN
on silica gel with added CuCl, and they concluded that the
copper was coordinately bound to nitrogen based on
differences in the Cu 2p 3/2 and N Is binding energies upon
the adsorption of CuCl onto silica supported PPAN.20
Different preparative methods or the presence of silica gel
makes direct comparison of these results impossible. As
mentioned previously, the nitrogens in PPAN may not be
capable of producing a nitrogen hyperfine interaction, thus


85
making the EPR method unsuitable for determining whether the
copper is bound to nitrogen. In any case, the fact that the
copper signal in PPAN*CuC12 was significantly different from
that of CuCl2, coupled with the fact that repeated washing
of the PPAN with HC1 and distilled water failed to alter or
reduce the intensity of the signal, implies that the copper
was indeed coordinated to the surface of the PPAN, although
it was not possible to determine whether the copper was
bound to nitrogen, carbon, or oxygen.
Bis (2,2'-bipyridine) ruthenium (II)
trifluoromethanesulfonate (Ru (bipy)2(CF3SO3)2)
This compound was prepared in situ by dissolving 0.11 g
of AgCF3S03 in an ethanol solution containing 0.10 grams of
Ru(bipy)2Cl2, heating the reddish solution to boiling to
coagulate the AgCl, and finally filtering and washing with
absolute ethanol. The deep red filtrate was presumed to
contain Ru(bipy)2(CF3SO3)2. The UV-visible spectrum of a
solution made in a volumetric flask was recorded before and
after the addition of a large excess of PPAN (5.10 grams) It
was found that there were no appreciable absorptivity
changes in the UV-visible spectra, neither in the wavelength
maxima or intensities of the absorptions, indicating that
there was negligible adsorption or binding of
Ru(bipy)2 (CF3SO3)2 to the surface of the PPAN. This
particular ruthenium complex was chosen for several reasons.


86
Being a very poor ligand and hence a good leaving group, the
trifluoromethanesulfonate anion seemed like a good candidate
for being displaced by PPAN, which had not proved to be a
particularly good ligand in previous experiments. In
addition, it was hoped that the presence of two bidentate
2,2'bipyridines bound to the ruthenium would facilitate (by
the "chelate effect") the binding of PPAN as a bidentate
ligand, thus lessening the steric strain associated with the
large PPAN structures. This strategy proved to be
ineffective since no adsorption of ruthenium complex was
detected.
Catalytic Results
The catalytic studies fall into two general catagories:
screening reactions were carried out to determine whether
PPAN catalysts were active towards various substrates, and a
specific reaction, that of ethanol over PPAN catalysts, was
used as a model reaction to study the effects of pyrolysis
conditions, metal dopants and oxide supports on the
catalytic activity of PPAN.


87
Screening Reactions
Isopropyl alcohol and ethylbenzene
Some preliminary experiments with early catalyst
preparations were undertaken to verify the catalytic
activity reported in previous studies. Manassen and
Wallach15 had reported the catalytic activity of PPAN
materials towards isopropyl alcohol. Therefore, the vapor
was passed over PPAN at 150C and although the results were
not quantified, acetone was detected by gas chromatography.
As a preliminary probe of the photocatalytic capabilities of
PPAN, suspensions of PPAN in isopropyl alcohol were
irradiated with visible and ultraviolet light using a medium
pressure mercury vapor lamp with pyrex and quartz reaction
vessels, respectively. The results, when compared to the
appropriate blank runs (no PPAN present), indicated that
PPAN had no appreciable photocatalytic activity towards
isopropyl alcohol.
The same series of experiments were carried out using
ethylbenzene as the substrate and the results were analogous
to those obtained for isopropyl alcohol. In the vapor
phase, ethylbenzene was converted to styrene, as previously
reported by Degannes and Ruthven,19 and no photocatalytic
reaction was observed to occur. The results of these
experiments served to verify that although the measured
surface areas were somewhat smaller, the catalytic


88
activities of the materials prepared in this laboratory were
comparable to those of earlier published reports.
Norbornene oxidation
It has been previously reported that PPAN is an active
dehydrogenation catalyst, however there have been very few
studies in which PPAN has been tested with respect to its
oxidative capabilities. Towards this end norbornene was
used as a gaseous substrate at 140C and atmospheric
pressure, as well as at 50C and 35 psi oxygen in a
suspension containing 1 gram of norbornene and 1 gram of
PPAN in 20 ml of acetonitrile. In both cases, there were no
detectable changes in composition under the reaction
conditions as measured by gas chromatography using a
carbowax column and a flame ionization detector. There are
several possible products which could conceivably be formed
from norbornene, namely norbornene oxide, norbornadiene, and
numerous products resulting from the fragmentation of the
bicyclic ring system. Norbornadiene would result from the
dehydrogenation of the carbon atoms symmetrically related by
a mirror plane to the doubly bonded carbons in norbornene.
Since PPAN has been shown to dehydrogenate cumene and
ethylbenzene, which both have a bulky aromatic system
conjugated to the bond being dehydrogenated, it was felt
that norbornene might be a good candidate to test whether
the existence of a 7r-allyl type system is necessary for


89
dehydrogenation to occur. On the other hand, one must
remember that norbornene, due to its bicyclic structure, has
a far greater steric problem than cumene or ethylbenzene.
If steric factors are of the most significance, norbornene's
lack of reactivity is readily explained since all of the
previously proposed structures for PPAN catalysts are quite
large and bulky, which would render them less suitable for
interaction with a sterically hindered substrate such as
norbornene. The conversion of norbornene to norbornene
oxide has been accomplished using ruthenium (II)
phenanthroline catalysts.53 The fact that norbornene is
stable in the presence of PPAN implies that either
norbornene is too sterically hindered to react with PPAN, or
PPAN is not a particularly good oxidation catalyst or both.
These rationalizations are supported by the available
literature reports since the use of PPAN as an oxidation
catalyst has been reported only twice (in one case, silica
and metal salts were also present, while the other report
employed PAN for the purpose of silver crystallite
deposition on the surface of oxide supports21) and the
polymeric structure of PPAN has been fairly well
characterized as being a large extensively crosslinked
heterocyclic structure.


90
Syngas conversion
The hydrogenation of carbon monoxide with dihydrogen is
an industrially important reaction since mixtures of
hydrogen and carbon monoxide, called syngas, are a major
by-product of the petroleum mining and refining industry.
Unlike carbon dioxide, carbon monoxide is fairly reactive
due to its valence deficient electronic structure, as
evidenced by its tendency to form numerous metal-carbonyl
complexes. The catalytic mechanism of PPAN has been
described in the literature as one in which the catalyst
dehydrogenates a substrate and subsequently becomes reduced.
The catalyst is oxidized to its original state by a short
treatment in oxygen at reduced temperatures and has been
shown to produce water. Therefore, if carbon monoxide were
to participate in this catalytic cycle, it would have to
replace the oxidant since a product further reduced would be
undesirable except for emission control applications. By
using a mixture of hydrogen and carbon monoxide it was hoped
that the PPAN could adsorb the hydrogen molecules and then
catalytically reduce the carbon monoxide, completing the
catalytic cycle. A 1:2 mixture of carbon monoxide and
hydrogen was passed over a nitrogen pyrolyzed sample of PPAN
at temperatures up to 300C and the only product observed
using gas chromatography was a small methane impurity which
was contained in the original carbon monoxide. Since
ruthenium is frequently used as a Fischer-Tropsch catalyst,


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J



THE PREPARATION AND PROPERTIES OF PYROLYZED
POLYACRYLONITRILE CATALYST MATERIALS
By
JEFFREY LEE CLARK
A DISSERTATION PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN
PARTIAL FULFILLMENT OF THE REQUIREMENTS
FOR THE DEGREE OF DOCTOR OF PHILOSOPHY
UNIVERSITY OF FLORIDA
1987

DEDICATION
To IC, who almost never lost the faith.

ACKNOWLEDGEMENTS
There are many people who facilitated this work through
actual assistance as well as thoughtful discussion and
encouragement, and although there are too many to mention,
their contribution is not overlooked. I am especially
appreciative to the glass, machine and electronics shops for
their invaluable assistance in constructing and repairing
the equipment used in this study. I would also like to
thank Dr. Willie Hendrickson of the 3M corporation, Ann
Livesey of the U.S. Army, and Ngai Wong of the University of
Florida for their generous assistance in this work. In
addition, I am grateful to the University of Florida, the
U.S. Army, and Geo-Centers Inc. for financial support during
this endeavor. Finally, I am deeply indebted to Dr. Russell
S. Drago for the freedom and support he provided me during
this study.
in

TABLE OF CONTENTS
Page
ACKNOWLEDGEMENTS iii
LIST OF TABLES v
LIST OF FIGURES vi
ABSTRACT viii
INTRODUCTION 1
BACKGROUND 4
EXPERIMENTAL 2 8
Materials 28
Preparations 29
Methods 3 4
RESULTS AND DISCUSSION 50
Pyrolysis Studies 50
Basicity Studies 76
Catalytic Results 86
SUMMARY AND FUTURE RECOMMENDATIONS 132
APPENDIX: XPS SPECTRA 138
REFERENCES 156
BIOGRAPHICAL SKETCH 161
IV

LIST OF TABLES
Page
1. Elemental Analysis Results 52
2. Surface Area Results 74
3. CEES Products 98
4. XPS Results; Elemental Ratios 108
5. XPS Results; Ionization Energies 109
6. Effects of Adsorbed Metal Salts; Series 1 114
7. Effects of Adsorbed Metal Salts; Series 2 117
8. Miscellaneous Catalyst Results 119
9. Glass Spheres Results 122
10. Elemental Analysis Results For Suported
Catalysts 125
11. Catalytic Results For Supported Catalysts 127
12. GC-MS Results for Discolored Products 130
v

LIST OF FIGURES
Page
1. Proposed Pyrolysis Reaction 6
2. Imine-Nitrone Structure 7
3. Proposed Catalytic Mechanism 11
4. Dehydrogenation of Acridine 13
5. Possible Dehydrogenation Mechanisms 14
6. Proposed Structure and Pyrolysis Product for
Polycyanoacetylene 16
7. Initial Pyrolysis Apparatus 3 6
8. Modified Pyrolysis Apparatus 38
9. Initial Catalytic Evaluation System 41
10. Oven Design 43
11. Improved Catalytic Evaluation System 45
12. Syringe Design 46
13. Representitive Temperature-Time Profiles 51
14. Isothermal Thermogravimetric Analysis 54
15. Thermogravimetric Analysis, 2.5°C/minute 56
16. Thermogravimetric Analysis, 5.0°C/minute 57
17. Thermogravimetric Analysis, 10.0°C/minute 58
18. Composite of Temperature Programmed
Thermogravimetric Analysis 59
19. Thermogravimetric Analysis, 0.5°C/minute 61
20. Thermogravimetric Analysis and Differential
Thermal Analysis 62
vi

21. Temperature Programmed Differential Scanning
Calorimetry 65
22. Differential Scanning Calorimetry, 0.5°C/minute 66
23. Isothermal Differential Scanning Calorimetry 68
24. Plot of Residual Nitrile Content Versus Area
of Exotherm 7 0
25. Naptheridine Binding Modes 79
26. Electron Paramagnetic Resonance Spectra of PPAN
and Co(DMGH)2 81
27. Electron Paramagnetic Resonance Spectra of PPAN
and Copper (II) Chloride 83
28. Polymerization Mechanisms 94
29. Structures of Nerve Gases and
Dimethyl-methylphosphonate 103
30. Selectivities of Metal Doped Catalysts 133
Vll

Abstract of Dissertation Presented to
the Graduate School of the University of Florida
in Partial Fulfillment of the Requirements for
the Degree of Doctor of Philosophy
PREPARATION AND PROPERTIES OF PYROLYZED
POLYACRYLONITRILE CATALYST
MATERIALS
By
Jeffrey Lee Clark
December, 1987
Chairman: Russell S. Drago
Major Department: Chemistry
The preparation and properties of pyrolyzed
polyacrylonitrile (PPAN) catalyst materials was
investigated. Specifically, the effects of variations in
the pyrolysis conditions (heating rate, maximum temperature,
atmosphere, etc.) were examined using elemental analysis,
Differential Scanning Calorimetry and Thermogravimetric
Analysis. In addition, the effects of metal additives and
oxide supports on the catalytic activity of PPAN catalysts
were studied using the dehydrogenation of ethanol as a test
reaction. It was found that small amounts of metal
additives were capable of profoundly affecting both the
activity and selectivity of PPAN catalysts. Supporting PPAN
on oxide supports served to increase the surface area by
Vlll

greater than an order of magnitude. However, the catalytic
activity seemed to be more a reflection of the catalytic
activity of the support used because blank runs using only
the support resulted in very similar selectivity and
activity. In addition, PPAN catalysts were found to
decompose chloroethyl ethylsulfide at temperatures as low as
200°C under both aerobic and anaerobic conditions.
IX

INTRODUCTION
In the last twenty years, due to the increased
consumption of manmade materials and energy, the area of
catalysis has blossomed into a science in its own right.
Due to its more general application to spectroscopic
techniques, homogeneous catalysis has traditionally been
better understood than heterogeneous catalysis, which is
less amenable to spectroscopic techniques, making study
more difficult. In fact, it has been frequently stated by
industrial catalyst chemists that heterogeneous catalyst
preparation is more an art than a science. However, recent
advances in surface analytical techniques have enabled
chemists to more fully understand the reactions occurring on
the surfaces of heterogeneous catalysts. At present,
though, the high relative cost of these techniques is
frequently prohibitive in many academic environments.
Metals and their complexes and salts have traditionally
been the catalysts of choice industrially. Indeed, one of
the first catalytic chemical plants (excluding biological
"plants" of course) was built by Germany during World War II
1

2
and utilized a catalyst composed of mixed metal oxides to
convert CO and H2 (syngas) into liquid fuel using the
Fischer Tropsch Process.1 Both the petroleum and polymer
industries initially led research efforts in catalysis. With
the advent of automobile exhaust emission controls, however,
extensive research was dedicated to developing effective
emission control catalysts.2 In many cases, the catalysts
employed in industrial operations consist of small metal
crystallites or complexes supported on the surface of an
inorganic oxide (such as silica gel (Si02), titanium dioxide
(Ti02) and alumina (A1203)), whose primary function is to
uniformly disperse the active catalyst species. This
provides greater efficiency per unit of metal species by
increasing the catalytic surface area available for chemical
reaction. Due to the high cost of precious and strategic
metals, interest has been directed towards less expensive
alternatives to precious metal catalysts. One area of
investigation has involved the use of heat stable organic
materials. In particular, organic pyropolymers have been
examined as possible catalyst candidates due to their
paramagnetism, semiconducting tendencies and heat stability.
One such polymer, pyrolyzed polyacrylonitrile (PPAN), has
been studied extensively as a catalyst for several
reactions. At present there are very few published reports
on the effects of metal additives or of supporting PPAN
catalysts on inorganic oxide supports. This study involves

3
a systematic investigation into the effects of variations
the preparation of PPAN catalyst materials.
in

BACKGROUND
It was reported in 1958 by Burlant and Parsons3 that
upon treatment with thermal or ultraviolet radiation,
polyacrylonitrile (PAN) undergoes a chemical reaction
leading to discoloration and the appearance of an Electron
Paramagnetic Resonance (EPR) signal. Since that time,
research has been carried out on this process for three
general reasons: (1) textile chemists are concerned with
eliminating this phenomenon, which causes premature decay of
PAN containing fabrics; (2) catalyst chemists have been
interested in exploiting the catalytic properties inherent
in pyrolyzed polyacrylonitrile; and (3) materials scientists
have been interested in this process since it is the first
step in the conversion of acrylonitrile to carbon fibers,
which are an increasingly used component of composite
structural materials. Pyrolyzed polyacrylonitrile films
have even been proposed as a cost-effective amorphous
semiconducting material to replace silicon crystals in solar
4
collectors.
The earliest studies on the pyrolysis of
polacrylonitrile were infrared studies in which it was
observed that the carbon-nitrogen triple bond stretch
disappeared, the bands in the spectrum broadened
4

5
considerably, and a large broad absorption grew in the
region where C=N, C=0 and C=C would appear. However, the
spectrum was too featureless to permit specific assignments.
On the basis of this evidence, it was proposed that the
reactions occurring during pyrolysis and polymerization were
as illustrated in Figure l.5'6'7 a recent spectroscopic
study on the vacuum pyrolysis of PAN8 concluded that an
intermediate phase existed at pyrolysis temperatures between
200 and 260°C. This material was said to be an intrinsic
semiconductor with an extensively delocalized n electron
system as depicted in Figure lb. They also suggested that
the reaction which occurs in this temperature range could
involve bonding and conjugation between adjacent chains
without interruption of the carbon backbone, due to the
atactic nature of the PAN starting material. This reaction
has also been proposed to occur during the alkaline
degradation of PAN.9 It has also been proposed that
partially oxidized species such as the imine-nitrone
copolymer illustrated in Figure 2 could be present and
contribute to the optical properties of PPAN (not pyrolyzed
under vacuum), based on studies done on synthetic model
compounds which have similar absorptions in the UV-visible
region.
PPAN exhibits a strong, fairly narrow EPR signal at
about g=2, and this has been used as a probe of the
pyrolysis reaction. The g value increases with increasing

6
nC H2=CÍ H
i
C=N
POLYMERIZATION
a.
CH CH CH
\\ \\ %
N N N N
1
PYROLYSIS
CHo CHp CH2 c H2
CH CH CH
C C C C
7 \ / \ / \ / %
N N N N
I
PYROLYSIS
CH CH CH C^H
c.
\
/Cx\ /C. C.\
N N N N
Figure 1: Proposed Pyrolysis Reaction

7
O
t
Figure 2: Imine-Nitrone Structure

8
thermal treatment time, becoming essentially constant after
about five hours.12 It was learned that the concentration
of unpaired spins per gram (as determined by an EPR
technique) was strongly dependent upon the pyrolysis
conditions as well as the temperature of measurements. In
addition, the number of unpaired spins per gram was
decreased by the presence of air during the measurement.
The important parameters of the pyrolysis seemed to be the
rate and duration of heating in addition to the atmosphere
in which the pyrolysis was being carried out. For example,
a sample pyrolyzed in an ammonia atmosphere contained more
unpaired spins per gram than an identical sample pyrolyzed
in air or nitrogen.13
There have also been studies to determine what volatile
products are formed during the pyrolysis reaction and it was
found that ammonia, hydrogen cyanide, acetonitrile,
acrylonitrile monomer, propionitrile, methacrylonitrile,
isobutylacrylonitrile and vinyl acrylonitrile were formed
during the reactions.14 It was also found that the relative
amount of volatiles produced generally increased with
temperature in the range studied (300-800°C).
Unfortunately, the temperature-time profile of these
experiments consisted of very fast rise times (<1 sec) to
the desired temperature followed by maintenance at that
temperature for 10-20 seconds after which time the products
were analyzed by Gas Chromatography (GC). The temperature-

9
time profile refers to a plot of the temperature versus
time, which enables one to make better comparisons of the
thermal history of each sample. It has been known for some
time that there is a strong exothermic reaction which occurs
between 200 and 300°C which is capable of generating enough
heat in a bulk sample to cause ignition of the sample15
concomitant with drastic weight losses resulting in a
material which has little catalytic activity. (Several
times during our investigations at the University of Florida
these runaway reactions were observed due to the
catastrophic failure of a temperature controller.) Since
the active catalytic species is presumed to be the product
of the reaction in Figure 1, it is logical to choose the
pyrolysis conditions designed to minimize weight losses
since the cyclization-dehydrogenation reactions depicted in
Figure 1 would result in a theoretical weight loss of less
than 5%.
There have been many reports describing the catalytic
activity of pyrolyzed polyacrylonitrile. PPAN has been
shown to be capable of isomerizing alkenes and
dehydrogenating alcohols,16 decomposing formic acid and
nitric oxide,17'18 dehydrogenating ethylbenzene19 and
cumene,20 and oxidizing ethylene to ethylene oxide.21
Manassen and coworkers, in a series of publications,
described the dehydrogenation of alkenes and of alcohols and
the isomerization of alkenes using PPAN catalysts. It was

10
found that unlike conventional dehydrogenation catalysts, no
gaseous hydrogen was evolved. Instead, the hydrogen was
thought to be either physically adsorbed or chemically bound
to the surface of the catalyst, and while prolonged heating
at elevated temperatures in a nitrogen atmosphere did not
restore the catalytic activity, a short treatment in air at
temperatures as low as 140°C completely restored the
activity with the concomitant production of water. The
dehydrogenation of cyclohexene to benzene was also observed
to follow a different pathway using PPAN catalysts in that
no disproportionation occurred as with commercial
dehydrogenation catalysts. When cyclohexene was passed over
various types of charcoals or graphite at elevated
temperatures, significant amounts of cyclohexane and other
products were formed while PPAN catalysts produced benzene
exclusively. In addition, as with the alcohols tested, no
gaseous hydrogen was produced and prolonged heat treatment
under a nitrogen atmosphere failed to restore the catalytic
activity while a short treatment in air was found to
completely restore the catalytic activity. It was also
found that dehydrogenation reactions tended to deactivate
these catalysts quicker than isomerization reactions, again
suggesting that the hydrogenation of the catalyst surface is
responsible for the deactivation of the catalyst. These
authors proposed the reaction illustrated in Figure 3 to be
the mechanism for the catalytic activity of PPAN catalysts.

11
H
Figure 3: Proposed Catalytic Mechanism

12
It has been demonstrated by Braude and coworkers22 that
dihydroquinoline can by hydrogenated by acridine while the
reverse reaction does not occur, as shown in Figure 4.
On the basis of these observations it was proposed that
higher annelation of a system of condensed heterocyclic
aromatic rings will result in poor H-donating properties.
It is also known that annularly condensed aromatic molecules
become less and less stable with an increasing number of
rings while the hydrogenated compounds gain in stability,23
which may help to explain why a compound containing the
proposed structure (Figure 1) of PPAN catalysts could be
both a good hydrogen acceptor and a poor donor.
Furthermore, model compounds were synthesized which were
incapable of forming hydroaromatic structures while
retaining aromaticity and were found to be catalytically
inactive.24 In addition, the transformation of a
hydroaromatic structure to an aromatic structure by air
oxidation has been shown to occur with acridine25 and model
compounds of condensed pyridine rings.26
Manassen and coworkers used the reaction of
5-ethyl-5-methyl-l,3-cyclo-hexadiene over PPAN catalysts in
an attempt to elucidate the form of the hydrogen transfer.
The scheme in Figure 5 illustrates the possible products
based on hydride ion abstraction, hydrogen atom abstraction
and proton abstraction. A comparison of the product
distributions obtained by passing

13
Figure 4: Dehydrogenation of Acridine

14
(a)
Hydri de
ion abstraction
5-ethyl-5-methyl -
1.3-cyc1onexadiene
abs tract i o*
(b) Hydrogen atom
5-ethyl-5-methv1 -
1,3-cyc1ohexadi ene
+ J
ethyl radical
(c) Proton abstraction
5-ethy1-5-methy 1 -
1,3-cyclohexadiene
methyl
carbanion
Figure 5: Possible Dehydrogenation Mechanisms

15
5-ethyl-5-methyl-l,3-cyclohexadiene over various catalysts
led the authors to conclude that the behavior of PPAN
catalysts is somewhat intermediate between that of
commercial dehydrogenation catalysts, which abstract H-atoms
and acidic alumina, which is a proton abstractor. Pyrolyzed
polycyanoacetylene, the structure of which is illustrated in
Figure 6, was found to exhibit only H-atom abstraction,
leading to the proposal that the hydride transfer aspect of
PPAN catalysis originates from another structural element
than the proposed structure of the active catalyst, since
both PPAN and pyrolyzed polycyanoacetylene should
theoretically have the same structure. They found support
for this hypothesis by the observation that treating the
PPAN with dimethyl sulfoxide can greatly enhance the
percentage of ortho-ethyl-toluene formed.
It is of interest to note that different preparations of
PPAN catalysts can have very different catalytic activities.
For example, Manassen and coworkers used the following
general preparative method for synthesizing their catalysts.
Polyacrylonitrile was spread out in a thin layer in a draft
oven, slowly heated to 230°C and kept at this temperature
for 12 hours, during which time the color went from white to
brown via yellow. The brown powder was pelleted at 8000
psi, crushed, sieved and then calcined at either 350 or
450°C for 30 minutes, during which time the color went from
brown to black. It was found that catalysts calcined at

16
Figure 6: Proposed Structure and Pyrolysis Product for
Polycyanoacetylene

17
450°C underwent decomposition, presumably yielding large
crosslinked structures, which were apparently somewhat
acidic resulting in a high activity for double bond shifts
as well as ethyl shifts. (It was the material calcined at
450°C which was used for the 5-ethyl,5-methyl-l,3-
cyclohexadiene experiment.) PPAN catalysts which were
calcined at 350°C showed much less activity for double bond
or ethyl shifts, which illustrates the importance of the
pyrolysis conditions on the activity of these catalysts.
These results are in agreement with electrical measurements
which have shown that drastic changes occur above 350°C.26
Another theory to account for the catalytic properties
of organic pyropolymers correlated the number of unpaired
spins per gram (as as measured by an EPR technique) to the
catalytic activity for the decomposition of nitrous
oxides.28'29 This theory includes a scaling factor called
the intrinsic activity of the free spins, which presupposes
that all of the free spins are not active in catalysis. The
assumption is that structural rearrangements preceding
graphitization bring about exchange interactions between the
free spins rendering some of them catalytically inactive. A
correlation was developed between the width of the EPR line
and the relaxation time, T1 (as determined by the saturation
technique) and the intrinsic catalytic activity of the free
spins.30 This concept seems somewhat related to the
electronic theory of catalysis on semiconductors, described

18
by Vol1kenstein31 and Hauffe.32 In the electronic
mechanism, the semiconductor catalyst acts as an electron
reservoir by either donating electrons to or accepting
electrons from the substrate in guestion. This mechanism is
in contrast to that proposed for alumina, in which the
reaction is suggested to occur on Bronsted acid or proton
donating sites through the formation of an adsorbed
carbonium ion. Cutlip and Peters15 examined the kinetics of
dehydration of t-butyl alcohol over PPAN catalysts at 240 to
280°C and applied both single site and dual site models to
describe the kinetics. After a considerable amount of
mathematical manipulation, these authors concluded that it
was impossible, based on their data, to uneguivocally
determine the reaction mechanism since a statistical best
fit could not be obtained for any of the models chosen.15
Since the electronic theory of catalysis greatly increases
the possible number of rate limiting steps in the mechanism
(dissociation of adsorbed molecules, transfer of electrons
between adsorbed species and the catalyst), it was only
possible to gualitatively evaluate this mechanism, and it
was proposed that the rate-limiting step would probably be
related to the concentration of free electrons in the
conduction band of the catalyst. This could possibly be
related to what Gallard-Nechtschein and coworkers described
as the intrinsic catalytic activity of the free spins and
its relationship to the catalytic activity of these

19
catalysts,30 thus raising the possibility that there are
unpaired electrons in addition to conduction band electrons.
One problem with reviewing the literature pertaining to
PPAN catalysis is that one is consistently faced with the
prospect of comparing apples to oranges in the sense that
many of the studies reported preparing their catalysts by
different methods, making it impossible to be certain
whether or not the preparations consisted of the same
chemical (or semiconducting, for that matter) species. In
many cases, the catalysts were only characterized by their
method of preparation and their catalytic activity.
Manassen and coworkers24 (as discussed earlier) noted a
difference in catalytic activity, as well as in elemental
analyses (C, H, N), by calcining in nitrogen for 30 minutes
at 350°C instead of 450°C. Another pertinent observation is
the fact that the sum total of the carbon, hydrogen and
nitrogen analyses is usually between 75 and 90%, indicating
that from 10 to 25% of the final catalyst is composed of an
element other than carbon, hydrogen or nitrogen. The most
likely candidate is oxygen since many of the catalyst
preparation methods involve pyrolysis in air for some
period, and none have sought to rigorously exclude oxygen
from the preparations. This gives further evidence that
structures such as the imine-nitrone depicted in Figure 2
may constitute a considerable proportion of PPAN catalysts,

20
as well as possibly being responsible for the Bronsted type
acidity.
A recent spectroscopic study done on silver backed
polyacrylonitrile films also implicated structures such as
the conjugated imine in Figure 1 as products in the UV
degradation of PAN films using light in the 250-400 nm
region.33 (It has been known for some time that UV light
and strong bases bring about similar structural changes as
heat treatment, producing dark colored, intrinsically
paramagnetic solids.) In the UV degradation study, changes
in the silver backed PAN films were monitored using Fourier
transform infrared reflection absorbance (FTIR-RA) under
both oxidative and non-oxidative conditions. Their results
indicated that under oxidative conditions, oxygen-containing
species such as alcohols, carboxylic acids, hydroperoxides
or ethers could be present. Since UV light is capable of
generating ozone, ozonolysis products are also theoretically
possible, although the presence of ozone was not detected.
In addition, the appearance of an N-H stretch and a C=N
stretch, as well as the loss of the carbon nitrogen triple
bond stretch, suggested that cyclization had occurred. The
film became discolored and developed an EPR signal; however,
when the film was redissolved in dimethyl sulfoxide, the EPR
signal disappeared while the color remained. This indicated
that discoloration is a necessary but not sufficient
condition for the presence of paramagnetism. Although

21
analogies can be drawn between the UV and thermal
degradation of PAN, it must be emphasized that there are
distinct differences since thermal degradation generally
results in an insoluble material. For example, only one
hour of heating at 160°C results in a solid of which only
20% can be extracted with dimethylformamide, the solvent of
choice for PAN. This may be attributed to more extensive
crosslinking occurring in the thermal process than in the
photochemical reaction, suggesting that the reactions in the
thermal process are considerably more complex. An
additional complication is the fact that all
polyacrylonitrile is not identical since the polymerization
conditions and polymerization catalysts used will have a
strong effect on the composition of the final product due to
defects and inevitable chain ends which may or may not
consist of polymerization catalyst residues. In studies of
the production of carbon fibers it has been learned that
pre-oxidation (prolonged isothermal heating at low to
moderate temperatures (150-250°C)) minimizes the exothermic
reaction and yields a flameproof material which has a
concentration of approximately 10% oxygen by weight.34
Although most of the catalytic studies of PPAN-based
catalysts had chosen simple, industrially unimportant
reactions in an attempt to establish a relationship between
the catalytic activity and the chemical or electronic
structure of the material, there have been a few studies of

22
industrially important reactions. Degannes and Ruthven18
investigated the oxidative dehydrogenation of ethylbenzene
to styrene (at atmospheric pressure and from 180 to 280°C)
and found that the reaction was zero order in oxygen and
approximately second order in ethylbenzene. The zero order
dependence on oxygen suggests that the reaction rate is
controlled by the rate at which ethylbenzene can be adsorbed
and dehydrogenated. Although most of these experiments were
for kinetic purposes carried out at low conversions under
differential conditions, a limited series of integral
experiments showed that conversions greater than 80% at
325°C could be achieved with no significant by-products.
This is in contrast to commercially available
dehydrogenation catalysts which do not operate oxidatively,
resulting in an endothermic process requiring temperatures
in excess of 500°C to achieve conversions of 50% to styrene,
with benzene and toluene being significant side
products.35,36,37
There have been very few studies in the literature
dealing with two potentially important areas: doping PPAN
catalysts with metals and supporting PPAN catalysts on
inorganic oxide supports. In one study, acrylonitrile was
polymerized using 2,2'-azobis[2-methyl propionitrile (AIBN)
in the presence of silica gel and the resulting material was
pyrolyzed at several different temperatures.20 Metals were
then added to these materials by slurrying with ethanol

23
solutions of the corresponding metal chlorides followed by
filtration and washing with ethanol. X-ray photoemission
spectroscopy (XPS) data were cited to propose copper binding
to nitrogen atoms in the PPAN since it was noted that only
the Cu 2p 3/2 and the N Is binding energies shifted upon
addition of CuCl to the silica supported PPAN samples. The
reactions studied were the oxidations of cumene and
ethylbenzene at 100°C and under 1 atm. of oxygen. For
cumene, the total conversion was 63% with the selectivity
being 63% cumyl alcohol and 28% acetophenone; for
ethylbenzene, the total conversion was 21% with a
selectivity of 87% acetophenone and 13% 1-phenylethanol.20
These results for ethylbenzene are quite different from the
vapor phase results obtained by Degannes and Ruthven19 who
observed styrene as the only product. This apparent
contradiction is not too surprising when one considers the
differences in the catalyst preparation and catalytic
reaction conditions. The silica-supported PPAN catalyst
contained of course Si02 and was also different in that it
was pyrolyzed for 12 hrs. at 190°C. The previous work by
Degannes and Ruthven had employed a catalyst which had
slowly been heated to 230°C in air, was maintained at that
temperature for 16 hrs., and then was calcined at 400°C in
an atmosphere of nitrogen for 4 hrs. In addition, one study
passed gaseous ethyl- benzene over the catalyst from 180 to

24
325°C, while the other reaction occurred in the liquid phase
(or at the liquid-solid interface) at 100°C.
The data available in the open literature concerning the
catalytic activity of PPAN are sketchy at best, which is
rather surprising since these few studies suggest that PPAN-
based materials could be promising catalysts for
dehydrogenation as well as oxidation and dehydration.
Perhaps one reason why there are so few reports of PPAN
catalysts is because of the formidable problems associated
with studying a material without being able to employ
common solution techniques such as nuclear magnetic
resonance spectroscopy, UV-vis spectroscopy and infrared
spectroscopy.
It was the goal of this investigation to increase the
growing body of knowledge concerning catalysis using PPAN-
based materials in three general areas: the effects of the
pyrolysis conditions on the catalytic activity of PPAN
catalysts; the effects of metal additives on the catalytic
activity of PPAN catalysts; and the effects of supporting
PPAN catalysts on oxide supports such as silica (Si02),
titania (Ti02) and alumina (A1203). As in many of the
earlier studies, a simple model reaction was chosen for ease
in handling the analysis. In this case, ethanol was chosen
as the substrate and was shown to be capable of undergoing
dehydrogenations as well as dehydrations.

25
The effects of variations in the pyrolysis conditions
were investigated by several methods. Differential scanning
calorimetry (DSC) and thermogravimetric analysis (TGA) were
carried out on PAN samples both isothermally and in the
temperature programmed mode in an attempt to learn more
about the destructive exotherm and concomitant weight loss
characteristic of the pyrolysis reaction. In addition,
several different atmospheres were used in the pyrolysis
reaction in an attempt to learn whether the pyrolysis
atmosphere affects either the composition or catalytic
activity of these materials. Finally, the effect of
variations in the pyrolysis reaction heating rate on the
composition and catalytic activity of the resulting
materials was investigated.
In the hopes of obtaining useful information about the
effects of metal additives on the catalytic activity of PPAN
catalysts, three general approaches were taken: the
intrinsic basicity of PPAN was studied by titrating with
dilute acid, paramagnetic metal species were deposited on
the surface of PPAN in an attempt to learn more about the
metal environment using electron paramagnetic resonance
spectroscopy (EPR), and the effects of various metal
additives on the catalytic activity of PPAN were examined in
the dehydrogenation of ethanol.
The investigation of the effects of supporting PPAN on
oxide supports was carried out by preparing samples of PPAN

26
supported on silica gel and alumina and then comparing the
catalytic activities of these materials in the
fore-mentioned ethanol reaction with unsupported PPAN
catalysts as well as pure silica gel and alumina. The
surface areas of these materials were also measured in order
to make activity comparisons on a surface area basis. As
stated previously, the goal of this research was to expand
the general knowledge about PPAN based catalytic materials
in the hope that a greater understanding of the
structure-reactivity relationships could be obtained so that
in the future it may be possible for chemists to engineer
low-cost organic catalysts which are tailored to optimize a
specific reaction. Although the work of one graduate
student is not sufficient to achieve these lofty goals, it
is possible to address some apparent inconsistencies in the
literature regarding the selectivity of PPAN catalysts in
the reaction of ethylbenzene and the significance of the
support interactions, if any. The report of silica
supported PPAN complexes by Bai and co-workers20 was very
brief and many important experiments were either not carried
out or not reported. For example, the activities of the
metal-silica-PPAN catalysts were ranked as follows: Cu I >
Cu II > Co II > Mn II. Unfortunately, no comparison was
made between doped and undoped catalyst preparations. In
addition, no blank runs using only silica were attempted,
making it impossible to determine what role (if any) the

27
silica support plays in the catalytic reaction. Although
ethanol was used as the substrate in the present study for
ease in handling, it was felt that the results obtained
would be applicable to the ethylbenzene reaction since both
involved primarily dehydrogenation reactions.
In addition, a limited number of experiments were
carried out to determine the feasibility of using PPAN
catalysts to decompose dimethyl methylphosphonate (DMMP) and
chloroethyl-ethylsulfide (CEES), two compounds of military
interest as simulants for chemical warfare agents.
Preliminary experiments were also carried out to determine
whether PPAN catalysts are photochemically active or active
towards syngas conversion.

EXPERIMENTAL
Materials
N,N-Dimethylformamide was reagent grade, purchased from
Aldrich and used without further purification.
Acrylonitrile was reagent grade, purchased from Aldrich and
used without further purification. AIBN (2,2'-azobis[2-
methyl propionitrile]) was purchased from Eastman Chemicals
and used without further purification. Polyacrylonitrile
was reagent grade, purchased from Aldrich and used without
further purification. Alumina (neutral) was purchased from
Fischer and used as supplied. Titanium dioxide (anatase,
Ti02) was supplied by Baeyer and used without further
purification. Metal complexes were all reagent grade and
used as supplied. Zirconium basic carbonate and zirconium
basic acetate were generously donated by Mr. Brady Crom and
Dr. Tom Wilson, of Zirtech Inc., Gainesville, Florida, and
were used without further purification. Vacuum distilled
chloroethyl ethylsulfide (CEES) was generously provided by
Dr. Yu-Chu Yang of the U. S. Army. Silica gel (Si02), Grade
62 with a mesh size of 60-200, was provided by Davison and
used without further purification. Ethylbenzene was reagent
28

29
grade, purchased from Aldrich and distilled before use to
remove traces of toluene and benzene. Dimethyl
methylphosphonate (DMMP) was reagent grade, supplied by the
U.S. Army and used without further purification. Silver (I)
trifluoromethanesulfonate (AgCF3S03) was purchased from
Aldrich and used without further purification. Bis(2,2'-
bipyridine) ruthenium (II) chloride (Ru(bipy)2cl2) was
generously provided by Dr. E. Stine and used without further
purification. Glass spheres (8-58 /¿m, Standard Reference
Material 1003a) were obtained from the National Bureau of
Standards and used without further purification.
Preparations
Polymerization of Acrylonitrile
Acrylonitrile (100 g were added to 1500 ml of distilled
water in a flask fitted with a reflux condenser and
maintained at 60°C in a silicone oil bath under an inert
atmosphere. AIBN (1.00 g) was added to this mixture with
vigorous stirring. The mixture was allowed to stir for 1.0
hr after which another gram of AIBN was added. The
suspension was allowed to stir overnight and then the white
polyacrylonitrile powder was recovered by filtration and
washed with copious amounts of acetone. The resulting
polymer was dried at 50°C in a vacuum oven. It was found

30
that vigorous stirring using an overhead stirrer was
required during the reaction to prevent agglomeration of the
polymer into large lumps. In addition, to prevent loss of
the monomer the cooling water was passed through a saltwater
ice bath, using a copper heat exchange coil, before entering
the condenser. Due to its greater density, argon was found
to be superior to nitrogen in preventing the loss of the
highly volatile monomer.
Diphenvlglvoximato Cobalt (II) (Co(DPGH)2)
This synthesis was carried out under argon in
schlenckware apparatus using the method of Tovrog.38
Methanol was dried by distillation from calcium hydride and
stored over 4 A molecular sieves. Dry methanol (200 ml) dry
methanol was added to a 1 L three-neck round bottom flask
and the oxygen was removed by bubbling argon through the
system. Cobalt (II) acetate (12.45 grams) and
diphenylglyoxime (24.00 grams) were added to the reaction
flask and the mixture was allowed to stir for several hours
until the cobalt (II) acetate dissolved. After several more
hours of stirring, the brown suspension was filtered through
a schlenckware frit yielding a brown solid.
Analyzed: Carbon — 59.23%, hydrogen — 4.11%, nitrogen —
9.85%
Calculated for Co(DPGH)2'2H20: Carbon — 58.99%, hydrogen
— 4.57%, nitrogen — 9.77%

31
PAN on Silica
Two preparations of silica gel supported PAN were
formulated with loadings of approximately 7 and 17% PAN by
weight. The appropriate amount (2.00 grams or 5.00 grams)
of PAN was dissolved in 250 ml of N,N-dimethylformamide.
Silica gel (30 g) was added to the solution and the
suspension was allowed to stir on low heat (90°C) for 18
hrs. The suspension was then rotary evaporated to dryness
and subsequently pyrolyzed. While the 7% PAN on silica gel
became an off-white free flowing solid after rotary
evaporation, the 17% PAN on silica gel became caked up and
required grinding in a mortar and pestle before pyrolysis.
PAN on alumina
This sample was prepared as a 7% loading of PAN on
alumina. A solution of 3.66 grams of PAN and 250 ml
dimethylformamide was added to 50.00 grams of alumina and
the suspension was allowed to stir overnight at 90°C. The
dimethylformamide was removed by rotary evaporation and as
in the case of the 17% PAN on silica gel, the resulting
solid was caked up and very hard. The material was
pyrolyzed after grinding in a mortar and pestle.

32
PAN on Titanium Dioxide
The same general procedure was followed to prepare a
solid which was approximately 15% PAN. PAN (3.00 g) was
dissolved in 300 ml of N,N-dimethylformamide to which was
added 20 g of anatase (Ti02) and the resulting suspension
was stirred for 18 hrs. at 90°C. The dimethylformamide was
removed by rotary evaporation and the caked up material was
ground in a mortar and pestle and subsequently pyrolyzed.
Metal Incorporation for Catalyst Studies
For the purposes of comparing the effects of various
metal additives on the catalytic activity of PPAN catalysts,
a different method of metal incorporation was used in an
attempt to have equimolar amounts of metal species per gram
of PPAN. Two series of catalysts were prepared using 3.45 x
10-4 moles of metal complex per gram of PPAN. Within each
series of catalysts, the PPAN starting material was all from
the same batch in order to eliminate differences due to
pyrolysis conditions. A weighed amount of metal complex was
dissolved in absolute ethanol and stirred with the
appropriate weight of PPAN to obtain a concentration of 3.45
x 10-4 moles metal/gram of PPAN. This suspension was
stirred overnight without heat, rotary evaporated and then
dried in a vacuum oven at room temperature.

33
Copper-Lithium Catalyst Preparation
PAN (20.00 g) was dissolved in 200 ml of
N,-N-dimethylformamide (DMF) with stirring and then 1.43 g
of LiCl was added. The solution was allowed to stir on low
heat for 18 hrs after which time the DMF was removed by
rotary evaporation. (Before rotary evaporation several
films were prepared by filling a small petrie dish about one
half full of the solution and drying in a vacuum
desiccator.) The resulting hard amber colored chunk of
plastic was ground up in a Waring Blender and then mixed
with 0.40 g of copper powder. The mixture was allowed to
sit with occasional stirring for about two weeks. During
this time, the color of the polymer slowly turned from amber
to green and the copper metal ceased to be observable. This
material was subsequently pyrolyzed as will be described.
Another batch of catalyst was prepared in which the copper
metal was added directly to the PAN-DMF-LiCl solution and
allowed to stir for about two weeks until all of the copper
was dissolved. The dark brown solution was rotary
evaporated to yield a dark brown plastic which was
subsequently ground up and pyrolyzed. A lithium chloride
catalyst was prepared in the same manner as the copper-
lithium catalyst except that copper was not added in this
case.

34
Ruthenium Catalyst
A solution of Ru3(CO)was prepared by dissolving 0.20
g of Ru3(CO)22 in 200 ml of toluene in a 500 ml erlenmeyer
flask. Powdered PAN (20.00 g, Aldrich) was added to this
solution which was allowed to stir on low heat for two days.
During this time the suspension became green looking. The
material was filtered using a glass frit resulting in a
light green polymer powder. The filtrate was bright orange
and virtually indistinguishable from the original solution.
Upon drying, the polymer powder became somewhat off white in
color. The FT-IR diffuse reflectance spectrum of this
powder indicated that it contained Ru3(CO)12; however the
color of the filtrate indicated that most of the Ru3(CO)12
remained in solution. This material was subseguently
pyrolyzed as will be described.
Methods
Pyrolysis Reaction
The pyrolysis reaction was carried out using a variety
of methods for two general reasons. First, as more was
learned about the pyrolysis reaction, modifications were
made to the apparatus in order to obtain more homogeneous
products. Second, the means of temperature control was
upgraded as funding allowed. The pyrolysis tubes were
constructed of 1.00 in. diameter pyrex tubing with a glass

35
frit at one end and a 24/40 ground glass joint at the other.
The ends were then tapered to about 1/4 inch diameter to
accommodate stopcocks and fittings for tygon hose
connections. Since highly poisonous gases were produced
during these reactions, all operations were carried out in a
fume hood. Using tygon tubing, the mineral oil bubblers
were attached to both ends of the pyrolysis tube to monitor
the carrier gas flow as well as for leak detection. A
schematic diagram of the apparatus is shown in Figure 7.
Initially a Fisher mechanical temperature controller was
used to monitor and control the temperature of the system
which was contained in a commercially available tube furnace
manufactured by Lindberg. A mercury-glass thermometer was
also used as a back-up for monitoring the temperature. It
was subsequently discovered that this temperature controller
had a fluctuation of approximately + 50°C at any given
set-point. This proved to be unacceptable since it had been
demonstrated in other studies (and born out in this
investigation) that accurate temperature control in the
200-300°C region of the reaction is crucial in controlling
the strong exothermic reaction which occurs in this
temperature region. Failure to control this exothermic
reaction results in a runaway reaction causing the PAN to
char very quickly with a dramatic loss in weight. To
circumvent this problem, an Omega CN-300 digital temperature
controller was purchased. Unfortunately, electrical

36
Figure 7:
Initial Pyrolysis Apparatus

37
problems were encountered in that the relay responsible for
turning the oven off and on had a propensity to stick in the
on position causing the oven to heat to 800-900°C, resulting
in damage to the tube furnace and the pyrolysis tube. In
theory, this problem should not have occurred since the tube
furnace was rated to draw only 6 amperes while the
temperature controller was rated for a 10 ampere load. This
problem (which was in all probability caused by power surges
in the line) was alleviated by including a second relay, or
slave relay, rated for 25 amperes in the circuit. In this
configuration, the relay in the temperature controller
served only to switch the slave relay, which directly
controlled the oven. This served to reduce the current
passing through the temperature controller's relay, thus
increasing its lifetime. During the course of these
investigations, several other modifications were made to the
system. Initially, the pyrolysis tube and tube furnace were
positioned horizontally; however this proved to be a problem
since the pyrolysis reaction produced noxious liquids which
tended to settle towards the bottom of the tube and
impregnate the product. It was found that by placing the
apparatus in a vertical configuration, these liquids were
able to drain out of the pyrolysis tube into the tygon
tubing. Figures 7 and 8 schematically illustrate two
different versions of the pyrolysis apparatus positioned
vertically. Glass wool was packed into the ends of the tube

38
TEMPERATURE
Figure 8:
Modified Pyrolysis Apparatus

39
furnace to prevent heat loss. In addition, it was found
that positioning the control thermocouple outside the
pyrolysis tube resulted in a systematically higher
temperature reading. Therefore the thermocouple was
inserted directly into the PAN near the center of the tube
furnace as shown in Figure 8.
One problem with the Omega CN-300 temperature controller
was that it was manually controlled, which made the exact
duplication of the temperature-time profile rather
difficult. In addition, the temperature was raised
incrementally making the temperature-time profile a step
function. Since one of the goals of this investigation was
to compare catalysts produced using different pyrolysis
conditions, it was imperative that the temperature-time
profiles be highly reproducible. Towards the end of this
study, an Omega CN-2000 programmable temperature controller
was purchased which enabled exact duplication of a given
temperature-time profile. In addition, this controller was
capable of increasing the temperature continuously, thus
eliminating the need to raise the temperature incrementally.
During the course of the pyrolysis experiments, several
aspects of the pyrolysis conditions were varied to determine
the effect on the catalytic activity as well as the
elemental composition. All elemental analyses were
performed by the University of Florida Micro-analytical
Services operated by Melvyn Courtney. One line of

40
investigation involved using the same temperature-time
profile and using ammonia, carbon monoxide, air or nitrogen
as the carrier gas. In addition, variations in both the
rate of increase in temperature and the duration of heating
were examined with respect to the effect that these factors
have on the catalytic activity and elemental composition.
Catalytic Studies
As in the case of the pyrolysis apparatus, the
catalytic evaluation system was modified and upgraded
throughout the course of these investigations. What was
needed was a system capable of operating at atmospheric
pressure in the temperature range from 100 to 300°C. In
addition, a high degree of reproducibility in the reaction
conditions was desirable to enable accurate comparison of
the catalytic activities of different catalyst formulations.
The initial screening studies employed a catalyst evaluation
system as depicted in Figure 9. A carrier gas (usually
nitrogen or air, both unpurified) flowed through a bubbler
containing neat substrate and the resulting gas stream was
assumed to be saturated in substrate. The reactor tubes
were approximately 1 cm I.D. with a sintered glass frit in
the center to support the catalyst. The heating system
consisted of a homemade tube furnace controlled by a
variable AC power supply and monitored with a mercury
thermometer located near the center of the tube furnace

41
Figure 9:
Initial Catalytic Evaluation System

42
outside the reactor tube. The tube furnaces were
constructed as in Figure 10 using pyrex or vycor tubing and
heavy gauge nichrome wire. The following substrates were
used in this system to determine if PPAN catalysts had any
activity towards them: CEES, DMMP, ethanol, methanol, CO/H2
(syngas), ethylbenzene, norbornene and 2-propanol. In the
case of syngas, the gases were introduced with a system of
two bubblers which enabled one to vary the ratio of carbon
monoxide and hydrogen. The system contained enough volume
for adequate mixing and the resulting mixture was passed
through an additional bubbler (corresponding to the
substrate bubbler in Figure 9) which was filled with mineral
oil and used to monitor the gas flow. The mineral oil
bubbler at the end of the system was used to monitor the
effluent flow as well as to aid in the detection of leaks in
the system. In the reactions involving CEES and DMMP,
Clorox bleach was used in the last bubbler to hydrolyze the
potentially harmful starting materials and reaction
products.
Although a system such as that illustrated in Figure 9
was useful for determining whether a particular catalyst had
any catalytic activity towards a certain substrate, there
were several inherent problems in the system which prevented
accurate catalyst activity comparisons. The two main
problems with this system were the inability to accurately
control and monitor the temperature and rate of influent

43
GLASS
Figure 10
Oven Design

44
supply. In addition, the amount of substrate that was
supplied by the bubbler was so small that it generally took
more than a week of continuous running to pass 1 ml of
substrate through the system. Figure 11 is a schematic
drawing of a catalytic evaluation system which was designed
to eliminate these difficulties. To overcome the
temperature control problem, a thermocouple well was built
into the reactor tube allowing accurate measurement of the
temperature near the catalyst bed. The K thermocouple was
connected to a digital, time-proportioning control, solid
state temperature controller (Omega CN 300) which enabled
the temperature to remain at + 2° C of the set point. It
was found that the temperature as measured by the
thermocouple in the thermocouple well was consistently
5-10°C cooler than the temperature measured by the
thermometer outside the reactor tube. In addition, a
mechanical syringe was designed and constructed as
illustrated in Figure 12 in order to both stabilize and
increase the feed rate. Generally a 5 ml Hamilton Gas-tight
syringe was employed and the mechanical syringe, operating
with a 10 RPM motor, delivered 4.6 ml in approximately 7
hrs, making the Weight Hourly Space Velocity (WHSV) equal to
approximately 0.6 hour--'-. The WHSV denotes the ratio of the
mass flow rate of feed to the mass of the catalyst used as
defined by the following equation: WHSV= pV/W where p is
the density of the feed, W is the weight of the catalyst and

45
GAS IN :
MEG HANICAL
SYRINGE
TUBE FURNACE
THERMOCOUPLE CO
/
EXTERNAL -1
RELAY
TEMPERATURE
CONTROLLER
DRY ICE/acetone
BATH
MINERALOIL
BUBBLER
SEPTUM
TYGON TUBE
GLASS BEADS
CATALYST
FRI T
^*GAS OUT
MINERALOI L
BUBBLER
LIQUID
TRAP
Figure 11: Improved Catalytic Evaluation System

46
Figure 12: Syringe Design

47
V is the characteristic volumetric flow rate of the fluid.
The space time is defined as the reciprocal of the space
velocity. This corresponds to a very fast flow rate, almost
as high as used in many effecient, industrial processes.
Therefore the conversions obtained using this system (Figure
11) tended to be much lower than in the earlier catalytic
set-up illustrated in Figure 9.
Basicity Studies
Two types of basicity studies were attempted in this
investigation: the direct titration with dilute HC1 of
slurries of PPAN in distilled water, and the determination
of metal complex uptake from slurries of PPAN in agueous
solutions of several metal complexes using UV-vis or EPR
spectroscopy. In the direct titration method, a weighed
sample of PPAN was slurried in a 100 ml erlenmeyer flask
containing about 50 ml of distilled water (Millipore
Nanopure Water System) and titrated with a 0.01 molar
solution of HCl, using a Fisher pH meter to monitor the pH.
Since the titrations were often lengthy, the buret tip and
pH electrode were sealed to the flask using parafilm to
prevent evaporation. In the metal binding studies,
solutions of metal complexes were slurried with PPAN
preparations for about 24 hours and then washed and dried
under nitrogen. The dried PPAN was subseguently examined
in the EPR for evidence of metal binding. In another set of

48
experiments, a solution with a known absorbance was slurried
with PPAN and then filtered and washed. The filtrate and
washings were combined and reduced by blowing nitrogen to
the original volume of solution. The absorbance of this
solution was then checked by UV-vis and compared to the
absorbance of the original solution.
Photolysis Reactions
A Hanovia medium pressure mercury lamp was used as the
light source for photolysis reactions. Suspensions of PPAN
in various substrates were placed in quartz or pyrex
reaction vessels and stirred using a magnetic stirrer.
Photolysis products were analyzed by gas chromatography.
Thermal Analysis
Differential Scanning Calorimetry (DSC) and
Thermogravimetric Analysis (TGA) were performed by Ann
Livesey of the U.S. Army on a Dupont 9900 DSC and a Kahn
TGA. These analyses were carried out using commercially
available PAN supplied by Aldrich Chemical Company. In
addition, TGA and DSC results were obtained at the
University of Florida using a Perkin Elmer Series 7 Thermal
Analysis System.

49
Surface Area Measurements
A Micromeritics Digisorb 2600 was used to measure the
surface areas of many of the catalysts. In addition,
surface areas were also obtained courtesy of Dr. Willie
Hendrickson, 3M Corporation. The measurements were based on
the BET method. On analyses performed on the Digisorb 2600,
the samples were degassed for 12 hours at 90°C, after which
time the samples were weighed. By weighing after degassing,
the contribution due to adsorbed water was minimized. The
sample weights were entered into the computer and the
analysis, calculations and report printout were performed
automatically by the instrument.
Gas Chromatography-Mass Spectrometry Analysis
GC-MS Analysis of the catalytic reaction products from
several substrates was performed by Dr. Dennis Rohrbaugh of
the U.S. Army.

RESULTS AND DISCUSSION
Pyrolysis Studies
One area of investigation in this study was the effects
that different pyrolysis conditions had upon the chemical
and catalytic properties of PPAN. Variations were made in
the temperature-time profile of the pyrolysis reaction as
well as the atmosphere under which the pyrolysis was carried
out in order to determine what effects these factors had
upon the elemental analysis and catalytic activity of these
preparations. Table 1 contains the elemental analysis
results obtained by varying the atmosphere under which the
pyrolysis was carried out. The last two entries differ from
the other samples in that they were pyrolyzed using a
programmable temperature controller. In addition, the
temperature programs for these two samples were identical;
therefore, the only difference between these two
preparations was the atmosphere under which the pyrolysis
was carried out. Figure 13 gives some representative
temperature-time profiles for some of the preparations
listed in Table 1. It should be noted that with the
exception of sample number 5, all of the manually controlled
pyrolysis reactions employed approximately the same
50

51
time (days)
Figure 13: Representitive Temperature-Time Profiles

52
Table 1: Elemental Analysis Results
Sample C% H% N% C/N Total Pyrolysis Maximum
ID Atmosphere Temp.(C)
Number
1*
52.41
2.63
24.44
2.14
79.48
air
300
2*
69.25
3.00
20.95
3.31
93.20
nitrogen
450
3*
59.47
2.40
23.21
2.56
85.08
air-250
N2-400
400
4*
74.11
0.87
17.19
4.31
92.17
nitrogen
800
5*
68.73
3.73
21.40
3.21
93.86
nitrogen
400(fast)
6*
70.25
4.00
21.82
3.23
96.07
ammonia
400
7*
69.99
3.15
22.38
3.13
95.47
ammonia
400
8*
69.76
2.70
21.05
3.31
93.51
CO
400
g**
69.39
3.78
21.59
3.21
94.76
nitrogen
400
10**
65.02
3.19
22.64
2.87
90.85
air
400
*manual temperature controller
**programmable temperature controller

53
temperature-time profile within the limits of experimental
error. Although at first glance one sees no apparent
relationships between the different preparations, several
general conclusions can be drawn. Reactions carried out in
an air atmosphere tend to have lower total percentages of
carbon, hydrogen and nitrogen than reactions carried out
under atmospheres of nitrogen, ammonia or carbon monoxide.
Presumably, this is due to increased oxygen incorporation
into the products of air pyrolyzed samples. In addition,
air pyrolyzed samples tend to have a lower percentage of
carbon than samples pyrolyzed in nitrogen, carbon monoxide
or ammonia. This effect appears more dramatic when one
looks at the C:N ratios of less than 3:1 for air pyrolyzed
samples while samples pyrolyzed in other atmospheres have
C:N ratios of greater than 3:1. Another general observation
is that all other factors being equal (temperature-time
profile, atmosphere), higher pyrolysis temperatures seem to
result in lower percentages of hydrogen in the final
product. These results are in general agreement with
reported results.3
Thermooravimetric Analysis Results
The TGA results also indicated that the pyrolysis
conditions are very important in determining the composition
of the final PPAN product. Figure 14 shows the effects of
isothermally heating at 280, 290 and 300°C on the TGA

WEIGHT (%)
54
Samóle: POLYACRYLONITRILE
Size: 10.32 mg
Method: RAMP/ISOTR 2B0 FOR 60
Comment: PT PAN / N2 100ML/MIN
TGA
105
65 _
60 —r * i i i t : : r ' t 1 ' i
0 5 10 15 20 25 30 35 40 45 50 55 60 65 70
Time (min)
Figure 14: Isothermal Thermogravimetric Analysis

55
thermograms of PAN. The method used was to ramp the
temperature rapidly (20°C/minute) to the given temperature
and then monitor the weight loss as a function of time. It
can be seen that the onset times for weight loss were
shorter and that the total weight loss was greater for
higher isothermal temperatures. Figures 15, 16, and 17 are
the results of temperature programmed TGA scans for
temperature ramping at 2.5°C/minute, 5.0°C/minute and
10°C/minute, respectively. The typical sample size was
approximately 10.0 mg and the analyses were carried out in a
platinum boat under nitrogen flowing at 100 ml/minute.
These scans indicate that the weight remains essentially
constant until at least 280°C; however, Figure 18, which is
a composite of Figures 15, 16, and 17, exhibits somewhat
anomalous behavior. Specifically, the scan corresponding to
a program rate of 5°C/minute seems to exhibit rather
uncharacteristic behavior. For the purposes of catalyst
preparation, 500°C is probably the maximum temperature of
interest and at that temperature one would expect the 5°
C/minute scan to fall between the scans corresponding to 10°
C/minute and 2.5°C/minute. Instead, the scan at 5°C/min
exhibits a smaller weight loss than either of the other two
scans. It is most likely that either an instrument
malfunction or an experimental error was the cause of this
discrepancy. If one looks at Figure 16, it can be seen from
the derivative curve that this sample is gaining weight from

56
Samóle: POLYACRYLONITRILE
Size: 13.12 mg
Method: PAN 900 8 2.5
Comment: PT PAN / N2 100ML/MIN
TGA
Temgerature (*C)
Figure 15: Thermogravimetric Analysis, 2.5°C/minute

57
Samóle: POLYACRYLONITRILE
Size: 10.05 mg
Method: PAN 900 § 5
Comment: PT PAN - N2 100 ML/MIN
TGA
Temperature (*C)
Figure 16: Thermogravimetric Analysis, 5.0°C/minute

58
Samóle:
Size:
Method:
Comment:
no —
100-
qn -
— 30 -
c.
Cl
60-
50-
1
JO -)
J
30 -L
” 0
POLYACRYLONITRILE
10.02 mg
PAN 900 0 10
PT PAN / N2 100ML/MIN
TGA
- 2E
- 19-c
CJ1
- 14>
c.
IS
â–¡
- 10—
- 6
- 2
100 200 300 ' 400 ' 50(3
600
700 800
—r- -2
900
Temperature (*C)
General VI.OJ OuPont 9900
Figure 17: Thermogravimetric Analysis, 10.0°C/minute

59
Sample: POLYACRYLONITRILE
Size: 10.02 mg
Method: PAN 900 @ 10
Comment: PT PAN / N2 100ML/MIN
TGA
File: A: PAN. 13
Ooerator: ABL
Run date: 01/13/S5 13:42
Figure 18
Composite of Temperature Programmed
Thermogravimetric Analysis

60
about 500°C to about 700°C. Since this is not apparent in
Figures 15 and 17, it is doubtful that the increase in
weight is due to nitrogen incorporation from the atmosphere.
Most likely, the atmosphere in the TGA chamber was
contaminated with oxygen, which was responsible for the
increase in weight. Unfortunately, it was impossible to
repeat these experiments to determine the cause of the
error. The data obtained before 325°C appear to be
reliable, however, and several important conclusions can be
drawn about the pyrolysis reaction.
As earlier studies have indicated, the reaction begins
at higher temperatures when faster heating rates are used.
This has been attributed to the induction period for the
exothermic polymerization of the nitrile groups39>40'41
which can reach explosive speeds resulting in the
destruction of the polymer chain.42 Figure 19 contains the
TGA (obtained at the University of Florida) corresponding to
a temperature program rate of 0.5°C/minute and the results
qualitatively agree with the scans done at faster heating
rates. The onset temperature for weight loss is
considerably less than in the scans shown in Figure 18. In
addition, the percentage weight loss is also less than for
the other runs; however, since this run was carried out
using a different instrument, the absolute percentages
should not be taken too literally since the instruments may
not have been identically calibrated. Figure 20, from some

61
Samóle Weight* 0.742 mg
poiyacry loftltrl le
PERKIN-ELMER
7 Series Thermal Analysis System
Figure 19: Thermogravimetric Analysis, 0.5°C/minute

sample weight
62
Figure 20:
Thermogravimetric Analysis and
Differential Thermal Analysis

63
work by Grassie and McGuchan, shows Differential Thermal
Analysis (DTA) curves for polyacrylonitrile heated for
10°C/minute in air, nitrogen and vacuum.43 It can be seen
from Figure 20 that air pyrolysis results in greater weight
loss than nitrogen pyrolysis, while the greater weight loss
for the vacuum pyrolysis can be attributed to the ease of
volatilizing high-boiling fractions. In addition, it can be
seen that the onset temperature for weight loss is lowest
for vacuum and lower for air than for nitrogen. This would
seem to suggest that there are processes occurring in the
air atmosphere at lower temperatures than in the nitrogen
atmosphere. Most likely, the lower temperature reactions
are due to oxygen incorporation reactions not occurring
during the nitrogen pyrolysis. The Differential Thermal
Analysis (DTA) curve in Figure 20 is also interesting in
that it is tilted due to the fact that the sample is giving
off so much heat that the system becomes hotter than the
programmed rate of temperature rise (the DTA instrument
measures the difference in temperature between a reference
and a sample when both are heated under identical
conditions). It should be noted that these results have
severe implications for the pyrolysis of bulk samples of PAN
due to the exothermicity of the nitrile polymerization
reaction. The destructive reaction could occur at
relatively low temperatures if the heat of polymerization is
not dissipated rapidly enough to prevent the interior of the

64
polymer from reaching the critical temperature.42 Since the
catalyst preparation method used in this study employed a
pyrex tube containing about 15 grams of PAN, heat
dissipation was potentially a serious problem.
Differential Scanning Calorimetry Results
The DSC results also suggested that there is some sort
of an induction period associated with the pyrolysis
reaction. Figure 21 shows the effects of heating rate on
the exotherms. As in the temperature programmed TGA
analysis, the onset temperatures increase with the heating
rate. If there were no induction period, one would expect
the onset temperature to be lower with increasing heating
rate. A comparison between the onset temperatures for the
DSC and TGA results of temperature programming at 2.5, 5.0,
and 10°C per minute demonstrates that there is fairly good
correlation between the onset temperatures for weight loss
and the onset temperatures for the exotherm, implying that
the reaction resulting in weight loss is an exothermic
reaction. The shapes of the exotherms in Figure 21 are
quite similar but the curves differ in the onset
temperatures and the amount of heat given off. Figure 22
contains the DSC results for a heating rate of 0.5°C/minute,
and again the results qualitatively agree with those for
faster heating rates in that the onset temperature for heat
loss is less while the amount of heat loss seems to be

MUAL/SEC
65
187 207 227 2<7 267 287 307 317
TEMPERATURE (°C)
â–¡ SC
86/B1/08 TIME: 14:03
Figure 21:
Temperature Programmed Differential
Scanning Calorimetry

Heat Flow (mW)
66
PERKIN-ELMER
7 Series Thermal Analysis System
Temperature (C)
Figure 22
Differential Scanning Calorimetry, 0.5°C/minute

67
reduced with respect to the results obtained at faster
heating rates. Again, direct comparisons should not be
taken too literally due to the fact that two different
instruments were used to obtain these results. In general,
it seems that faster heating rates result in a greater
amount of heat being given off, correlating with the TGA
results which indicated that more weight was lost with
faster heating rates. It is also interesting to note that
there seems to be an endothermic process occurring at higher
temperatures than the exothermic process. In the exothermic
process, the amount of heat absorbed also seems to slowly
increase with increasing temperature, although this could be
an instrumental artifact. Figure 23 shows the effects of
isothermally heating at 240, 245, 250, 255, and 260°C. The
samples were heated at the rate of 1°C per minute until
reaching the desired temperature and then maintained at that
temperature for a period of time. The isothermal DSC
results qualitatively agree with the isothermal TGA results
in that lower isothermal temperatures result in longer onset
times. Also, the amount of heat given off increases with
increasing temperature, in agreement with TGA results which
showed that higher isothermal temperatures resulted in
greater weight loss. In addition, the shape of the exotherm
varied with isothermal temperature in that higher isothermal
temperatures resulted in larger, sharper exotherms. As in
the temperature programmed DSC runs, there seems to be an

68
247 257 267 277 287 297 307 317 327
TEMPERATURE CO
â–¡ SC
Figure 23: Isothermal Differential Scanning Calorimetry

69
endothermic process occurring after the exothermic process
and the amount of heat absorbed increases with increasing
isothermal temperature. Again it is not known at this time
whether this is an instrumental artifact. From the thermal
analysis results, several conclusions about the pyrolysis
reactions can be drawn. The reactions occurring during the
pyrolysis are dominated by a strongly exothermic reaction
concomitant with a dramatic loss in weight, preceded by an
induction period. The onset times for weight loss and heat
loss are strongly dependent upon the rate of heating and the
absolute temperature. In addition, there appear to be some
inconsistencies in treating this pyrolysis reaction as
merely a nitrile polymerization reaction since the dramatic
weight losses observed in this study suggest that a
considerable amount of chain destruction is taking place.
Earlier studies on a commercial acrylic fiber containing
methyl acrylate and acrylonitrile found a high correlation
between the nitrile content and the heat evolved during the
pyrolysis.44 The amount of unreacted nitrile groups was
determined by infrared analysis before and after pyrolysis
(Figure 24). The results of these analyses suggest that the
exothermic reaction is associated with the disappearance of
the nitrile groups, although it seems like an
oversimplification to attribute this reaction to a simple
nitrile polymerization reaction in the case of a
polyacrylonitrile homopolymer. In a copolymer containing

Nitrile Content, % (2240 cm
70
Í I
100 —
cw ^
SO —
" o
o o
SO
40 ,—
’ ^ o
/ Cs
?K
o.
ZC c
I I
40 50 80 100 *20 140 160
Residual Exotherm (cal/gram)
Figure 24: Plot of Residual Nitrile Content Versus
Area of Exotherm

71
relatively unreactive polymer subunits interspersed with the
nitrile containing subunits, it might be possible to view
the pyrolysis as a simple nitrile polymerization reaction.
If one calculates the theoretical weight loss associated
with the reaction in Figure 1, the proposed catalyst
formation reaction, the result is about 5% weight loss
principally due to hydrogen loss during cyclization and
aromatization. This does not account for the fate of
inevitable chain ends, defects and polymerization catalyst
residues which are present to various extents depending upon
individual sample preparation techniques. The weight losses
obtained for isothermal runs at 280, 290, and 300°C were
about 25, 30, and 35%, respectively (Figure 18), after 1
hour of heating at these temperatures. These percentages
seem far greater than what one would expect from the nitrile
polymerization reactions, even with the contributions from
the reactions undergone by the defects and impurities
previously mentioned. Furthermore, the nitrile
polymerization reaction should theoretically produce no
nitrogen containing volatiles; however, these were observed
in this as well as in all other studies. Although it is
conceivable that the nitrogen could have originated from
atmospheric sources, this is unlikely since the nitrile
moiety is far more reactive than the dinitrogen molecule.
In conclusion, while these and other results indicate that
the disappearance of the nitrile functional group is

72
intimately related to the exothermic reaction(s) occurring
during the pyrolysis, it is unlikely that a simple
intramolecular nitrile polymerization-aromatization is the
only reaction occurring to a significant extent. The
polymers used in this and other studies were atactic
polymers, making the likelihood of extensive intramolecular
cyclization more remote than in an isotactic polymer. In
reality one must view the nitrile groups as being randomly
oriented about the polymer backbone. This conformation
would increase the probability of reactions between adjacent
chains resulting in extensive crosslinking. The observation
that polyacrylonitrile becomes harder and more brittle after
heating supports this interpretation. Although the thermal
analysis studies, which employed about 10 mg of sample, may
not be directly applicable to the bulk pyrolysis of 15 to 20
grams of material, the results indicate the there may be a
serious design flaw in the pyrolysis apparatus used in these
studies in that the dissipation of the heat produced in the
exothermic reaction is crucial in preventing the destructive
runaway reaction. The apparatus employed about 20 grams of
PAN in a 1 inch diameter tube and had typical gas flow rates
of less than 100 ml/minute. The combination of rather
densely packed material and relatively slow gas flow rate
may have resulted in a heat dissipation problem. Although
the pyrolysis apparatus probably was not producing the best
catalysts possible, it proved to be adequate for the

73
purposes of this study. Some catalytic material was
prepared in our controlled temperature pyrolysis since
little or no activity was observed for materials heated too
quickly.
Surface Area Results
The results of the BET surface area determinations are
contained in Table 2. Since these results were determined
using different instruments and conditions (degas time,
degas temperature, etc.), these factors should be taken into
account when comparing the data. Several conclusions may be
drawn from these results. In general, the surface areas of
the pure PPAN materials are quite low, on the order of 0-9
square meters per gram. These results are considerably
lower than those reported by other workers who have obtained
values of 19.0-19.2 m2/gram,19 17.7 m2/gram15 and 18
m2/gram.16 These discrepancies can in part be explained
when one observes that the surface area of the PPAN'CuLi
catalyst increases after passing 15 ml of ethanol over two
grams of catalyst. In addition, it was observed that the
products collected from the catalytic reaction frequently
were amber colored, indicating that some of the waxes and
oils produced in the pyrolysis reaction had impregnated the

Table 2: Surface Area Results
Sample
Max
BET Surface
C
H
N
Comments
Temp(C)
Area (M2/g)
(%)
(%)
(%)
PPAN on
Silica3
400
260 (272)c
high load
PPAN on
Silica3
400
243 (263)c
2.55
0.34
0.66
low load
Silica3
—
(332)c
0.20
0.32
0.43
PAN
—
(286)c
67.12
5.82
26.17
PPAN3
400
5 (6)c
65.02
3.19
22.64
air pyr.
PPAN3
400
7 (9) c
69.39
3.78
21.59
nit. pyr
Alumina3
—
153c
0.03
0.24
0.00
PPAN on
alumina3
400
134 (13 4)c
3.55
0.52
0.73
nit. pyr.
PPANb
400
>lc
59.47
2.40
23.21
air pyr.
PPANb
450
4C
69.25
3.00
20.95
nit. pyr.
PPAN Cu-Lib
500
>lc
60.20
4.07
18.26
freshd
PPAN Cu-Lib
—
lc
—
—
—
spente
aThese samples were all pyrolyzed using an identical program
on an Omega CN 2000 temperature controller.
bThese samples were pyrolyzed using an Omega CN 300 manual
temperature controller.
cThese surface areas were generously provided by Dr. W.
Hendrickson of the 3M corporation while the other numbers were
obtained using a Micromeritics Digisorb 2600 surface area
analyzer.
dAfter pyrolysis.
eAfter passing 15 ml of ethanol over 2.00 g of this material
at 300°C.

75
catalyst, possibly filling up some of the pores and reducing
the surface area. The passage of ethanol through the
material may serve to wash out these residues and increase
the surface area. Granted, the surface areas of these two
samples are very small and possibly within the error limits
for this instrument. However, both samples were run on the
same instrument and although the absolute difference in the
"spent" and "fresh" catalyst samples is relatively small,
the surface areas differ by a factor of five. Degannes and
Ruthven19 reported that the surface areas of fresh and used
catalysts remained identical within the limits of
experimental error. The studies which reported surface
areas in the neighborhood of 20 m2/gram all used a similar
pyrolysis method in that they spread the PAN thinly on trays
and carried out the reaction in a draft oven. This
pyrolysis method is probably superior to the method used in
this investigation for two reasons. First, there is a
greater capacity for heat dissipation in a draft oven;
second, with the PAN spread in a thin layer is less of a
propensity for the pores on the surface to become clogged
with low molecular weight residues from the pyrolysis
reaction, thus lowering the surface area.
The surface area results for the silica and alumina
supported PPAN catalyst preparations were approximately 250
and 130 m2/gram, respectively, indicating that the surface
areas of these supports were not significantly altered by

76
the addition of up to 17% PAN by weight. These surface
areas were more than an order of magnitude greater than the
largest of the unsupported PPAN catalyst preparations, and
these results will be discussed later with respect to
catalytic activity comparisons between different
preparations.
Basicity Studies
Titration Results
Several attempts were made to determine the number of
basic sites on the surface of PPAN catalysts by directly
titrating with dilute HCl. The PPAN samples were either
used as is after the pyrolysis reaction or slurried with
concentrated sodium hydroxide, filtered and then washed with
distilled water. For a number of reasons it was impossible
to obtain quantitative data from this titration method. For
some unknown reason, the pH meter would not stabilize on a
reading in suspensions of PPAN materials. The meter
readings would continually randomly jump around when
immersed in PPAN suspensions but would stabilize immediately
upon being immersed in a solution not containing PPAN. In
addition, the meter would drift up as much as 4-5 pH units
over a period of several hours, making the time at which the
reading was taken very important. Throughout the titration,
the calibration of the pH meter was checked and found to be

77
accurate. Therefore, the instability in the readings was
caused by the PPAN itself, suggesting that some unknown
reaction was occurring at the electrode and/or on the
surface of PPAN. These results are in general agreement
with the only other titration study which demonstrated that
long-time exposure to dilute acid will result in some acid
uptake, indicating a generally basic nature.36 Although, as
stated previously, no quantitative results could be obtained
from this titration study, it is possible to make some
general observations about the basic properties of PPAN
materials. Although PPAN is essentially insoluble in all
acids, bases and solvents, it is by no means inert, as
evidenced by the instability of the pH electrode and the
slow uptake of acid. In addition, PPAN samples pyrolyzed in
air seem capable of absorbing more acid than samples
pyrolyzed in a nitrogen atmosphere.
Metal Binding Studies
As in the case of the titration studies, the results of
the metal binding studies were somewhat inconclusive in that
no direct evidence of metal binding was ever established.
The general method used was to add an EPR active metal
complex to PPAN by adsorption from solution and then look
for nitrogen hyperfine in the EPR. It was assumed that if
the polypyridine type structure in Figure 1 is the basic
site, the metal species would bind to the nitrogen and the

78
unpaired electron on the metal ion would interact with the
nuclear spin of the nitrogen (3/2), resulting in a nuclear
hyperfine interaction. Copper and cobalt were chosen as the
metal species since nitrogen hyperfine is known to be
readily observable for both at liquid nitrogen temperatures.
There have recently been several studies on the use of 1,8-
napthyridine based ligands in the formation of metal
complexes and these studies indicate that the napthyridine
moiety can accommodate several bonding modes, as shown in
Figure 25. Unlike bipyridines, the nitrogens in 1,8-
napthyridines are rather close together, resulting in a
significantly reduced "bite angle", thus making
napthyridines much poorer candidates for bidentate ligands.
In addition, its steric bulk can make the napthyridine
moiety a poor candidate for a monodentate ligand in certain
instances. There are numerous examples in the literature of
napthyridine based ligands forming complexes with many of
the metals in the periodic table, including the lanthanides
and the rare earth metals.45-50 The three general bonding
modes for these are depicted in Figure 25. These materials
are frequently highly colored due to the presence of a metal
to ligand charge transfer band, and have also been known to
exhibit fluxional behavior.51'52
i

79
Figure 25: Naptheridine Binding Modes

80
Diphenvlglvoximato cobalt (II) (Co(DPGH)2
This compound was prepared and stored in a nitrogen
atmosphere. The EPR samples were prepared in a glove box
using a 50:50 toluene:methylene chloride mixture which had
been exposed to three freeze-pump-thaw cycles to remove
dissolved oxygen. A saturated solution of Co(DPGH)2 was
slurried with a small amount of PPAN, the mixture was
filtered and EPR samples were prepared from the filtrate and
the solid material. The EPR spectra of solid samples of
Co(DPGH)2 and PPAN after being exposed to Co(DPGH)2 in
solution are given in Figure 26. The PPAN sample was run at
a higher sensitivity to detect any small changes in the
spectrum upon addition of the cobalt complex. The spectra
demonstrate that PPAN has little or no effect upon the EPR
spectra for Co(DPGH)2, other than the superimposition of the
narrow signal at about g=2 which is characteristic of PPAN.
Both the lineshape and g value of the cobalt complex remain
unchanged upon addition of PPAN. It should be noted that
solutions (50:50 toluene:methylene chloride mixture) of
Co(DMGH)2 and its pyridine adduct were prepared and the EPR
spectra were consistent with earlier published results which
showed the typical eight line cobalt spectrum with five line
nitrogen hyperfine in the case of the pyridine adduct
(Co(DPGH)2'2Pyridine).38 These results suggest that the
electronic structure of the cobalt is not perturbed upon
addition of Co(DPGH)2, implying that the Co(DPGH)2 present

81
3300 GAUSS
Electron Paramagnetic Resonance Spectra of
PPAN and Co(DMGH)2
Figure 26:

82
in the sample is merely physically adsorbed and not
chemically bound to the surface. Since pyridine binds to
this complex, the results show that PPAN is a poorer base
than pyridine, although at this time it is not possible to
discern whether the chemical structure, electronic structure
or steric bulk is responsible for PPAN's poor binding
ability. Since Co(DPGH)2 is a square planar complex with a
nearly planar ligand, the steric requirements for axial
adduct formation are not particularly stringent. This
suggests that the steric bulk of the PPAN is probably not
the over-riding cause for the poor binding ability of PPAN
materials.
Cupric chloride (CuCl2)
Cupric chloride (0.20 g) was added to a strongly basic
suspension of PPAN in distilled water and the resulting
mixture was allowed to stir for several hours, after which
time the mixture was filtered and the PPAN*CuC12 washed with
copious amounts of distilled water. The EPR spectra of this
material as well as that of solid CuCl2 are presented in
Figure 27. The EPR of PPAN*CuC12 consists of the usual PPAN
signal at about g=2, with another signal appearing at
slightly greater than g=2 as well as a much smaller but
discernable signal at about g=2.5. After washing with 0.1
molar HCl and copious amounts of distilled water the signal
remained unchanged. Since a strongly basic solution could

83
Figure 27: Electron Paramagnetic Resonance Spectra of
PPAN and Copper (II) Chloride

84
result in the formation of Cu(OH)2 on the surface of the
PPAN, it was hoped that washing with hydrochloric acid would
result in the protonation of any remaining basic sites and
the formation of chloride salts on the surface of the PPAN.
A comparison of the EPR spectra of solid cupric chloride and
PPAN*CuC12 reveals that the PPAN*CuC12 spectrum is not a
simple addition of the spectra for CuCl2(s) and PPAN as in
the case of Co(DMGH)2. Although the characteristic PPAN
signal appears unchanged, the copper signal in PPAN‘CuC12 is
markedly different from that of CuCl2(s), implying that
there is some sort of electronic interaction between the
PPAN and copper's unpaired electron, unlike the case of
Co(DMGH)2. Although the g values and lineshapes of the
copper signals are quite different, there is no evidence of
a nitrogen hyperfine interaction. Bai and coworkers20
prepared silica gel supported PPAN materials and studied
their physical and catalytic properties. Their work
presented XPS results for CuCl, PPAN on silica gel and PPAN
on silica gel with added CuCl, and they concluded that the
copper was coordinately bound to nitrogen based on
differences in the Cu 2p 3/2 and N Is binding energies upon
the adsorption of CuCl onto silica supported PPAN.20
Different preparative methods or the presence of silica gel
makes direct comparison of these results impossible. As
mentioned previously, the nitrogens in PPAN may not be
capable of producing a nitrogen hyperfine interaction, thus

85
making the EPR method unsuitable for determining whether the
copper is bound to nitrogen. In any case, the fact that the
copper signal in PPAN*CuC12 was significantly different from
that of CuCl2, coupled with the fact that repeated washing
of the PPAN with HC1 and distilled water failed to alter or
reduce the intensity of the signal, implies that the copper
was indeed coordinated to the surface of the PPAN, although
it was not possible to determine whether the copper was
bound to nitrogen, carbon, or oxygen.
Bis (2,2'-bipyridine) ruthenium (II)
trifluoromethanesulfonate (Ru (bipy)2(CF3SO3)2)
This compound was prepared in situ by dissolving 0.11 g
of AgCF3S03 in an ethanol solution containing 0.10 grams of
Ru(bipy)2Cl2, heating the reddish solution to boiling to
coagulate the AgCl, and finally filtering and washing with
absolute ethanol. The deep red filtrate was presumed to
contain Ru(bipy)2(CF3SO3)2. The UV-visible spectrum of a
solution made in a volumetric flask was recorded before and
after the addition of a large excess of PPAN (5.10 grams) It
was found that there were no appreciable absorptivity
changes in the UV-visible spectra, neither in the wavelength
maxima or intensities of the absorptions, indicating that
there was negligible adsorption or binding of
Ru(bipy)2 (CF3SO3)2 to the surface of the PPAN. This
particular ruthenium complex was chosen for several reasons.

86
Being a very poor ligand and hence a good leaving group, the
trifluoromethanesulfonate anion seemed like a good candidate
for being displaced by PPAN, which had not proved to be a
particularly good ligand in previous experiments. In
addition, it was hoped that the presence of two bidentate
2,2'bipyridines bound to the ruthenium would facilitate (by
the "chelate effect") the binding of PPAN as a bidentate
ligand, thus lessening the steric strain associated with the
large PPAN structures. This strategy proved to be
ineffective since no adsorption of ruthenium complex was
detected.
Catalytic Results
The catalytic studies fall into two general catagories:
screening reactions were carried out to determine whether
PPAN catalysts were active towards various substrates, and a
specific reaction, that of ethanol over PPAN catalysts, was
used as a model reaction to study the effects of pyrolysis
conditions, metal dopants and oxide supports on the
catalytic activity of PPAN.

87
Screening Reactions
Isopropyl alcohol and ethylbenzene
Some preliminary experiments with early catalyst
preparations were undertaken to verify the catalytic
activity reported in previous studies. Manassen and
Wallach15 had reported the catalytic activity of PPAN
materials towards isopropyl alcohol. Therefore, the vapor
was passed over PPAN at 150°C and although the results were
not quantified, acetone was detected by gas chromatography.
As a preliminary probe of the photocatalytic capabilities of
PPAN, suspensions of PPAN in isopropyl alcohol were
irradiated with visible and ultraviolet light using a medium
pressure mercury vapor lamp with pyrex and quartz reaction
vessels, respectively. The results, when compared to the
appropriate blank runs (no PPAN present), indicated that
PPAN had no appreciable photocatalytic activity towards
isopropyl alcohol.
The same series of experiments were carried out using
ethylbenzene as the substrate and the results were analogous
to those obtained for isopropyl alcohol. In the vapor
phase, ethylbenzene was converted to styrene, as previously
reported by Degannes and Ruthven,19 and no photocatalytic
reaction was observed to occur. The results of these
experiments served to verify that although the measured
surface areas were somewhat smaller, the catalytic

88
activities of the materials prepared in this laboratory were
comparable to those of earlier published reports.
Norbornene oxidation
It has been previously reported that PPAN is an active
dehydrogenation catalyst, however there have been very few
studies in which PPAN has been tested with respect to its
oxidative capabilities. Towards this end norbornene was
used as a gaseous substrate at 140°C and atmospheric
pressure, as well as at 50°C and 35 psi oxygen in a
suspension containing 1 gram of norbornene and 1 gram of
PPAN in 20 ml of acetonitrile. In both cases, there were no
detectable changes in composition under the reaction
conditions as measured by gas chromatography using a
carbowax column and a flame ionization detector. There are
several possible products which could conceivably be formed
from norbornene, namely norbornene oxide, norbornadiene, and
numerous products resulting from the fragmentation of the
bicyclic ring system. Norbornadiene would result from the
dehydrogenation of the carbon atoms symmetrically related by
a mirror plane to the doubly bonded carbons in norbornene.
Since PPAN has been shown to dehydrogenate cumene and
ethylbenzene, which both have a bulky aromatic system
conjugated to the bond being dehydrogenated, it was felt
that norbornene might be a good candidate to test whether
the existence of a 7r-allyl type system is necessary for

89
dehydrogenation to occur. On the other hand, one must
remember that norbornene, due to its bicyclic structure, has
a far greater steric problem than cumene or ethylbenzene.
If steric factors are of the most significance, norbornene's
lack of reactivity is readily explained since all of the
previously proposed structures for PPAN catalysts are quite
large and bulky, which would render them less suitable for
interaction with a sterically hindered substrate such as
norbornene. The conversion of norbornene to norbornene
oxide has been accomplished using ruthenium (II)
phenanthroline catalysts.53 The fact that norbornene is
stable in the presence of PPAN implies that either
norbornene is too sterically hindered to react with PPAN, or
PPAN is not a particularly good oxidation catalyst or both.
These rationalizations are supported by the available
literature reports since the use of PPAN as an oxidation
catalyst has been reported only twice (in one case, silica
and metal salts were also present, while the other report
employed PAN for the purpose of silver crystallite
deposition on the surface of oxide supports21) and the
polymeric structure of PPAN has been fairly well
characterized as being a large extensively crosslinked
heterocyclic structure.

90
Syngas conversion
The hydrogenation of carbon monoxide with dihydrogen is
an industrially important reaction since mixtures of
hydrogen and carbon monoxide, called syngas, are a major
by-product of the petroleum mining and refining industry.
Unlike carbon dioxide, carbon monoxide is fairly reactive
due to its valence deficient electronic structure, as
evidenced by its tendency to form numerous metal-carbonyl
complexes. The catalytic mechanism of PPAN has been
described in the literature as one in which the catalyst
dehydrogenates a substrate and subsequently becomes reduced.
The catalyst is oxidized to its original state by a short
treatment in oxygen at reduced temperatures and has been
shown to produce water. Therefore, if carbon monoxide were
to participate in this catalytic cycle, it would have to
replace the oxidant since a product further reduced would be
undesirable except for emission control applications. By
using a mixture of hydrogen and carbon monoxide it was hoped
that the PPAN could adsorb the hydrogen molecules and then
catalytically reduce the carbon monoxide, completing the
catalytic cycle. A 1:2 mixture of carbon monoxide and
hydrogen was passed over a nitrogen pyrolyzed sample of PPAN
at temperatures up to 300°C and the only product observed
using gas chromatography was a small methane impurity which
was contained in the original carbon monoxide. Since
ruthenium is frequently used as a Fischer-Tropsch catalyst,

91
samples of ruthenium doped PPAN were prepared by refluxing
PPAN in a methanol or ethanol solution of RuCl3*3H20 for 24
hours. The mixture was then filtered and washed with
methanol or ethanol and dried under a stream of nitrogen.
When a mixture of carbon monoxide and hydrogen was passed
over the catalyst prepared in methanol using the same
catalytic conditions as before, dimethyl ether was the only
observable product, whereas the catalyst prepared in ethanol
produced only acetaldehyde, diethyl ether and ethyl acetate,
as determined by GC analysis. These results indicated that
the only reactions occurring involved the adsorbed solvent
molecules since the same results were obtained using
nitrogen as the carrier gas instead of syngas. The fact
that dihydrogen is not active towards PPAN catalysts is not
too surprising since Manassen and coworkers16 observed that
gaseous hydrogen was never produced during the catalytic
reaction or after prolonged heating in an inert atmosphere.
Their work indicated that hydrogen binds PPAN as a species
intermediate between a hydrogen atom and a hydride ion.15
Since the residual ethanol solvent produced mostly
acetaldehyde, it is reasonable to assume that the surface of
PPAN must be becoming reduced or hydrogenated. This is
substantiated by the fact that a short air treatment would
restore the original catalytic activity with the concomitant
production of water. Therefore, ethanol and carbon monoxide
were passed over PPAN simultaneously to determine whether

92
the reduced PPAN was capable of hydrogenating carbon
monoxide. The results obtained indicated that there was no
difference in product distribution between reactions carried
out with carbon monoxide or nitrogen atmospheres at
temperatures as high as 300°C, implying that PPAN was not
capable of hydrogenating carbon monoxide. Recent results by
other workers employing a ruthenium doped PPAN catalyst have
shown that these materials are capable of undergoing
Fischer-Tropsh type chemistry producing low molecular weight
alkane and alkene products.54
Polymerizations
Since PPAN is paramagnetic and generally contains a
large number of unpaired spins per gram, it seemed like a
good candidate for a free radical polymerization catalyst.
In free radical polymerization, an initiator is usually
required to generate enough free radicals to start the chain
reaction. Only a small amount of initiator is necessary
since each step in the reaction involves the consumption of
a free radical accompanied by the production of a new,
larger free radical until the reaction is terminated by a
step such as disproportionation, which consumes without
producing free radicals. Since the EPR spectrum of PPAN
most closely resembles that of an organic free radical,
albeit with a larger linewidth, the use of PPAN based
materials as free radical polymerization initiators was

93
investigated. Polymerizations were attempted using pure
styrene as well as solutions containing 10 mL of styrene, 10
mL of methyl methacrylate and 25 mL of carbon tetrachloride.
In the former case, 0.5 grams of PPAN were added to 25 mL of
styrene; in the latter case, 1.0 gram of PPAN was added to
the solution. There was no noticeable evolution of heat
upon addition of the PPAN in either case. The suspensions
were allowed to stir overnight at room temperature, after
which time NMR samples (after removing the PPAN) were
prepared. The proton NMR results indicated that no
appreciable polymerization had occurred. In addition, a
styrene-methyl methacrylate-PPAN mixture was irradiated with
pyrex filtered light from a medium pressure mercury lamp,
and the NMR results again indicated that no polymerization
had occurred. Since the bandgap of PPAN has been measured
to fall in the visible region (approximately 1.5-3.0 eV55),
it was also surprising that strong excitation with visible
light failed to produce any free radical products. These
results suggest that although the EPR signal associated with
PPAN resembles that of an organic free radical, the chemical
reactivity of PPAN must be fundamentally different from that
of organic free radicals or peroxides which are known for
initiating these polymerizations. Of the numerous types of
polymerization mechanisms, free radical polymerization and
coordination polymerization, illustrated schematically in
Figure 28, would seem to be the most likely mechanisms of

94
FREE RADICAL
Rad• + CHZ=CH
Rad C — CH
Rad CH0—CH* + CH==CH
Ra d C H2— C H — C H2— C H
etc.
COORDINATION
Cat — C HnC H,
f 2 o
CH =C H
2 2
Cat- CH2CH2CH2CH3
C H =C H
2 2
etc.
Figure 28: Polymerization Mechanisms

95
activity for PPAN catalysts. Both mechanisms involve
coordination or bonding to the catalyst surface during chain
growth, although they differ in that coordination
polymerization involves the addition of monomer at the
catalyst surface, thus enabling stereoselective control,
whereas free radical polymerization involves the migration
of the radical down the polymer chain as monomer units are
added, resulting in a non-stereoselective product. In view
of the above mechanisms, there are several possible
explanations for PPAN's lack of reactivity, and all can be
generally related to the physical or electronic structure of
PPAN. It is possible that steric constraints imposed by the
PPAN structure preclude interaction with these monomer
units, although the fact that PPAN is capable of
dehydrogenating ethylbenzene in the vapor phase implies that
it is capable of interacting with styrene to some extent.
It is also of interest to note that Degannes and coworkers19
reported no polystyrene production in the ethylbenzene
dehydrogenation reaction. In fact they observed no side
products at all and concluded that the removal of hydrogen
from the surface was very rapid, and that the overall rate
was controlled by the rate at which the catalyst can
abstract hydrogen from ethylbenzene.19 It should also be
noted that PPAN cannot dehydrogenate ethylbenzene under the
conditions used for the polymerization reactions. The
possibility that the chain transfer rate of carbon

96
tetrachloride was large enough to make the polymerization
rate insignificant was ruled out by attempting the
polymerizations in neat styrene and methyl methacrylate and
observing no products. In summary, it has been shown that
PPAN does not seem to behave like an organic free radical
even though the similarity between the EPR spectra strongly
suggests this.
Chloroethvl ethylsulfide (CEES)
The decomposition of CEES is of military importance due
to its structural similarity to mustard gas
(bis(2-chloroethyl) sulfide). In fact CEES is frequently
referred to as "one armed mustard." It was of interest to
determine whether PPAN catalysts were capable of decomposing
CEES to any extent. The catalytic set-up used was that
shown in Figure 9 with CEES in the first bubbler and Clorox
bleach in the exit bubbler. Bleach is known to detoxify
mustard gas and is the principal decontaminating agent
presently used by the U.S. Army. The products were analyzed
by GC, GC-IR and GC-MS, and although some of the product
assignments may be uncertain, it is clear that PPAN
catalysts are quite active towards CEES in that
approximately 30 to 40% of the CEES was decomposed at 200°C
using nitrogen as the carrier gas. A sample of the products
was trapped using a dry ice/acetone bath, and GC-IR analysis
indicated that hydrogen sulfide and carbon disulfide were

97
the major products that could be identified using the
Georgia State Crime Lab Library and the EPA vapor phase
library. The gas chromatograms also indicated the presence
of several additional minor products. After running for
several days at 200°C, there was very little product
formation and the catalyst appeared to be dead. However,
increasing the temperature to 300°C caused the decomposition
effeciency to increase to approximately 40%, although it is
possible that some of this increase at the higher
temperature was due to the desorption of adsorbed products.
In addition, using air as the carrier gas also served to
restore the original activity at temperatures as low as
140°C. These experiments were repeated at the CRDEC,
Aberdeen Proving Ground, in an attempt to identify the minor
products using GC-MS. Table 3 contains the results as well
as possible structures of the products based on a library
search of the Aldrich Library of Mass Spectra. The
percentages in parentheses were taken 1 day later after
sitting at room temperature overnight, indicating that at
room temperature the decomposition products of CEES are
unstable and subject to rearrangement. In particular, the
higher molecular weight polythioethers, as well as
ethyldisulfide and thiophene, seem to decompose upon
standing, while the areas of the first peak and the
tetrahydrothiophene increased. The nearly 6% increase in
the concentration of tetrahydrothiophene can not be

Table 3: CEES Products
Peak
Retention
Percent of
M/Z
Possible
Number
Time (sec)
Total
Structures
1
41
66.11
(69.62)*
32
Air
64
so2
76
cs2
62
ch3ch2sh
2
68
2.41
(0.72)*
84
Thiophene
3
79
5.14
(11.94)*
88
Tetrahydrothiophene
4
223
11.40
(6.48)*
122
(CH3CH2S)2
5
228
10.47
(10.76)*
124
CEES
6
263
0.24
(-)*
105
Ethynyl
pyridine
7
374
3.76
(")*
150
(CH3CH2S)2chch3
8
465
0.27
(-)*
150
ch3ch2sch2ch2sch2ch3
* After sitting in the trap at room temperature overnight

99
accounted for by the hydrogenation of thiophene. Therefore
it was most likely formed by the fragmentation and
cyclization of the polythioethers which would also result in
the production of ethyl sulfhydryl. The increase in the
percentage of the air peak may be ethyl sulfhydryl from the
scission of both the disulfide and the polythioethers;
however in the absence of conclusive evidence, such
speculation is of limited value. The percentage
decomposition of 40% reported earlier was based on gas
chromatographic analysis using a flame ionization detector
and a chromasorb packed column, while the GC-MS experiment
used a 15 meter SE-54 capillary column and was analyzed
using a thermal conductivity detector. Since flame
ionization detectors do not fully detect oxidized compounds,
the only detectable product in the first peak eluted would
have been ethyl sulfhydryl. Most likely this was what the
GC-IR library search identified as hydrogen sulfide since
neither of the libraries searched contain ethyl sulfhydryl.
The experiments performed at the University of Florida and
the CRDEC, Aberdeen Proving Ground, involved the use of
different catalyst preparations, a different purity of CEES
and different reaction conditions. In all cases, however,
the decomposition effeciency was quite good, on the order of
about 40% at 200°C and 80% decomposition at 300°C. It
should be remembered that the space velocity of the CEES was
very low and none of the integrated areas have been

100
corrected for detector sensitivity since no internal
standard was used. Using higher space velocities would
undoubtedly decrease the decomposition efficiency.
Nonetheless these preliminary results, obtained with an un¬
optimized catalyst, were encouraging.
The elucidation of the reaction mechanism(s) would be
impossible without detailed knowledge of reaction kinetics
and thermodynamics. Two significant observations can be
made on the basis of the data presented: there were no
chlorine containing products accounted for and there were no
products with an odd number of carbon atoms. The fact that
there were no chlorine containing products was evidenced by
the lack of the characteristic chlorine isotope pattern,
suggesting that the chlorine must be bound to the surface of
the catalyst, possibly accounting for its slow deactivation
over a period of weeks under reaction conditions. If this
was the case, then washing the catalyst with water or some
other solvent to remove bound chlorine (which could be in
the form of bound hydrochloric acid since this and other
studies have demonstrated PPAN's ability to slowly absorb
acids), could regenerate the catalyst. The fact that all of
the products contain an even number of carbon atoms suggests
that the ethyl group (with or without the sulfur) remains
intact during the catalytic reaction, while the formation of
the disulfide indicates that the -SCH2CH3 moiety was a
probable reaction intermediate. The production of ethynyl

101
pyridine could possibly be an artifact of the
instrumentation or, more likely, it could result from the
decomposition of the catalyst. In any event, the fact that
the molecular weight was odd numbered strongly suggests the
presence of nitrogen, with the only logical sources being
the catalyst or molecular nitrogen from the atmosphere. The
products obtained from the thermal reaction of CEES over
PPAN were strikingly different from those obtained from the
solution hydrolysis of CEES in mixtures of water and various
organic solvents in that there were no oxygenated products
detected. Mustard gas hydrolyzes to form hydrochloric acid
and mustard chlorohydrin (2-hydroxyethyl ethylsulfide),
which subsequently hydrolyzes to form thiodiglycol and
another molecule of hydrogen chloride. CEES, on the other
hand, hydrolyzes in one step to form one molecule of
hydrogen chloride and one molecule of 2-hydroxyethyl
ethylsulfide. The mechanism for this reaction has been
shown to involve a cyclic sulfonium ion intermediate,56 and
the solvent effect on the rates indicated an SN-1
mechanism.57 Since the reaction over PPAN was carried out
partly in the presence of air, there should have been water
present since it is known to be a product of the oxidative
regeneration of the catalyst. It was therefore most
surprising that oxygen failed to turn up in any of the
products. One possible explanation was that the chlorine,
presumably bound to the surface of the PPAN since none was

102
observed in the products, poisons the surface of the
catalyst by preventing the oxidative regeneration of the
active catalytic species. It should be noted that PPAN was
not demonstrated to be catalytic with respect to the
decomposition of CEES since only a relatively small amount
could be passed over the PPAN sample using a bubbler as the
source of substrate feed. The results are encouraging,
though, and since the observed products were guite different
from the normal hydrolysis products, this system merits
further study as a means of mustard gas decontamination.
Dimethyl methvlphosphonate (DMMP)
This compound is also of interest to the military since
it is freguently used as a simulant for a general class of
compounds referred to as G-Agents, or nerve gases, all of
which have the common feature of a central phosphorous atom
with various organic substituents. The structural formulas
for DMMP and various nerve agents are shown in Figure 29 for
comparison. It can be seen that DMMP lacks the presence of
a pseudohalogen or halogen (specifically fluorine or
cyanide) bound to the central phosphorous atom, while all
contain a phosphite ester or phosphorl group. The
experimental set-up used to test the catalytic activity of
PPAN towards DMMP was the same as that used for the CEES
experiments and is schematically illustrated in Figure 9.
The influent bubbler contained neat DMMP and the effluent

103
H3C
0CH-
I
P=0
I
OCH.
H C
3 \
H3C
N —
/ I
0
II
P-O-C H
2 5
CN
DUMP
Dimethyl methyl-
phosphonate
labun (GA)
Ethyl N,N-di methy1 -
phosphoramidocyanidate
F CbU
I I 5
H_C—P—0—C-C(CH_)
3 || | 3
0 H
Soman (GD)
Pinaccolyl methyl-
phosphonofluoridate
F CH3
I I °
H C-P-O-C-H
3 ii |
0 CH3
Sarin (GB)
Isopropy1 methy1 -
phosphonofluoridate
Figure 29: Structures of Dimethyl Methylphosphonate
and Nerve Gases

104
bubbler contained bleach to hydrolyze the unreacted DMMP.
In the temperature range of 100-210°C under aerobic
conditions, both on-line GC analysis as well as GC-MS of the
products trapped in a dry ice/acetone bath indicated that no
decomposition of the DMMP had occurred since the only
compound detected was the DMMP starting material (barring
minor impurities). The catalyst used for this experiment
was pyrolyzed using a programmable temperature controller in
an air atmosphere. In fact the sample was from the same
batch as used in the CEES reaction, thus eliminating the
possibility that this particular sample of catalyst was
inactive. The results suggest that PPAN catalysts are
unsuitable candidates for use in the decomposition of DMMP;
however the possibility exists that PPAN may be effective on
live agents as a result of their differences from DMMP in
chemical reactivity due to the presence of fluorine or
cyanide substituents. It is also quite possible that PPAN
supported on alumina would be active towards DMMP since it
has been shown, using inelastic electron tunnelling
spectroscopy, that alumina is capable of adsorbing DMMP
dissociatively at temperatures above 20°C. At temperatures
above 200°C, decomposition occurs by way of dealkylation of
the phosphorous.58'59 Since platinum supported on titania
and on alumina have also been shown to decompose DMMP,60 it
is possible that the addition of platinum to PPAN based
catalysts could render these materials active towards DMMP.

105
Ethanol Catalysis
In this series of experiments, the reaction of ethanol
over PPAN based materials was used as a test reaction to
determine whether the presence of metal additives or the
attachment of PPAN to the surface of oxide supports (such as
alumina, silica, and titania) had any effect on either the
activity or product selectivity of these catalysts.
Metal additives
Through the course of these experiments, numerous
methods were employed for the deposition of metal species
onto the surface of PPAN catalysts. For example, lithium
chloride and copper were incorporated into PPAN catalysts by
dissolving lithium chloride in a dimethylformamide (DMF)
solution containing PAN, and then copper was added to the
solution either after or before the stripping of the DMF by
rotary evaporation. In Method A, where copper metal (finely
divided powder) was added directly to the DMF solution of
PPAN, all of the copper dissolved yielding a brown solution
after about 2 weeks. This solution was subsequently rotary
evaporated to yield a hard brown solid which was then
pulverized in a blender. In Method B, the dried PAN
containing lithium chloride was pulverized and mixed with
copper powder, and after about 2 weeks with occasional
mixing, all of the copper disappeared resulting in a light
green solid. These materials were subsequently pyrolyzed as

106
previously described. This preparation technique was
eventually abandoned for several reasons. Most importantly,
the goal of these experiments was to compare the activities
of catalysts containing several different metal additives;
however, of all the metals tried, (copper, cobalt, iron and
zinc), copper seemed to be the only metal capable of
incorporating into PPAN containing lithium chloride in this
unusual manner. Another reason for discontinuing this line
of experimentation was that the catalytic activity of these
preparations proved to be very difficult to reproduce from
one batch to the next. Also, the surface areas of these
preparations were very low, probably as a result of
dissolving the PAN in DMF followed by rotary evaporation.
Techniques such as spray drying, plasma polymerization, and
cryogenic grinding (which are not available at the
University of Florida) are capable of producing solids with
much higher surface areas than can be produced by rotary
evaporation followed by grinding in a blender. Another
problem associated with this preparative technique was the
inability to completely remove all of the DMF from these
materials (even after evacuation in a desiccator for 2
weeks), due to its high polarity and low vapor pressure. It
can therefore be assumed that when these materials were
pyrolyzed, there were significant amounts of DMF present,
and it is not known what effect this would have on the
outcome of the pyrolysis reaction. There were two types of

107
evidence, mass balance results and X-ray Photoelectron
Spectroscopy (XPS) results, which indicated that DMF was
incorporated into the starting material. The mass before
pyrolysis was greater than the sum of the weights of PPAN,
lithium chloride and copper used, indicating the
incorporation of DMF or some other material from either the
atmosphere or the pyrex vessel. Films for XPS analysis were
prepared containing PAN and lithium chloride as well as PAN,
lithium chloride, and copper (using both preparative methods
outlined earlier for copper incorporation). The XPS
analysis, performed courtesy of Dr. Willie Hendrickson of
the 3M corporation, were intriguing because they proved to
be somewhat of a mystery in their own right. In the lengthy
process of getting the samples to 3M and then receiving the
results, the samples were mixed up and it is not known which
of the four copper containing samples was the one prepared
by adding the copper metal to the DMF solution containing
lithium chloride and PAN. Had the samples not been
discarded after analysis it would have been easy to identify
since unlike all the other samples which were clear to light
green, this sample was dark brown and nearly opaque.
Table 4 consists of relative elemental ratios which were
calculated on the five surfaces examined, while Table 5
lists the binding energies for the ionizations of interest.
Both the original spectra as well as the high resolution
spectra are contained in the Appendix. There are several

108
TABLE 4: XPS Results; Elemental Ratios
Number Sample
O/C
N/C
Cl/C
Cu/C
Li/C
1 LiCl,
PAN
. 18
. 12
. 04
-
. 04
2 LiCl,
Cu, PAN
. 17
. 17
. 02
. 002
. 03
3 LiCl,
Cu, PAN
. 25
. 11
. 03
. 005
. 04
4 LiCl,
Cu, PAN
. 15
. 19
. 03
. 01
. 00
5 LiCl,
Cu, PAN
.45
.03
. 10
. 04
. 04

109
Table 5: Ionization Energies (eV)
Sample
0 (KLL)
Cu (2P)
0 (IS)
N (IS)
C (IS)
1.
LiCl, PAN
987.98
—
540.08
407.82
293.16
2 .
LiCl, PAN,
Cu
975.95
989.91*
931.00
532.06
542.75*
398.46
409.15*
285.19
295.82
3.
LiCl, PAN,
Cu
979.95
934.54
532.06
399.80
285.19
4.
LiCl, PAN,
Cu
978.62
934.54
532.06
399.80
285.19
5.
LiCl, PAN,
Cu
977.29
934.54
532.06
399.80
285.19
6.
CuCl**
—
935.0
—
—
—
7.
PAN-3**
—
—
531.9
399.0
—
8.
PAN-3-Cu**
932.4
531.8
399.6
* Analysis was performed on the side untreated with copper
** This data was compiled from reference 19.

110
conclusions which can be drawn from this data about the
effects of copper incorporation in these systems. First of
all, there is a dramatic change in the binding energies of
the carbon, nitrogen and oxygen electrons upon copper
incorporation. This is illustrated by the results for both
sides of sample two contained in Table 5. In method A, the
copper metal was placed in contact with only one side of the
film, therefore the binding energies of one side should
resemble sample one (no copper) while the binding energies
of the other side should resemble those of the other copper
containing samples. These predictions are born out by the
data in Table 5. This also proves that the sample prepared
by method B must be either sample three, four or five. Of
these three samples, number five was most likely the one
prepared by Method B since it was significantly different
from any of the other spectra in that there seemed to be
considerably more copper and oxygen on this surface as well
as less carbon and nitrogen, presumably due to the greater
copper loading. Since this sample was presumably prepared by
Method B, where the copper metal was added to the DMF
solution, it is possible that DMF incorporation or possibly
even coordination to copper could be responsible for the
increase in oxygen concentration at the surface. It must be
remembered however that the XPS technique is only capable of
providing a chemical analysis of the outermost 30 to 60 A
angstroms of a given surface. Therefore the ratios

Ill
presented in Table 4 do not necessarily represent the
elemental ratios in the bulk sample. The copper present in
the samples also appears to be exclusively in the +2
oxidation state. This was not surprising for the green
samples since green color is frequently characteristic of
copper in the +2 oxidation state; however, the fact that the
brown sample is in the same oxidation state leads one to
suspect that the copper in the two types of preparations
exists in markedly different coordination environments. It
is of interest to note that anhydrous cupric chloride
(eriochalcite) is a brownish yellow while hydrated cupric
chloride is green. Cupric oxides (paramelaconite and
tenorite), on the other hand, are black. Another
observation which can be extracted from Table 4 is that
samples one, two and three all have lithium and chloride
present in a one to one ratio, while samples four and five
have more chlorine than lithium on their surfaces.
Coincidentally, these two samples also have the highest
copper concentrations at the surface, which suggests that
some type of copper-chloride species exists at the surface.
However, this does not account for the lithium which is
known to be present in these samples. Most likely, the
lithium excess in the interior is ion-paired with a species
derived from the polymer backbone or it exists imbedded in
the polymer matrix as the hydroxide.

112
The XPS results are also of interest for comparison to
the results reported by Bai and coworkers20 on their silica
gel supported PPAN catalysts. The only XPS results reported
in their communication (also contained in Table 5) were
those for cuprous chloride and a silica supported PPAN
catalyst before and after treatment with cuprous chloride.
This material was prepared by polymerizing acrylonitrile in
the presence of silica gel and then pyrolyzing the material
at 190°C for 12 hours. In their analysis of the XPS data,
it was concluded that the copper was bound to nitrogen in
the polymer based on the differences in the binding energies
of the nitrogen IS and copper 2P electrons (0.6 and 2.6
electron volts, respectively). However, upon comparing
these differences to those obtained for the PAN films, it
can be seen that variations larger than these exist between
films of differing copper concentrations, with elements in
addition to copper and nitrogen (most notably oxygen).
Therefore it is highly doubtful that these differences in
binding energies constitute conclusive evidence of copper
nitrogen bonding. In addition, other studies using copper
on silica without PAN reported a value of 932.7 eV for the
copper 2P 3/2 ionization,61 which is very close to the value
reported by Bai and coworkers, casting further doubt as to
whether their XPS data can be used as evidence for copper-
nitrogen binding.

113
Another general method of metal incorporation involved
the adsorption of metal salts from solution. However, as
the metal binding studies indicated, it was not possible to
observe significant metal uptake from solution. In
addition, since different metal species would tend to have
different adsorption characteristics, it was difficult to
compare the effects of eguimolar amounts of different metal
species using adsorption as the deposition method. The
final method chosen (as described previously) involved the
rotary evaporation of suspensions containing known amounts
of PPAN and the metal species. Two series of catalysts were
prepared in this manner and tested for catalytic activity
towards ethanol. The members of each series were all
prepared from the same batch of PPAN with the goal of
minimizing the differences caused by variations in the
pyrolysis reaction since earlier results indicated that the
elemental compositions tended to vary somewhat from one
pyrolysis to the next. Table 6 summarizes the catalytic
results for the first series of metal doped PPAN catalysts,
and the results for the undoped PPAN catalyst which was used
to prepare this series have been included for comparison.
The particular batch of PPAN used to prepare this series was
pyrolyzed in a nitrogen atmosphere using a manually operated
temperature controller and a relatively fast heating rate.
The results clearly indicate that the addition of metal
additives can affect both the catalytic activity as well as

114
Table 6: Effects of Adsorbed Metal Salts, Series 1
Catalyst
Conversion
(%)
Acetaldehyde
(%)
Diethyl
Ether
(%)
Ethyl
Acetate
(%)
PPAN
. 4
.4
—
—
PPAN-CuC12
. 06
. 06
—
—
PPAN-Ni(NO3)2
2-12*
2-10*
.2*
2.0*
PPAN-Fe(NO3)3
2.6
. 6
trace
2.0
PPAN-Zrx(CH3COO)
y 1.7
. 3
. 7
.7
PPAN-Zrx(C03)v
1.4
1.4
trace
—
* Produced during air regeneration.
WHSV= 0.6
Temperature= 270°C (except PPAN-CuC12 at 180°C)
3.45 X 10-4 moles of metal/gram of PPAN

115
the selectivity of PPAN catalysts. For example, the overall
activities of the nickel (II) nitrate and iron(III) nitrate
were roughly equivalent within the limits of experimental
error, although the iron doped catalyst was more selective
towards ethyl acetate while the nickel catalyst was more
selective towards acetaldehyde. Another general observation
which can be made is that the addition of metal species
increases the catalytic activity in all cases except
copper (II) chloride. The lower reaction temperature in
this case probably accounts for the lower reactivity. In
addition, it was observed that the use of air as the carrier
gas instead of nitrogen results in a several-fold increase
in overall activity accompanied by a decrease in
selectivity, as evidenced by the results for the nickel (II)
nitrate doped catalyst. Using nitrogen as the carrier gas
resulted in the production of acetaldehyde exclusively with
an approximate conversion of 2% while the use of air as the
carrier gas resulted in the production of approximately 10%
acetaldehyde as well as significant quantities of diethyl
ether and ethyl acetate. It is possible that proportional
amounts of the minor products were produced using nitrogen
as the carrier gas but the quantities were small enough to
escape detection. Although the nickel catalyst was the only
member of this series which was subjected to air as the
carrier gas, subsequent experiments with other PPAN

116
catalysts demonstrated the general validity of this
observation.
Table 7 contains the catalytic results for a second
series of metal containing catalysts as well as the results
for the PPAN starting material. As was shown in the first
series of catalysts, the addition of metal species always
increased the conversion efficiency while the use of air as
the carrier gas also served to increase the conversion
efficiency (generally by greater than an order of magnitude)
at the expense of selectivity in most cases. Variations in
the metal species also affected the selectivity as well as
the activity of these catalysts. For example, the
manganese (II) chloride doped catalyst produced exclusively
acetaldehyde and diethyl ether under aerobic conditions,
while the chromium (II) chloride doped catalyst produced
acetaldehyde, diethyl ether, and ethyl acetate with a much
different product distribution. The PPAN used to prepare
these catalysts was pyrolyzed using a programmable
temperature controller set to the program described earlier.
Another difference between this batch of PPAN and the one
used for the first series is that this was pyrolyzed in an
air atmosphere while the other was pyrolyzed in a nitrogen
atmosphere. A comparison between the catalytic activities
of the two undoped starting materials reveals very little
difference in either conversion efficiency or product
selectivity. In general, however, the air pyrolyzed samples

117
Table 7: Effects of Adsorbed Metal Salts, Series 2
Conversion Acetaldehyde Diethyl Ethyl
Ether Acetate
(%)
(%)
(%)
(%)
Dopant
n2
air
n2
air
n2
air
n2
air
0.25
3.0
0.25
3.0
0.0
0.3
0.0
0.4
MnCl2
0.3
9.0
0.3
9.0
0.0
2.0
0.0
0.0
CoCl2
0.5
6.0
0.5
6.0
0.0
0.3
0.0
1.0
CrCl2
0.3
13.0
0.3
11.0
0.4
0.5
0.3
1.5
FeCl2
1.0
15.4
1.0
14.0
0.0
0.2
0.0
1.2
WHSV= .6
Temperature= 270°C
3.45 X 10-4 moles metal/gram PPAN

118
seem to be slightly more active towards ethanol, as these
and other results indicate. For example, two identical
samples of PAN were pyrolyzed in air and nitrogen using the
exact same temperature program, and the air pyrolyzed sample
was slightly more catalytically active.
Table 8 contains catalytic data as well as the
elemental analyses of several earlier preparations,
including some results for the copper-lithium catalyst whose
preparation was described earlier. Although it is difficult
to make accurate comparisons since there are differences in
the reaction temperatures as well as the space velocities,
one can see that the performance of these catalysts is
roughly equivalent to that previously presented for the
undoped catalysts in series one and two. The fact that
these pyrolyses were carried out in atmospheres of air,
nitrogen and ammonia, yielding catalysts of comparable
activity with different elemental compositions, suggests
that the formation of the catalytic moiety contained in PPAN
is independent of the pyrolysis atmosphere and intrinsic to
the pyrolysis of polyacrylonitrile. It is also of interest
to note that the spent catalysts invariably contains a
higher percentage of carbon than the fresh catalysts
implying that one possible mode of catalyst deactivation may
be that the PPAN surface becomes covered with coke or other
carbon containing fragments.

119
Table 8: Miscellaneous Catalyst Results
Sample
o\o
u
H%
N%
M2/g
Conversion
WHSV
T (C)
PAN
67.12
5.82
26.17
—
—
—
—
PPAN-air
59.47
2.40
23.21
. 29
0.4
0.6
180
Dead
61.65
2.42
23.52
—
—
—
—
ppan-n2
69.39
3.78
21.59
4.45
0.2
0.6
170
PPAN-CuLi
60.20
4.07
18.26
0.24
1.25
0.3
250
Dead (air)
62.96
4.35
19.22
—
0.04
0.6
170
Dead (N2)
61.84
3.90
19.28
1.20
0.03
0.6
170
PPAN-NH3
69.94
3.15
22.38
—
0.4
0.6
270

120
Several control experiments were also carried out to
determine whether the dispersion of the metal species on an
inactive support, such as 8-58 /xm glass spheres, could
result in a material of comparable catalytic activity to the
corresponding metal species supported on PPAN. The goal was
to use the same amount of metal species per unit of surface
area as in the metal doped PPAN experiments. Towards this
end, a rough measurement of the surface area of the glass
spheres was carried out using the following method. Samples
of PPAN (with a known surface area) and glass spheres were
placed in a vacuum oven at 45° C overnight. The samples
were weighed immediately upon removal from the oven. The
samples were then placed in a dessicator saturated with
water vapor and the percentages of weight increase (due to
water adsorption), were determined by reweighing the samples
after 24 hours. The results of these experiments indicated
that the surface area of the glass spheres was somewhere
between about 1.2 M2/g and 8.9 M2/g since the PPAN sample
had a percentage weight gain of approximately 12% while that
for the glass spheres varied from approximately 2% for
untreated spheres, to 12% weight gain for spheres which had
been doped with Ni(N03)2*6H20. Since Ni(N03)2’6H2O is
hygroscopic, it is possible that some of the weight gain in
the surface area experiments could have been due to the
rehydration of the metal complex, however the proportion of
metal complex to support is small enough to be negligible

121
for practical purposes. At any rate, the value of 1.2 m2/g
obtained for the untreated spheres should serve as a lower
limit for the surface area of the glass spheres. Catalytic
experiments were carried out using 35 g of glass spheres
which were washed with ethanol and then rotary evaporated,
35 g of glass spheres containing 3.45 x 10-4 moles of
FeCl2*3H20, and 5 g of glass spheres containing the same
number of moles of Ni(N03)2•6H20. In addition, all
experiments were carried out under nitrogen and air
atmospheres. The results of these experiments are contained
in Table 9. If one calculates the amount of surface area
covered by 3.45xl0-4 moles of a complex one obtains an area
of approximately 75 m2 assuming that a metal complex would
cover approximately 36.0 A2 of surface area. This indicates
that these experiments employ at least a monolayer of metal
dopants both with PPAN and glass spheres preparations (the
PPAN would probably have more monolayers). In the cases
where 35 grams of glass spheres were used, a significant
amount of conversion was observed. The conversion obtained
using undoped glass spheres was approximately 1.0% to
acetaldehyde under both air and nitrogen atmospheres, while
using 35 grams of spheres doped with FeCl2*3H20 resulted in
conversions of 1% and 9% under nitrogen and air,
respectively. Based on the surface area results, the amount
of catalyst used in the glass spheres experiments had
aproximately 5 times as much surface area as one gram

122
TABLE 9: Glass Spheres Results
Untreated
Spheres
FeCl2 * 3H20
Ni(N03)2•6H20
Weight (grams)
35
35
5
Surface Area
Mvgram
1.2
—
8.9
Total Surface
Area Mvgram
42
—
35
Conversion (%)
Air
1
9
0
Nitrogen
1
1
0
Space Time (hr)*
59
59
8
Temperature= 270°C
*Since the density of PPAN is approximately 1.8 g/cm3 while
that for the glass spheres is approximately 2.4 g/cm3, the
space time for the glass spheres should be multiplied by a
factor of 1.3 to compensate for the difference in density
between the PPAN and the glass spheres. Therefore, the
density corrected space times for the glass spheres
preparations would be 77.0 hr. for both the untreated
spheres and the iron chloride doped spheres while that for
the nickel nitrate doped spheres would be 10 hr. For
comparison, the space time of a typical PPAN experiment was
approximately 1.7 hr.

123
of PPAN, and in addition, the space time was greater using
35 g of glass spheres. Both of these factors should tend to
increase the conversion effeciency, however this was not
observed. (Although the numerical values of the space time
as presented have little physical meaning, the difference
between them is of significance in comparing the activity of
the various catalyst preparations.) In fact, with the
exception of the results for the undoped sample under
nitrogen, all of the results obtained demonstrate equal or
lower conversion effeciencies than for the corresponding
metal doped PPAN preparations. These results coupled with
the fact that the space time and surface area were
undoubtedly greater for the glass spheres experiments than
for the PPAN experiments suggest that PPAN is an
intrinsically more active catalyst or support than the glass
spheres used in this study. This conclusion is reinforced
by the results obtained using 5 grams of glass spheres
containing 3.45 x 10-4 moles of Ni(NO3)2’6H20. In this
case, the catalyst bed depth as well as the total surface
area should have been roughly equivalent to that obtained
using one gram of PPAN, however no catalytic products were
observed under air or nitrogen atmospheres. Another
striking difference between PPAN catalyst preparations and
the glass sphere preparations was that unlike PPAN
catalysts, the glass sphere catalysts did not produce any
detectable amounts of diethyl ether or ethyl acetate.

124
Oxide supported catalysts
As described previously, preparations were made which
contained PPAN on the surface of silica gel, titanium
dioxide, and aluminum oxide. As can be seem from the
surface area data in Table 2, the addition of PPAN to the
oxides resulted in a moderate decrease in the surface areas
for silica gel and aluminum oxide (titanium dioxide was not
measured), although the surface areas were still more than
an order of magnitude greater than for the unsupported
catalysts. Table 10 contains the elemental analyses for
these materials before and after the addition of PPAN, the
pyrolysis of the sample and catalysis, using ethanol as the
substrate (except titanium dioxide). In the case of
titanium dioxide, the starting material was white and after
the addition of PPAN the material became somewhat off white.
During the process of pyrolyzing the sample in nitrogen
(which employed the programmable temperature controller
using the same program as for alumina and silica), the
material gradually became brown and then proceeded to turn
white again. The elemental analysis results revealed that
nearly all of the carbon, nitrogen and hydrogen had
disappeared. Although the loss was not as significant for
silica and alumina, the proportion of the organic material
lost far exceeds any results for the pyrolysis of
unsupported PAN. In addition, during the catalytic

125
Table 10: Elemental Analysis Results For
Supported Catalysts
Sample
C%
H%
N%
Al2°3
0.03
0.24
0.0
a12°3*
0.64
0.21
0.0
ai2o3-ppan
3.55
0.52
0.0
ai2o3-ppan*
2.11
0.22
0.42
ai2o3-ppan**
1.29
0.16
0.15
Ti02
0.06
0.0
0.43
Ti02-PAN
9.67
1.11
3.47
Ti02-PPAN
0.05
0.08
0.08
sio2
0.02
0.32
0.0
Si02-PAN
6.62
1.00
2.47
Si02-PPAN
2.55
0.34
0.66
* After passing 5 ml of ethanol
** After passing 10 ml of ethanol

126
reaction, even more organic material was lost from the
alumina and silica catalysts (the titania supported material
was not tested for catalytic activity since there was no
PPAN left on the surface). These results seem to suggest
that these oxide materials are capable of somehow catalyzing
or promoting the decomposition and volatilization of the
polyacrylonitrile on the surface during both the pyrolysis
and the catalytic reaction, titania having the most
pronounced effect since all of the organic material was lost
during the pyrolysis.
The supported catalysts as well as the pure supports
were tested for activity towards ethanol and the results are
presented in Table 11. While the product distributions
obtained using the silica containing materials were
reasonably consistent with the results obtained for pure
PPAN materials (doped or undoped), the alumina containing
materials yielded a markedly different product distribution
in that very high conversion (30 to 50%) to diethyl ether
was observed. However, since the pure alumina support has a
very similar product distribution to the alumina supported
PPAN, the difference can probably be attributed to the
alumina. The addition of PPAN to these materials does not
result in a significant change in their activity towards
ethanol. In fact, for silica, the activity seems to be
somewhat less than for the unsupported catalysts. These
results are rather surprising for a number of reasons, the

127
Table 11: Catalytic Results For Supported Catalysts
Products
Acetaldehyde
(%)
Diethyl
(%)
Ether
Ethyl
(%
Acetate
)
Carrier Gas
air
n2
air
n2
air
n2
sio2
5.0
0.1
3.0
—
1.3
—
Si02-PPAN-1
6.0
0.1
0.6
0.3
1.1
—
Si02-PPAN-2
1.0
0.2
0.3
—
0.1
—
ai2o3-ppan
3.6
0.1
39.0
45.0
0.6
—
A1203
3.8
0.3
46.0
35.0
—
—
WHSW= 0.6
Temperature= 270°C

128
most important being that these materials have surface areas
which are an order of magnitude greater than the unsupported
catalysts, as shown in Table 2. The increased surface areas
would be expected to result in greater overall conversions,
although this proved not to be the case. The overall
conversion obtained using silica gel supported materials was
comparable to that obtained for unsupported catalysts, while
the alumina supported material had a greater conversion
efficiency; however, this was probably due to the alumina
and not the PPAN. Since Bai and coworkers20 never tested
the pure silica for catalytic activity, it is not possible
to elucidate the specific role played by the PPAN in the
reactions of cumene and ethylbenzene. Nonetheless, the
results from this laboratory lead one to suspect that the
catalytic contribution due to the support may be more
important than previously realized. It is possible that
their active catalyst was actually copper on silica gel with
the PPAN having little effect. This could also help account
for the discrepancy between the products formed in the vapor
phase using pure PPAN and those formed in the liquid phase
using silica supported PPAN. Of course one must remember
that the preparative method used by Bai and coworkers was
much different in that acrylonitrile was polymerized in the
presence of silica gel and pyrolyzed for a shorter period of
time at a lower temperature. It is quite likely that these
modifications could produce a completely different catalyst.

129
One puzzling feature of the catalytic reaction using
all types of PPAN based catalysts was that occasionally, for
no apparent reason, the products of the catalytic reaction
would be clear amber colored and minor impurities would
appear in the baseline of the gas chromatograms. The
discolored samples appeared randomly; for example, after
doing several catalytic runs which produced colorless
products, one run might produce colored products and then
the products of the next run would be colorless again. On
one occasion, while using PPAN supported on alumina, the
discolored products were analyzed by GC-MS and found to
contain a number of products containing four carbon
segments. Table 12 lists the possible identities of the
components of the product mixture. Although the reliability
of the GC-MS technique for the absolute identification of
unknown species is somewhat questionable, there are
certainly enough compounds with butyl groups to enable one
to safely assume that their presence (or possibly isobutyl,
butylene, isobutylene and butadiene) is not an artifact of
the instrumentation. The two logical sources of the butyl
groups would be either carbon-carbon bond formation from the
ethanol substrate or decomposition of the polymeric
catalyst. Since acrylonitrile is a four carbon monomer and
two ethanol molecules would combine to form a four carbon
species, it is difficult to determine the origin of these
products on the basis of carbon number. However, the

130
Table 12: GC-MS Results for Discolored Products
Retention Time (sec.)
Possible Identification
41
CH2CHCHCH2 and CH2CH0
43
ch3ch2oh
45
ch3ch2och2ch3
49
ch3ch2ch2oh
52
ch3ch2ch2cho
55
ch3co2ch2ch3
66
ch3ch2ch2ch2oh
75
ch3ch2ch2ch2och2ch3
86
ch2chch2ch2och2ch3
119
ch3ch2ch2co2ch2ch3
124
ch3co2ch2ch2ch2ch3

131
formation of the butyl groups from ethanol under these
conditions is not a trivial transformation. In addition,
none of the compounds listed in Table 11 are colored, so
that the amber color must be due to an intensely colored
minor impurity. One possible explanation is that during the
catalytic run the temperature overshot considerably and went
unnoticed. The increased temperature could result in the
further pyrolysis of the PPAN and it was not uncommon at all
for the contacts in a temperature controller to temporarily
stick together causing a temperature deviation or
occasionally even a melt-down of the entire glass oven. The
color could also be due to residues from the original
pyrolysis which had become trapped in the pores of the PPAN
(or alumina in this case) to be later washed out by the
ethanol substrate. The fact that the elemental analyses of
spent alumina supported catalysts showed a loss in organic
material (PPAN) during the catalytic reaction would support
either interpretation.

SUMMARY AND FUTURE RECOMMENDATIONS
Although a considerable amount of knowledge was gained
from the present study, there is still a great deal to be
learned about the structure of PPAN and how it relates to
the physical and catalytic properties of these materials.
The results of this and other studies show promise that a
thorough understanding of the structure-reactivity
relationships for PPAN based catalysts could yield a new
class of "tailor made" catalysts produced by molecular
engineering techniques. These materials would offer
advantages over many conventional industrial catalysts which
employ precious metal catalysts such as rhodium, platinum
and ruthenium, the continued supply of which is uncertain
since most of the worlds known reserves are located in South
Africa and the USSR.
It should be emphasized that no attempt was made to
optimize these catalyst preparations or the catalytic
reaction conditions used in these experiments. In spite of
this, the results were quite encouraging, particularly the
ability to control product selectivity by the addition of
small amounts of metal additives. Figure 30 is a graphical
representation of the extremes of product selectivity
132

133
PRODUCT
SELECTIVITY
F e ( N 0 )
3 3
Z r (acetate)
X
y
CrCI2
Figure 30: Selectivities of Metal Doped Catalysts

134
measured in this study, and with the proper optimization of
these catalysts (percent metal, pyrolysis conditions,
catalytic reaction conditions, etc.), it should be possible
to attain even better product selectivities. However, the
academic setting is not the proper place for this type of
endeavor, which is most properly carried out in industry.
It is also interesting to note that the product
selectivities of the zirconium basic acetate and zirconium
basic carbonate, as well as iron (III) nitrate and iron (II)
chloride, are guite different from each other, indicating
that the oxidation state of the metal as well as the
identity of the anion are important in catalysis. Another
significant result was the discovery that PPAN catalysts are
capable of decomposing compounds such as CEES and could
potentially be used for the demilitarization of the World
War II reserves of mustard gas into marketable chemicals as
a cost effective alternative to incineration.
A considerable amount of information was learned about
the pyrolysis reaction through the use of Differential
Scanning Calorimetry, Thermogravimetric Analysis and
elemental analysis results, and consequently, a number of
conclusions can be drawn. First of all, the amount of
weight loss observed and the elemental analysis results, as
well as the increase in brittleness and loss of surface
area, lead one to suspect that considerably more
crosslinking and chain scission are occurring during the

135
pyrolysis reaction than was previously believed, which is in
agreement with some recently published spectroscopic
results.8 These factors, coupled with the fact that the
polyacrylonitrile starting material is an atactic polymer,
serves to raise doubts as to what extent the proposed
cyclized structure shown in Figure 1 is actually present,
particularly since metal binding to PPAN could not be
unequivocally demonstrated. This again raises the question
as to what actually is the catalytic species, and without
further spectroscopic information, it is impossible to be
certain. The Thermogravimetric Analysis and Differential
Scanning Calorimetry indicated that both the heat loss and
weight loss occurred simultaneously, and it has been shown
that the amount of heat loss correlates with the
disappearance of the nitrile groups in previously reported
studies. However, the structure of the resulting product
has yet to be determined unambiguously. The catalytic
species also seems to be intrinsic to the polyacrylonitrile
since it was generated under several different pyrolysis
atmospheres in products having different elemental
compositions. In addition, it was interesting to learn that
the addition of oxide supports (which were found not to
appreciably affect the catalytic activity despite the
greatly increased surface area) served to facilitate the
decomposition of PPAN, titanium dioxide having the most
pronounced effect. Titania has been found to exhibit the

136
strong metal support interaction (SMSI), which is
characterized by a strong interaction between metal
catalysts and the support.62 Although PAN is quite
different from a metal, it is possible that some sort of
interaction exists between PAN and the titania support
resulting in the decomposition or cracking of the PAN on the
surface. Since the PAN was diluted to a great extent by the
support, the dissipation of the exothermic heat produced
during the pyrolysis was probably not a problem in the case
of the supported catalysts, making the large proportion of
organics lost even more significant.
Before PPAN based catalysts can become commercially
viable, there are a number of problems which will need to be
overcome. Perhaps the most important obstacle is the lack
of a good spectroscopic probe of the pyrolysis reaction or
the catalytic reaction due to the inherent insolubility and
paramagnetism of PPAN. The solution to this problem will
probably lie within the rapidly expanding realm of surface
analytical techniques such as X-ray Photoelectron
Spectroscopy, Multiphoton Ionization Spectroscopy, Secondary
Ion Mass Spectrometry, Laser Induced Fluorescence, and Auger
Spectroscopy, just to name a few. The information thus
obtained could enable researchers to determine the
relationship between the pyrolysis conditions, the structure
and the catalytic reactivity, as well to as elucidate the
role of the metal dopants in dictating product selectivity.

137
With this added knowledge it may be possible to better
understand and control the pyrolysis reaction (heat
dissipation, optimum temperature-time profile, etc.) in
order to maximize the concentration of the desired catalytic
species. However, significant commercial applications will
probably have to be found to motivate such an effort.
Other areas which will need study will be the use of
different types of polyacrylonitrile as well as
polycyanoacetylene for the starting material, in addition to
long term stability studies in order to understand and
prevent catalyst deactivation. Alternative types of PAN
which could be tried are various isotactic forms of PAN
prepared by such techniques as anionic polymerization,63
electropolymerization64-66 and possibly plasma
polymerization (there have been no reports of the
preparation of isotactic polymers using plasma
polymerization, although polymers with very high surface
areas have been prepared using this technique), since
isotactic PAN would be more likely to produce more fused
ring structures with less crosslinking than the atactic
polymers used in this study. This would be desirable if the
catalytic species turned out to be the product of Figure 1
as other investigators have proposed. The research just
outlined could certainly consume the graduate careers of a
number of graduate students, but if the results of this
study are any indication, it will be well worth the effort.

.-’387:38
APPENDIX
XPS SPECTRA
FILE NAME: 27434A
Stored on 1, 7,1986
LiCl IN ACRYLONITRILE
SAMPLE Q FG = 4mA-'4eV AEPR OPEN
see counts div
27434A
GROUPS 5 THRU S
CD
CD
Q
TT
in
- JC/v*
in
0
in
ivj
r-
a
I
u
CD
cn
OJ
CL
OJ
967.98 EV
1412 COUNTS
<3
PJ
,
id
PJ
CL
OJ
u
u
LfJ
li
n
IV-
j'j
CD
J1
u
\j
ca
â–¡
rxj
61.6
ca
â–¡
''w'^^L*r-uA
1GGG EV
138

975.95 O(KLL)
139
FILE NAME: 27434C
Stored on 1, 7, 1986
LiC1 + Cu IN ACRYLONITRILE - SAMPLE(jJ side b FG=4mñr4eV APER OPEN
588 COUNTS DIV
27434C
GROUPS 5 THRU 8
975.35 EV
1753 COUNTS
o
CD
in —
CO
rvj (J
1ÜÚÜ EV
$ EV

—cjwia
140
FILE NAME: 27434B
Stored on 1, 7,1986
Li Cl + Cu IN ACRYLONITRILE - SAMPLE(?) side a FG=4mA'4eV APER OPEN
588 COUNTS DIV
274348
GROUPS 5 THRU 0
112.22 EV
498 COUNTS
u
O
IT)
rs.
tr
uo
<\J ~
03 M
tn w
m
*>1
O
W*'
Q-
Cl
—
\J
CM
w
0)
OJ
a
cu
w
w
J
u
CO
00
DO
©
ID
07
(VI
X>
cn
OJ
T>
cn
Q
pj
cu
(VI
LO
-
M — —
1
1889 EV
8 EV

827434 LiCI COMPOUNDS IN ACRYLONITRILE
141
Memory group
Subtract 310
0 thru 2
counts
3 Point Smooth
10S9 COUNTS DIV
27434D
REGION 13
GROUPS 2 THRU 9
289.28 EV
1333 COUNTS
Peak
E. V.
= 286.88
Height *
6507
Area = 130205
Peak
E. V.
= 286.72
Height *
6464
firea = 129792
Peak
E. V.
= 285. 12
Height **
5289
Area = 1 10403

B£743A L1C1 COMPOUNDS IN ACRYLONITRILE
142
Memory group 3 thru 5
Subtract 1537 counts
5 Point Smooth
5SS COUNTS DIV
27434D
REGION 14
GROUPS 5 THRU 3
532.52 EV
4254 COUNTS
Peak
E. V.
= 532.52
Height *
4225
Area =
66323
Peak
E. V.
= 532.36
He i ght *
4625
Area =
66445
Peak
E. V.
= 532.36
Height «>
3453
Area =
56742
532.3G

827^ LiCI COMPOUNDS IN ACRYLONITRILE
143
Memory group 6 thru 8
Subtract 2372 counts
7 Point Smooth
1888 COUNTS DIV
27434D
REGION 15
GROUPS 8 THRU 6
488.24 EV
737S COUNTS
418 EV
338 EV
Peak
E. V.
= 400.24
Height ”*
7281
Area =
78876
Peak
E. V.
= 399.92
Height
7067
Area =
78357
Peak
E. V.
= 399.92
Height ~
3949
Area =
47818

B27-134 l'iCI ¿OMPollNDS 'IN R¿RYLC^JITRíLE
144
Memory group 0 thru
Subtract 554 counts
1
Memory group 2 thru 2
Subtract 1540 counts
9 Point Smooth
209 COUNTS DIV
27434D
REGION IS
GROUPS 2 THRU 9
299.52 EV
112 COUNTS
219 EV
199 EV
Peak
E. V.
= 198.64
Height ~
1063
Area =
22553
Peak
E. V.
= 198.64
Height ■»
949
Area =
20427
Peak
E. V.
= 198.80
Height ~
1448
Area =
26954

145
Memory group 3 thru 3
Subtract 5758 counts
Memory group 4 thru 4
Subtract 6958 counts
Memory group 5 thru 5
Subtract 8107 counts
9 Point Smooth
2SS COUNTS DIV
27434D
REGION 17
GROUPS 5 THRU 3
935.24 EV
1463 COUNTS
945 EV
Peak E.V. = 935.40
Peak E.V. = 935.24
Height » 1464
Height "> 901
Area = 21404
Area = 12324
925 EV
B27434 LiCl -tCu COMPOUNDS IN ñCRYLONIT

146
Memory group 6 thru 8
Subtract 679 counts
9 Point Smoot h
58 COUNTS DIV
27434D
REGION 18
GROUPS 3 THRU 6
56.28 EV
272 COUNTS
65 EV
45 EV
Peak
E. V. =
56.20
He i ght "•
271
Area =
1075
Peak
E. V. =
55.88
Height «*
238
firea =
1830
Peak
E. V. =
55.40
Height *
167
Area =
1059
B27434 Li Cl C0I1P0UMD5 IN RCR VLON1TRILE

Cu Id POLVflCPYLONITR[LE
147
Memory group 6 thru 6
Subtract 10319 counts
Memory group 7 thru 7
Subtract 23757 counts
Memory group 8 thru 8
Subtract 9113 counts
5 Point Smooth
2886 COUNTS DIV
27434E
REGION 17
GROUPS 8 THRU S
833.64 EV
1486 COUNTS
845 EV 825 EV
Peak
E. V.
= 935.24
Height *
2877
Area =
41679
Peak
E. V.
= 935.88
Height ■»
14370
Area = 248296
Peak
E. V.
= 935.08
Height *>
4137
Area =
71062

148
588 COUNTS DIV
27434E
REGION 19
GROUPS 8 THRU 3
954.58 EV
3276 COUNTS
1S8S EV
8 EV
B27434 LiCI+Cu IN POLVflCRYLONII

OKLL
149
FILE N AME : 27434F
Stored on 1, 9,1986
LiCl + Cu IN POLYACRYLONITRILE - SAMPLE (?) FG=4mlV4eV APER OPEN
1888 COUNTS DIV
27434F
GROUPS 5 THRU 8
977.29 EV
7318 COUNTS
1880 EV
8 EV

150
FILE NAME: 27434G
Stored on 1, 9,1986
L1C1 + Cu IN POLYACRYLONITRILE - SAMPLE 3 FG=4mA/4eV APER OPEN
580 COUNTS DIV
27434G
GROUPS 5 THRU 8
25.38 EV
185 COUNTS
1888 EV
8 EV

151
Memory group
Subtract 456
0 thru 2
c ount s
3 Point Smooth
1888 COUNTS DIV
27434E
REGION 13
GROUPS 2 THRU 0
289.68 EV
11 S3 COUNTS
388 EV
288 EV
Peak
E. V.
= 285.44
Height •>
5530
Area = 119833
Peak
E. V.
= 285.44
Height «■
6594
Area = 94535
Peak
E. V.
= 286.56
He i ght *>
6103
Area - 1 13363
2B6.56

Cu IN POLYACRYLONITRILE
152
5 Point Smoot h
1QS8 COUNTS DIV
27434E
REGION 15
GROUPS 2 THRU 8
482.16 EV
1283 COUNTS
418 EV
3S8 EV
Peak
E. V.
= 400.24
Height nj
4735
Area =
43341
Peak
E. V.
= 400.56
Height
1371
Area =
11194
Peak
E. V.
= 399.76
Height "»
6767
Area =
73002

B27434 LiCl + Cu IN POL YRC4?YL0NITR1LE
153
Memory group 3 thru 3
Subtract 1523 counts
5 Point Smooth
1880 COUNTS DIV
27434E
REGION 14
GROUPS 5 THRU 3
532.68 EV
6167 COUNTS
545 EV
525 EV
Peak
E. V.
= 532.68
Height ">
6112
Area =
89184
Peak
E. V.
= 532.52
Height *
8142
Area -
124818
Peak
E. V.
= 532.52
He i ght '»
2817
Area =
50229

154
Memory group 0 thru 2
Subtract 936 counts
Memory group 1 thru 1
Subtract 554 counts
9 Point Smooth
29 COUNTS DIV
27434E
REGION 18
GROUPS 2 THRU 9
55.88 EV
155 COUNTS
65 EV
45 EV
Peak
E.V. =
55.88
He i ght *>
154
Area =
1197
Peak
E.V. =
56.04
Height «*
192
Area =
931
Peak
E.V. =
55.56
Height ■»
119
Area =
226

B2?434 LiCl +â–  Cu IN POLYRCRYLONITRILE
155
Memory group 3 thru 5
Subtract 1660 counts
Memory group 4 thru 4
Subtract 2313 counts
7 Point Smooth
5S8 COUNTS DIV
27434E
REGION IS
GROUPS 5 THRU 3
2S6.88 EV
247 COUNTS
218 EV 138 EV
Peak
E. V.
= 198.80
Height *»
1830
fires =
33548
Peak
E. V.
= 199.12
Height **
4611
Area =
83214
Peak
E. V.
= 198.48
Height <»
1515
Area =
31118

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BIOGRAPHICAL SKETCH
Jeffrey L. Clark was born on April 1, 1954, in Newark,
New Jersey. He graduated in 1972 from the Chapel Hill
Senior High School, Chapel Hill, North Carolina. He
received a BS with honors in chemistry in 1981 from Ithaca
College, Ithaca, New York, and has been attending The
University of Florida since 1982. From August to December
in 1986, he was employed by Geo-Centers, Inc., Newton Lower
Falls, Mass.
161

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
Russell S. Drago, Chairman
Graduate Research 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.
David E. Richardson
Assistant 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.
George E. Ryschkewitsch
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.
John G. Dorsey
Assistant Professor of
mistry

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 deqree of Doctor of Philosophy.
This dissertation was submitted to the Graduate Faculty of
the Department of Chemistry in the College of Liberal Arts
and Sciences and to the Graduate School and was accepted as
partial fulfillment of the requirements for the degree of
Doctor of Philosophy.
December 1987
Dean, Graduate School

UNIVERSITY OF FLORIDA
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