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LIBRARY
Michigan State
University

 

 

ABSTRACT
SPECTROSCOPIC STUDIES OF

ALKALI AND AMMONIUM ION SOLVATION
IN PROPYLENE CARBONATE

By
Diana M. Wied

The subject of this thesis is the study of alkali and
ammonium ion solvation in propylene carbonate by infrared
spectroscopy.

In propylene carbonate, solutions of the various cations
generally show anion independent far-infrared bands which have
been ascribed to the vibration of the cation in a solvent cage.

Isotopic substitutions of the lithium cation produced
frequency shifts in the expected direction, thus verifying the
assumption that the cation is involved in the observed vibration.
Not all band positions were anion independent, especially for the
lithium salts used. Evidence of contact and solvent separated
ion pairing was shown from the anion dependence of the solvation
band frequencies. In an attempt to determine the species present
in solution, Beer's law plots were made for two lithium and two
sodium salts. The results showed that for propylene carbonate,
the study of Beer's law plots is probably not a sensitive enough

method for use in determination of the species present.

Diana M. Nied

Observation of the spectrum of propylene carbonate in the
infrared region showed that the position of the carbonyl stretch
is shifted to lower frequency by the addition of alkali metal or
ammonium ion salt. The carbonyl oxygen is, therefore, the solvation

site on the solvent molecule.

SPECTROSCOPIC STUDIES OF
ALKALI AND AMMONIUM ION SOLVATION
IN PROPYLENE CARBONATE

By

q
Diana M9‘Nied

A THESIS

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

MASTER OF SCIENCE
Department of Chemistry

T971

ACKNOWLEDGMENTS

The author wishes to express deep gratitude to Dr. A. I. Popov,
without whose patience and understanding this work could not have
been possible.

Special thanks should be given to Paul Handy for help with
operating the far ir and to Dick Bodner for help with writing,
proofreading and for his moral support while the author was looking
for a job.

Thanks is due to the National Science Foundation for the

financial support for this work.

11

TABLE OF CONTENTS

ACKNOWLEDGMENTS ......................
LIST OF TABLES ......................
LIST OF FIGURES ......................
Chapter
I. INTRODUCTION ....................
II. HISTORICAL .....................
A. Far-Infrared Studies of Ionic Solvation .
8. Properties of Propylene Carbonate .......
C. Physicochemical Studies in Propylene Carbonate .
III. EXPERIMENTAL ....................
A. Purification of Reagents ............
l. Propylene carbonate ............
2. Nitromethane ...............
3. Salts ...................
B. Preparation of Solutions ............
C. Far-Infrared Measurements ............
I. General techniques ..........
2. Determination of solvation band positions. .
3. Beer' 5 law studies .............
D. Near Infrared Measurements ...........
E Raman Measurements ...............
IV. RESULTS AND DISCUSSION ...............
A. Far-Infrared Studies ..............
B.
l. General ...................
2. Peak positions ...............
3. Beer's law studies .............

iii

Page
ii

vi

TO

16
16
27

V.

B.
C.

Infrared Measurements . .......
Raman Mole Ratio Study ........

SUGGESTIONS FOR FURTHER STUDY .....

BIBLIOGRAPHY .........

iv

Table

LIST OF TABLES

Important Physical Properties of Propylene Carbonate . .

Ion Pair Formation Constants in Propylene Carbonate

Positions of the Solvation Bands in Propylene

Carbonate ......................

Solvation Band Area vs Concentration for Lithium

Perchlorate in PropyTEhe Carbonate ..........

Solvation Band Area !§_Concentration for Lithium
Bromide in Propylene Carbonate ..........

Solvation Band Area vs Concentration for Sodium

Perchlorate in PropyTEhe Carbonate ..........

Solvation Band Area vs Concentration for Sodium

Iodide in Propylene Carbonate. . . . . . .......

Page

24

28

. 28

. 29

LIST OF FIGURES

Figure Page

l. Absorption spectrum of pure propylene carbonate from
600 to 100 cm' .................. . . . l8

2. Solvation band absorption spectra in propylene
carbonate ............... . . . . . . . . 20

3. Solvation band spectra in propylene carbonate with
solvent subtracted .............. . . . . . 23

4. Beer's law plot for lithium perchlorate in propylene
carbonate ..................... . . 32

5. Beer's law plot for lithium bromide in propylene
carbonate .................. . . . . . 34

6. Beer's law plot for sodium perchlorate in propylene
carbonate ............... . . . ..... 36

7. Beer's law plot for sodium iodide in propylene
carbonate ............ . . . ...... . . 38

vi

I. INTRODUCTION

Since the vast majority of analytical chemical reactions occur
in solutions and involve ionic species, the importance of the
structure of these solutions and the nature of the species present
in them needs no emphasis. It should be noted, however, that our
knowledge of either the structure or of the species present is very
imperfect. While it is known that ions in solutions are solvated,
the extent of solvation is usually not known. The overall extent
of ionic association can be obtained from electrical conductance
measurements. However, these measurements give no information on
the nature of the associated species, and cannot distinguish between
contact ion pairs, solvent shared ion pairs or solvent separated ion
pairs.

Recently it has been found that far-infrared measurements
on ionic solutions in various solvents can be useful in the elucidation
of the nature of the solute species. Such studies have been carried
out in this laboratory on several solvent systems, and the present
thesis is a part of this project.

The solvents used for this study have to meet two requirements.
They have to be relatively transparent in the far ir from 600 to
80 cm-], and should be capable of dissolving a wide variety of salts.
Preliminary studies indicate that propylene carbonate meets these

requirements.

Alkali metal and ammonium salts were used in this investigation
because their cations have a single charge and because they do not

readily form complexes with non-aqueous solvents or with the anions.

II. HISTORICAL

A. Far-Infrared Studies of Ionic Solvation

 

The pioneering far-infrared studies of ionic solutions were
done by Evans and Lo (l). They observed bands in the far ir for the
tetrabutyl- and tetrapentylammonium chlorides and bromides in
benzene solutions at I20 and 80 cm'] respectively. These bands
were ascribed to the vibration of the respective ion pairs. Low
solubilities prevented a study of the position of the bands as a
function of concentration.

Edgell and co-workers (2,3) observed far ir bands which
could not be ascribed either to the solvent or to the solute, when
various alkali metal salts were dissolved in tetrahydrofuran (THF)
and in dimethylsulfoxide (DMSO). In DMSO the band positions were

I for lithium and

anion independent, appearing at 424 and 200 cm-
sodium salts respectively. In THF, however, the band frequencies
were dependent on the anion. This dependence was attributed to the
possible formation of contact ion pairs in that solvent.

Maxey (4,5) studied the solvation of alkali metal and ammonium
ion salts in dialkylsulfoxides by far ir spectroscopy. The positions
of the bands were anion independent. In DMSO band maxima for Li+,

+, Rb+, and Cs+ were observed at 429, 214, 199, 153,

+ +
NH4 , Na , K
124, and llO cm'1 respectively. The other solvents used showed

similar behavior.

Following the above work, Wuepper (6,7), McKinney (8),

Wong (8,9), and presently Handy have extended the range of solvents
to include 2-pyrrolidone, l-methyl-Z-pyrrolidone, l-vinyl-Z-
pyrrolidone, acetone, acetic acid, pyridine, and various picolines.
Because the bands could not be assigned either to the solute or to
the solvent vibration, they had to be due to a solute-solvent
interaction. Also, since their frequencies, in most cases, depended
only on the nature of the cation, the band was ascribed to the
vibration of the cation in a solvent cage. To verify this model,
isotopic substitutions on the cations and on the solvents were made.
Frequency shifts of these bands in the anticipated directions
supported the assignments. Anion dependence was ascribed to the
presence of contact ion pairs in solution.

In solvents of low and medium dielectric constant, it is
generally assumed that as the concentration of salt in solution is
increased, contact ion pairs are more likely to predominate. Likewise,
solvent-separated ion pairs are assumed to be the primary species
in dilute solution.

While it has been noticed that, in general, the intensities
of the solvation bands have been proportional to salt concentration,
no accurate measurements of intensity v§_concentrations have been
made. Edgell §t_al, (3) attributed the anion dependence of the band
positions in THF to contact ion pairing. In an effort to determine
the type of ion pair in solution, the integrated absorbance was
studied as a function of salt concentration for sodium tetracarbonyl-

cobaltate in THF. In any given THF solution, there could be an

equilibrium between the different types of ion pairs. If any one,
or several, of these species is giving rise to the band, and if
the position of the equilibrium changes with concentration, then a
Beer's law plot of the band intensity should show a deviation from
linearity. Studies of these plots could yield useful information
regarding the species in solution.

The result of the intensity y§_concentration study was a
straight line. The authors concluded that either the absorption
was due to a single species or that the ratio of various forms of
ion pairs did not change. However, this experiment had three major
flaws. One was that they used tetracarbonylcobaltate, which is a
large tetrahedral anion and not very likely to form ion pairs. Two,
sodium was used as the cation, and it in general does not interact
as strongly as does, say, lithium, due to sodium‘s smaller charge
to radius ratio. Finally, the concentration range used was only
from O.l to 0.3 M, probably not large enough to significantly shift
the equilibrium. Studies with a more judicious selection of solutes
and over a larger concentration range could prove to be more

clarifying.

B. Properties of Propylene Carbonate

 

Initial studies showed propylene carbonate to be an
appropriate solvent for this work, as it has a useful transparency
in the far ir and a reasonably good dissolving power.

Propylene carbonate (PC) is a colorless, moderately viscous

liquid with very low toxicity and no appreciable odor. It has been

classified as an aprotic dipolar solvent with no internal solvent
structure (l0). Because it is a racemic mixture of forms I and

II, PC has a very large liquid range of -48.9 to 242°C (ll). It

C) C)

*0 0X0
(1) L‘é CH3 (II) CH3

tends to supercool easily and can be kept at -80°C for several

months as a liquid (12). It has a high dielectric constant of

65.0 at 25°C (ll), and its donor number, which is a measure of its
solvating ability, is l5.l. The donor number of a solvent is defined
by Gutmann (l3) as the heat released, in kcal/mole, in the following

reaction, where S is the solvent in question. The reaction is

S + SbCT5 + S°SbCl5

carried out in l,2-dichloroethane at 25°C. The value of 15.l is
intermediate in the range of donor numbers found by Gutmann and
indicates that PC is a solvent of moderate solvating ability.
Conductance studies by Courtot-Coupez (l4) agree with the above
conclusion.

The list of important physical properties of PC appears in
Table l.

Propylene carbonate has only recently come to be considered

a useful non-aqueous solvent and was called by Friedman (l7) an ideal

Table l. Important Physical Properties of Propylene Carbonate.

 

 

 

Property Value Reference
Melting point -48.9°C ll
Boiling point 242°C ll
Density 25°C l.199 i 0.00l g/cc ll
Refractive index 20°C l.42l4 i 0.0004 ll
Dielectric constant 25°C 65.0 i 0.3 ll
Dipole moment 4.94 0 l5
Viscosity 2.530 cp l6

 

dipolar aprotic structureless medium. Before long it should

become one of the more important non-aqueous solvents.

C. Physicochemical Studies in Prgpylene Carbonate

 

Physicochemical studies in PC began to appear a few years
ago and since then their number has greatly increased.

Friedman and co-workers (16,18-20) have studied the heats
of transfer of a large number of alkali metal and alkyl ammonium
salts from water to PC by measuring the heats of solution of these
salts in both solvents. Later they began to study the heats of
transfer from DMSO to PC.

Butler and co-workers (21,22) studied the solubility and
complex formation of silver chloride in PC and in PC-water mixtures.
They also looked at the solvation of various ions by small amounts
of water present in PC (10). Association constants of these ions
with the water were determined.

Salomon (23-26) studied the thermodynamics of lithium chloride
and bromide as well as of lithium, sodium, and potassium iodide
solutions in PC. In his studies he determined the activity coefficient,
standard free energies, enthalpies, and did some electrochemical
measurements on these solutes. He calculated the ion pair
association constants from the ratio of the measured mean activity
coefficients to those calculated from the modified Debye-Huckel
equation, assuming complete dissociation. These data are presented
in Table 2. Lithium iodide, as mentioned below, has proved to be
very difficult to dry, and it is very likely that Salomon's lithium

iodide may have been wet, as it was dried for only 24 hr at 250°C.

Table 2. Ion Pair Formation Constants in Propylene Carbonate.

 

 

Salt K

 

f
LiCl 50 : lO
LiBr 2.5 t 1
Lil 0.0 i l
NaI 0.55 1 0.1
KI 0.0 i l

 

Mukherjee and Boden (27) studied the conductances and
viscosities of lithium and quaternary ammonium salts in PC. They
found evidence for contact ion pairs in lithium chloride and
bromide solutions but not for those of the perchlorate. Association
constants of 557 and 19 for the chloride and the bromide ion pairs
respectively were reported. Also, from ionic mobilities they
concluded that in PC the size of the solvated lithium cation is
larger than that of the tetrabutylammonium ion. The effective
sizes of the perchlorate, iodide, bromide, and chloride ions were
found to be the same (28).

Choux and Benoit (29) also studied solvation of various ions
in PC by measuring their enthalpies of transfer from water to PC.
McComsey (30) made some preliminary investigations on polarography
in PC. Jasinski and Carroll (31) developed a spectrophotometric

method of determining water in PC with reasonable accuracy down to 20 ppm.

III. EXPERIMENTAL

A. Purification of Reagents

 

l. Propylene carbonate. Propylene carbonate (PC) was

 

obtained from Aldrich, dried over molecular sieves (Linde 4A), and
distilled from fresh molecular sieves at 150°C and 40 torr. The
middle fraction was retained. Comparison of its near ir spectrum
with the standard PC spectrum showed no extraneous peaks. The water
content as determined by Karl-Fischer titration was 0.002 M, The
purified PC was stored over molecular sieves. No increase in water

content was observed after a storage of up to three months.

2. Nitromethane. Three liters of nitromethane (Matheson,

 

Coleman and Bell) was successively washed with one liter of saturated
sodium bicarbonate, one liter of half-saturated sodium bisulfite,

two one-liter portions of distilled water, one liter of five

percent sulfuric acid, one liter of half-saturated sodium bicarbonate
solution, and one liter of water. Excess water was removed by

drying the nitromethane with two three-hundred gram portions of
calcium chloride, for a few minutes each, and then over molecular
sieves overnight. The product was distilled over barium oxide at
99.9°C and comparison of its neat pmr spectrum with the standard
nitromethane pmr spectrum showed no impurities. Karl-Fischer
titration showed the purified nitromethane to be 0.0033 M_in water.

10

11

It was stored in a brown bottle over molecular sieves with the top

sealed with cellophane.

3. Spits, Lithium6 nitrate, bromide, and perchlorate were
synthesized by McKinney (32) and were dried at 200°C before use.

Lithium tetraphenylborate was synthesized by Wong (33) and
used without further purification.

Special care was taken in the drying of ammonium iodide and
perchlorate, sodium tetraphenylborate, and ammonium, sodium and
potassium thiocyanate as they are generally quite hygroscopic and
decompose at high drying temperatures. These salts were dried in
a vacuum oven for four days while the temperature was gradually
raised from 50 to 60°C. They were stored in an evacuated desiccator
charged with phosphorus pentoxide.

All other salts used were dried by oven heating for two days
at 170°C. Since the salts would occasionally fuse, they were
periodically removed from the oven, ground in a mortar, and replaced
in the oven to continue drying. These salts were all kept in
evacuated desiccators charged with Aquasorb (Mallincrodt). The
desiccators were periodically re-evacuated.

Water titrations done on the salts showed that lithium salts
could be kept for several months at a constant 0.2 weight percent
of water, an acceptable dryness for far-infrared studies. All other
salts could be dried to about 0.02% water and kept that way for

several months.

12

B. Preparation of Solutions

 

Since most of the salts used in this study are hygroscopic,
care was used in the preparation of their PC solutions. Transfer
of the salt from the desiccator to the volumetric flasks was done
as rapidly as possible. Solutions could be kept dry indefinitely
by sealing the top of the volumetric flask with cellophane. In
some mole ratio studies, the solutions were further dried by the

addition of molecular sieves.

C. Far-Infrared Measurements

 

1. General techniques. A few far ir measurements were

 

obtained on a Perkin Elmer Model 457 Grating Infrared Spectrophoto-
meter operating in the region from 600 to 250 cm']. It was run on
the slowest scan rate at an expansion of 10 giving a resolution of
about 3 cm']. The greater majority of far ir measurements were
obtained on a Digilab FTS l6 Fourier transform spectrophotometer.
The theory and operation of this instrument have been recently
described (34). It was operated in the single beam mode at a
nominal resolution of 4 cm']. The computer has storage space in
memory for the data for one spectrum. This spectrum is then
automatically subtracted from all sample spectra plotted in either
absorbance or transmittance scale and will remain the reference
spectrum until another is stored in its place.

A standard demountable liquid cell with polyethylene windows
and a 0.1 mm path length was used. It was found that this path

length minimized the noise and maximized the usable range of salt

13

concentrations. In order to maintain a constant path length, the
cell was not dismantled between runs. It was cleaned by rinsing

with dried acetone and vacuum drying.

2. Determination of solvation band positions. Spectra

 

consisting of just the solvation band could be obtained by employing
pure solvent as a reference. Usually all of the solvent peaks
would subtract well. However, for lithium salts some solvent bands
remained, and for lithium and lithium6 nitrate, the low solubility
and the poorly subtracted solvent band at 445 cm"1 interfered with
the accurate determination of their solvation band frequencies.

The minimum concentration of lithium salt which would produce
a distinct solvation band is approximately 0.25 M, Whenever
esolubilities allowed, several concentrations of each salt were
used, to determine if the position of the solvation band was
dependent on salt concentration. Since the molar absorptivity of
the cation-PC vibration becomes smaller as the atomic weight of the
cation increases, the usable concentration range diminishes with
increasing cation size. For rubidium and cesium thiocyanate,
saturated solutions at about 0.6 and 0.7 M respectively produced
barely visible bands.

Once a plot of a spectrum was obtained, the center portion
of the band was expanded and replotted to obtain a more accurate

measurement of the band position.

3. Beer's law studies. A Study of the integrated area of

 

the solvation band as a function of salt concentration was done for

lithium perchlorate, lithium bromide, sodium perchlorate, and

14

sodium iodide solutions in PC. Salt concentrations were varied

from 0.1 to 1.0 M, To do this study, accurate area measurements

had to be made on the solvation bands. Since the spectra of the
solvation bands with the solvent subtracted were too noisy and since
the baseline could not be determined with reasonable accuracy, a
different method of obtaining the spectra was developed. The
spectrum of three polyethylene plates in a mull cell was stored,

and then pure PC and various PC solutions were run in the demountable
cell. When all spectra were obtained, the baseline from the

solvent spectrum was traced onto each solution spectrum, and the
area of the band was found with a planimeter.

Whenever solute is dissolved in a solvent, the solvent can
be considered to be “diluted” by the solute. Therefore, in a
solution the far ir absorption spectrum of the solvent will be less
intense, making the solvation band appear smaller than it is. A
correction was applied for the sodium iodide and perchlorate Beer's
law studies by ratioing the area of a pure solvent peak at 330 cm-1
with the area of the same peak in the sample spectrum.

For lithium perchlorate and bromide solutions, however, this
correction method could not be used, since the only solvent
vibration which could be used for ratioing the areas become more
intense as the salt concentration increased. Therefore, for both

lithium salts, the unratioed areas were used.

 

0. Near Infrared Measurements

All near infrared measurements were made on a Perkin-Elmer

237B grating infrared spectrophotometer Operating between 4000 and

15

625 cm-]. Sodium chloride windows were used with a thin film of
sample. A thin sheet of polyethylene was used to calibrate the

instrument.

E. Raman Measurements

 

All Raman measurements were obtained on a laser Raman
spectrophotometer constructed at this university (35). The 5145 A
exciting line of an argon ion laser was employed. The instrument
was operated at a resolution of 1.2 cm'].

A mole ratio study was done observing the position of the
0] Raman active perchlorate band (36) as a function of the mole
ratio (R) of propylene carbonate to lithium perchlorate, where R was
varied from 20 to 1. Nitromethane was used as the inert solvent.

There are three major sources of difficulty in this experiment.
One is that the minimum concentration of perchlorate which will
produce a visible Raman band is 0.5 M, Two, lithium perchlorate is
only soluble in nitromethane to about 0.1 M, Finally, nitromethane

1 which could interfere with the

has a strong vibration at 917 cm-
determination of the perchlorate band position. The first two
difficulties were overcome by the fact that, although a salt may be
only slightly soluble in nitromethane, the addition of even a

small amount of PC (as when R = I) greatly enhances the solubility
of the salt in the mixed solvent. The third difficulty was overcome
by adding just enough nitromethane to dissolve the lithium

perchlorate. Molecular sieves were added to the flasks to keep the

solutions dry.

IV. RESULTS AND DISCUSSION

A. Far-Infrared Studies

 

1. General. Preliminary investigations showed propylene
carbonate to be a suitable solvent for far—infrared solvation studies.
In the 600 to 50 cm'1 region, the pure solvent has bands at 540, 445,
330, 200 (very weak), and 85 cm']. Unfortunately, the assignments
have not been made on the infrared vibrations of PC. However,
since spectral windows are of greater importance and since the
existing solvent bands, with the exception of the band at 85 cm'],
are very weak, knowledge of the vibrational assignments is not
necessary for our purposes. Figure 1 shows the absorption spectrum

of PC from 600 to 100 cm'].

2. Peak positions. When a lithium salt is dissolved in PC,

 

a new band appears at m 400 cm-]. This is the solvation band which
has previously been shown (4-9) to be due to a solvent-cation
interaction. Likewise, sodium salts dissolved in PC cause a band

at m 185 cm'] to appear. This is illustrated in Figure 2. The
dashed line is the spectrum of 0.68 M_lithium bromide and the dotted
line the spectrum of 1.48 M_sodium perchlorate in PC. Where the
dashed and dotted lines end, the solvent spectrum (solid line)
continues as normal. The two peaks are close in size, even though
the sodium ion concentration is twice as large as that of the lithium

ion. The decrease in intensity per unit of concentration with

16

17

Figure l

Absorbance spectrum of pure propylene carbonate from 600 to 100 cm'1.

18

 

 

 

 

ABSORBANCE

100

200

300

400

500

600

cm

19

Figure 2

Salvation band absorption spectra in propylene carbonate.

Pure solvent

 

............... Solvation band of 0.68 M LiBr

°°°°°°°°°°°°° Solvation band of 1.48 MNaClO4

 

20

 

Mannie.“

 

 

ABSORBANCE

100

200

300

400

500

600

cm

21

increasing ionic radius has been previously observed. Band
intensities analogous to those found by Wuepper (6) in 2-pyrrolidones
were observed in PC. Figure 3 illustrates the appearance of the
solvation bands when the solvent spectrum is subtracted.

When various alkali and ammonium ion salts were dissolved in

PC, bands appeared at approximately 400, 185, 185, 142, 115, and

l +

, Na+, K

112 cm' for Li+, NH4 , Rb+, and Cs+ salts respectively.
Unlike previous work, however, the positions of these bands are not
completely anion independent. Table 3 shows the Solvation band
positions of the salts used.

Isotopic substitutions of the lithium ion for the bromide,
perchlorate and the nitrate salt produced shifts of 25, 16, and
30 cm'] respectively in the expected direction. These shifts verify
the assumption that the cation is involved in the vibration, but
give no evidence as to the stoichiometry of the vibrating species.
The band positions were independent of the salt concentration. The
addition of a small amount of water caused the solvation bands to
disappear. This is expected in view of the fact that the donor
number of water, as determined by Erlich (37) is almost twice that
of PC and water should preferentially solvate the cations.

Salomon (26) and Mukherjee and Boden (27) have reported that
the contact ion pair formation constants for lithium bromide in PC
to be 2.5 i 1 and 19 respectively. For lithium perchlorate solutions,
the latter authors reported no evidence of association in PC. This

would mean that from 70 to 95% of the lithium bromide in solution

exists as contact ion pairs. Therefore, the solvation band at

22

Figure 3

Solvation band spectra in propylene carbonate with solvent subtracted.

 

Solvation peak of 0.92 M LiBr

----------------- Solvation peak of 1.68 M_NaC1O4

23

 

 

r
:“““’
’51-:
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0"-..
F-l
-..”)
tan"
ice-=5.
a;
e...__
““t.
~ -
C. \ -\
” s
\
\\
\

 

 

ABSORBANCE

100

200

300

400

500

600

cm

meeepsoeme

 

24

 

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e e em, 4 h Nam -mee
empm eom_ e h emm
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4 e om_ e h mm_ 4 H Rom e h m_e -eo_u
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e e mwm e H o_e -em
+mu +em +¥ +ez +412 +eb +me4 eoee<

 

 

AFIEU CC muwconcfimu mam—XQOLQ cw mvcmm cowum>~ow 2.3. $0 mcowpwmoa .m mrnwh

25

385 cm-1 for lithium bromide could be due to the vibration of the
lithium ion in a solvent cage in which the bromide ion has replaced
a solvent molecule. Likewise, the 397 cm'1 band for lithium
perchlorate would be due to the cation vibrating in a cage of
solvent molecules alone. The fact that the band for lithium
tetraphenylborate is also at 397 cm.1 verifies the latter conclusion,
for the tetraphenylborate anion is a very bulky symmetrical ion
not likely to form contact ion pairs. Within experimental error,
the band for lithium nitrate fits into this.category.

Solutions of thiocyanate salts gave interesting results.
Lithium thiocyanate solutions showed two new bands while the
ammonium and sodium salt both produced a band and a distinct
shoulder. With the latter two salts, there were no poorly subtracted
solvent bands in the 200 cm'] region with which these shoulders
could be confused. For lithium thiocyanate solutions at low

1 solvent band interfered with the

concentrations, the 445 cm-
determination of the solvation band positions. Although two
solvation bands for one salt have not been observed before, this
result is not unexpected. Nevertheless, unambiguous assignment of
the bands to specific species, as explained below, is not possible
as yet.

A possible interpretation of the data is that there exists
a type of "linkage isomerism" with the thiocyanate anion.
Thiocyanate ions are known to coordinate either through the sulfur

atom or through the nitrogen atom depending upon the nature of the

acceptor. It is possible that the bands produced here are due to

26

the cation vibrating in a solvent cage with one thiocyanate ion
replacing a solvent molecule, where either end of the thiocyanate
ion is oriented toward the cation. This could produce two bands
if the energy of interaction of the cation with the two ends of
the thiocyanate ion is different. However, in the spectra of all
six thiocyanate salts used, a small band appeared between 468 and
478 cm']. This peak indicates (38,39) that the anion coordinates
through the nitrogen atom, thus eliminating any linkage isomerism.

An alternate explanation is that contact and solvent
separated ion pairs could be present in solution, and that the two
bands are due to the vibration of the cations in two different
environments.

It is reasonable to assume that if lithium perchlorate
exists in PC as solvent separated ion pairs, then it is probable
that sodium perchlorate likewise forms the same type of species in
this solvent. Except for the thiocyanate and within experimental
error, the solvation bands of the sodium ion appear at the same
frequency, which is indicative of solvent separated ion pairing.
Thus, for sodium thiocyanate, the band at 188 cm"1 would be due to
the cation vibrating in a cage of solvent molecules only, and the
shoulder at 220 cm'] due to vibration in a cage with one thiocyanate
ion present. By a similar argument, all ammonium ion vibration
frequencies would indicate solvent separated ion pairs in solution.
For ammonium thiocyanate solution spectra, the shoulder at 186 cm.1
would be evidence of solvent separated ion pairs and the band at

220 cm'1 evidence of contact ion pairs.

27

However, if these arguments are correct, then one would
expect that at lower concentrations, only the band due to the
solvent separated ion pair would be present. The band due to the
contact ion pair should become larger as the salt concentration is
increased. This was not the case. For solutions of all three
salts, the bands appeared to have the same shape at all concentrations.
The situation for lithium thiocyanate solutions was even more
baffling. In this case, the band as lower frequency did not appear
in the same position as the bands for the lithium salt solutions

consisting of solvent separated ion pairs.

3. Beer's law studies. Studies of the integrated absorbance

 

of the solvation band as a function of salt concentration were done
for lithium perchlorate, lithium bromide, sodium perchlorate, and
sodium iodide solutions. The area of the bands from spectra

similar to those in Figure 2 were measured with a planimeter
calibrated to 60 units per square inch. The numbers assigned to

the areas have no absolute significance, but since only the relative
areas were needed, this was not important. The results appear in
Tables 4-7 and in Figures 4-7.

In PC solutions, lithium bromide and perchlorate very
likely exist mainly as contact ion pairs and solvent separated ion
pairs respectively. Since the solvation band frequencies did not
change in the range of salt concentration used, it is also likely
that the solutions consist of the respective type of ion pair at
all concentrations. Therefore, it would be expected that the

Beer's law plot for each of these salts would be a straight line.

28

Table 4. Solvation Band Area !§_Concentration for Lithium

Perchlorate in Propylene Carbonate.

 

 

 

Concentration (M) Area
0.075 118
0.154 226
0.266 415
0.310 482
0.451 690
0.572 931
0 . 640 1060
0.731 1216
0.858 1286
0.903 1432

 

Table 5. Solvation Band Area v§_Concentration for Lithium Bromide

in Propylene Carbonate.

 

 

 

Concentration (M) Area
0.132 196
0.267 390
0.377 557
0.477 725
0.528 805
0.680 1026
0.772 1084
0.919 1352

0.978 - 1344

 

29

Table 6. Solvation Band Area !§_Concentration for Sodium Perchlorate

in Propylene Carbonate.

 

 

 

Concentration (M) Area
0.060 29
0.184 107
0.308 156
0.343 162
0.466 235
0.492 235
0.591 302
0.591 281
0. 627 406
0.627 279
0.696 307
0.696 295
0.872 548
0.872 528
1.007 499

 

30

Table 7. Solvation Band Area v§_Concentration for Sodium Iodide

in Propylene Carbonate.

 

 

 

Concentration (M) Area
0.111 38
0.111 57
0.244 136
0.244 109
0.369 193
0.369 196
0.447 241
0.447 227
0.554 325
0.554 284
0.703 371
0.703 336
0.758 419
0.758 388
0.867 458
0.867 460
1.038 568
1.038 528

 

31

Figure 4

Beer's law plot for lithium perchlorate in propylene carbonate.

32

1500

 

iooo,

AIREAI

5001

 

 

0 4 1 I l
I

 

 

00 (L2 034 (16 (18

CONCENTRATION (M)

LO

33

Figure 5

Beer‘s law plot for lithium bromide in propylene carbonate.

AREA

34

 

 

 

 

1500
O
O
1000* .
O
500--
o f 4. i 4. i
0.0 0.2 0.4 0-6 0.8 1.0

CONCENTRATION (M)

35

Figure 6

Beer's law plot for sodium perchlorate in propylene carbonate.

AREA

36

 

 

 

 

 

500
400*r
O
O
300* . .
O
200"
100*- .
O i J. i 'r i
0.0 0.2 0.4 0.6 0.8 1.0

CONCENTRATION (M)

1.2

37

Figure 7

Beer's law plot for sodium iodide in propylene carbonate.

AREA

38

 

600

500‘

400‘

3001

200‘

100-

 

 

 

 

0 1 1
(10 (15 L0

CONCENTRATION (M)

39

If the molar absorbance of the solvation band arising from the cation
vibration in a solution of mainly contact ion pairs were different
from that in a solution of solvent separated ion pairs, then the
slopes of the two lines would be different. A least squares analysis
on Figures 4 and 5 showed them to have slopes of 1594 i 106 and

1388 i 132 respectively. Within the standard error the slopes are
the same.

The sodium perchlorate and iodide Beer's law studies also
produced a straight line for each salt. The slopes of Figures 6 and
7 were calculated by least squares to be 545 i 92 and 536 i 29
respectively. In the case of sodium thiocyanate solutions, the
area of several bands for solutions of 0.95 M sodium thiocyanate
was determined by counting the squares under the peaks, assuming
the baseline to extend horizontally from the lowest point of the
spectrum. These areas averaged out to approximately 470 units.

This value coincides well with the value of 480 units at 0.951fi
from the sodium iodide plot and with 475 units from the sodium
perchlorate plot.

At this time, it appears that the use of Beer's law plots
is not a sensitive enough method to reveal anything definite about
the composition of the vibrating species in PC-alkali metal

solutions.

B. Near-Infrared Studies

 

It has been assumed that when a PC molecule solvates a cation,

the negative end of the dipole, or the carbonyl group, is oriented

40

toward the cation. To verify this assumption, the near-infrared
Spectra of various PC solutions were observed. The only band that
shifted with increasing concentration was the carbonyl stretch at
1798 cm-1, which shifted linearly with concentration to about

1789 cm'] at 0.73 M salt. From these data it would be reasonable
to say that the carbonyl group is indeed the solvation site.

The ir vibrations of the thiocyanate ion were also observed
in the hopes of explaining the solvation peaks observed for
thiocyanate solutions in the far ir. However, the only thiocyanate
vibration which was not under a large solvent band appeared at

1553 cm".

Its frequency did not shift either with a change in the
concentration of salt or with the cation. Thus, it appears that
observation of the thiocyanate bands in the ir will not yield
useful information regarding the solvation of the cations in PC

solutions of thiocyanate salts.

C. Raman Mole Ratio Study_

 

A mole ratio study was done observing the shift of the 0]
Raman active perchlorate vibration as a function of R, where R =
the ratio of PC to perchlorate ion, with nitromethane as the inert
solvent. Bands for PC appeared at 958, 912, 850, 711, 630, 452,
and 187 cm_], a strong nitromethane band appeared at 917 cm-], and

the perchlorate band appeared at 932 cm'].

All band positions were
independent of R. This is a reasonable result, since lithium

perchlorate probably exists as solvent separated ion pairs in PC.

V. SUGGESTIONS FOR FUTURE WORK

1. Temperature studies are usually hindered by a small
solvent liquid range and by decreased solubilities of salts at low
temperatures. However, with PC solutions, far-infrared temperature
studies should be relatively easy, since the solvent has a practical
liquid range of about 320 C° and since alkali metal salts dissolve
in it fairly well.

2. Preliminary investigations on the pmr chemical shift of
protons on PC as a function of salt concentration have shown shifts
of only about 2 cps in going from pure PC to a 1.0 M_lithium
perchlorate solution. This is reasonable, since the nearest proton
is three bonds away from the solvation site, i.e., the carbonyl
oxygen. However, chemical shifts for sodium23 and lithium7 nmr
could yield useful information, since one would be observing the
change in electron density around the solvated species itself.

Mole ratio studies of chemical shift y§_R (mole ratio) could lead to
the determination of salvation numbers for the sodium and lithium
ions in PC. Again, temperature studies could be very easily carried
out and could possible give valuable information. Pulsed nmr could
reveal the relaxation times of the various metal ions as a function
of R. These may also lead to the determination of solvation numbers
in PC for the cations used.

41

42

3. The position of the far ir solvation band of lithium
thiocyanate solutions in PC was difficult to determine at low
concentrations due to the poorly subtracted solvent band at 445 cm'].
If there were no interfering solvent bands in the 400 cm'] region,
the spectra of lithium thiocyanate solutions in ethylene carbonate
(EC) could be observed. The only structural difference between the
two solvents is that the methyl group of PC is replaced by hydrogen
in EC. The solvating properties of the two solvents should be similar

enough to allow parallel conclusions to be drawn regarding the species

present in their thiocyanate solutions.

BIBLIOGRAPHY

10.

11.
12.

13.

14.

15.
16.
17.
18.
19.

BIBLIOGRAPHY

J. C. Evans and G. Y-S. Lo, J. Phys. Chem., 99, 3223 (1967).

 

F. Edgell, A. T. Watts, J. L. Lyford, and W. M. Risen,
. Amer. Chem. Soc., 99, 1815 (1966).

 

F. Edgell, J. Lyford, Rose Wright, W. Risen, and Alan Watts,
bid., 99, 2240 (1970).

HE C—aZ

 

B. W. Maxey and A. I. Popov, 1919,, 99, 2230 (1967).
B. W. Maxey and A. I. Popov, 1919,, 91, 20 (1969).

J. L. Wuepper and A. I. Popov, 1919,, 91, 4352 (1969).
J. L. Wuepper and A. I. Popov, 1919,, 99, 1493 (1970).

M. K. Wong, W. J. McKinney, and A. I. Popov, J. Phys. Chem.,
19, 56 (1971).

 

M. K. Wong and A. I. Popov, J. Inorg. Nucl. Chem., 99, 1205 (1971).

 

D. R. Cogley, J. N. Butler and E. Grunwald, J. Phys. Chem.,
19, 1477 (1971).

 

L. Simeral and R. L. Amey, Ibid., 19, 1443 (1970).

"Propylene Carbonate, Technical Bulletin", Jefferson Chemical
Company, Houston, Texas, 99,, 1962.

V. Gutmann, ”Principles of Coordination Chemistry in Nonaqueous
Solutions", Springer-Verlag, Vienna, 1968, p. 19.

J. Courtot-Coupez and M. L'Her, C. R. Acad. Sci., Ser. C,
1970, p. 271.

 

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43

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44

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