• TABLE OF CONTENTS
HIDE
 Title Page
 Acknowledgement
 Table of Contents
 List of Tables
 List of Figures
 Introduction
 Experimental
 Results
 Discussion
 Appendix
 Reference
 Biography
 Copyright














Title: Solvolysis kinetics of pyridine-boranes.
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Title: Solvolysis kinetics of pyridine-boranes.
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Table of Contents
    Title Page
        Page i
    Acknowledgement
        Page ii
    Table of Contents
        Page iii
    List of Tables
        Page iv
    List of Figures
        Page v
        Page vi
    Introduction
        Page 1
        Page 2
    Experimental
        Page 3
        Page 4
        Page 5
        Page 6
        Page 7
        Page 8
        Page 9
        Page 10
    Results
        Page 11
        Page 12
        Page 13
        Page 14
        Page 15
        Page 16
        Page 17
        Page 18
        Page 19
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        Page 21
        Page 22
        Page 23
        Page 24
        Page 25
        Page 26
        Page 27
        Page 28
        Page 29
        Page 30
        Page 31
        Page 32
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        Page 58
        Page 59
        Page 60
        Page 61
        Page 62
        Page 63
        Page 64
    Discussion
        Page 65
        Page 66
        Page 67
        Page 68
        Page 69
        Page 70
        Page 71
        Page 72
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        Page 77
        Page 78
        Page 79
        Page 80
        Page 81
        Page 82
        Page 83
        Page 84
    Appendix
        Page 85
        Page 86
        Page 87
        Page 88
    Reference
        Page 89
        Page 90
    Biography
        Page 91
        Page 92
    Copyright
        Copyright
Full Text











SOLVOLYSIS KINETICS OF

PYRIDINE-BORANES











By

ERNEST RODMAN BIRNBAUM


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










UNIVERSITY OF FLORIDA
January, 1961
















Tic author tiches to hia:'.: his fricnic and nscociatoc for many

helpful rwu:;cotions and extre-ely stinu3latin; discussions,

To his faculty:. co-..itoo, who have gone far beyond the call

of dut: to C:epc31t.o the writin3 of this dissertation, the author

er:roseco his sincere appreciation*

For efforts in obtaining equi;:icnit a .C dimicalo, for a firn,

but patient gtuidanco, and for an interest in the author's well-being,

the author is dccp;l:, grateful to )Dr. tschklcitsch, his research

director.

The author would als o like to thank the Research Corporation

for its generous financial support.


AC11CWLn^-irrTIT















TABLE OF COITINTS


Page


Chapter I

Introduction .* . . . . . .

Chapter II

Experiuental2.. .. *

Chapter III

ReCul.ts * .*. * *. * *

Chapter IV

Discussion . * *. * * *

Statistical Appendix . . . . . * * *

Reforences .* * * a * * * *


iii













LIST OF TABLES


Table Page

1. Analysis of purified pyridine-borane . . 5 5

2. Amines w * * * * 5

3, Pyridine-borane rate constant data . . . . 31

4, 11ixed-solvent experiments, 50,00C. . . . . . 32

5. Surmary of kinetic data for 4-picoline-borane . 40

6, Su a ru of kinetic data for 3-picoline-borane . 43

7. S'u-:ar. of kinetic data for 2-picoline-borane .. 6
8. Sumary of kinetic data for 2-othylpyridine-borane 4 . 19

9. Sunrmary of kinetic data for 2,4-1utidine-borane . 52

10. Sunmary of kinetic data for 2,6-lutidineoborane . .

11. Kinetic data for 2,h,6-collidine-borane . 5. . 7

12, Fit of rate constants to Arrhenius equation,
log k A+ B/T (sece,)1 . . . 58

13. Physicochemical data for the pyridines . . 77

!h, Test for homogeneity of several least squaes slopes . 88















LIST OF FIGURES


Figure

1.

2.

3.

4.

5.



7.

8,.

9.

10.

lla.

lib.

lc,.

12,

13.

3i.

15.

16.

17.

18.

19.

20.


es

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Fyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane

Pyridine-borane


* 0 . 0 9 *



* 09. 0 99900. 9 . 9

00. 090*99 9 99.


*,

0999


@9900

* 99*


43.120C.

43.120C.

4.5490C.

5.4L9C .

L7.6300.

47.630C

47.630C.

50-000.

50.000C.

50.000c.

506000C.,
50.OCC.
0.000oo.
5o.oooc.



$o.Oooc.


S6


Behavior

Behavior

Behavior


Pyridine-borano, 52.430C. . .

Pyridine-borane, 52.30C . .

Pyridine-borane, 52.330. .

Pyridine-borane, 54$260C. . .

Pyridine-borane, 54.260C. . .

Arrhenius Plot for pyridine-borane

Pyridine-borane 5~.300C. Iagnetica:

Pyridine-borane h5424C. Added pyr:

Pyridine-borane 54.24OC. Added pyr:


over

over

over


* 0 a *

several h

several hi

several h



* . *


alf-lives

alf-lives

alf-lives


ly stirred . . .


idine, 0,26M

idine, 0.421I


* . .9

* S 9 5 9


* 9

9


9.. SO 9* 9 *999*


Page

12

13

14

15

16

17

18

19

20

21

22

23

24

25

26

27

28

29

34

35

36

37


* 09

9..










LIST OF FIOURES--(Continued)


Page

Pyridine-borane 50.00C. Mixed solvent 30% water 70%
propanol by weight *. *. *. . . 38


Arrhenius plot for

Arrhenius plot for

Arrhenius plot for

Arrhenius plot for

Arrhenius plot for

Arrhenius plot for

Rate dependence on


4-picoline-borane . .

3-picoline-borane . .

2-picoline-borane . .

2-ethylpyridine-borane .

2,4-lutidine-borane .

2,6-lutidine-borane .

alcohol concentration .


29, Kirkwood plot for propanol-water mixtures at

30. Steric requirement of transition state relat
of methanesulfonic acid . . . .

31. Steric requirement of transition state relate
of a borane group .. . .


* 4 5 *


ive to


Ive to
...


that
* *t

that
...


32. Steric requirement of transition state relative to that
of boron trifluoride . . . *. . . .


Figures

21.


22.

23.

24.

25.

26.

27.

28.














CHAPTER I


INTRODUCTION


The earliest known addition compound between a tertiary amine

and a borane group, trimethylanine-borane, was synthesized in 1937 by

Burg and Schlesingera. Originally prepared by the base displacement

reaction between borine carbonyl and trimethylamine, the amine-borane

was also formed when trimethylamine and diborane were brought together.

Melting at 940C., trimethylamine-borane is a white, crystalline solid

which can be heated for severalhours at 125C. without a detectable

change in its physical properties. Even though a number of physico-

chemical investigations, including electron diffraction,2 x-ray

diffraction,3 Raman and infrared spectra,4,5j6 dipole moment measure-

ments,7 heats of reaction,8 B-U magnetic resonance spectra,9 and use

as a stereo-specific reducing agent,10 have been made on tritaethylamine-

borane, relatively little knowledge has been obtained about the reac-

tivity of its boron-nitrogen or boron-hydrogen bonds.

Five years after the discovery of trimethylamine-borane, Brown,

Schlesinger, and Cardonll reacted pyridine with diborane to produce

pyridine-borane. Utilizing the analogous reaction in nitrobenzene

solutions of pyridine and various alkyl-substituted pyridines, Brown

and Domash12 in 1956 prepared pyridine-borane and seven additional

alkyl-substituted pyridine-boranes, and measured the heats of reaction

with diborane calorimetrically,











The pyridine-boranes are either colorless liquids or white,

crystalline solids. In contrast to trimethylamine-borane, however,

they are not stable to heat, but readily evolve hydrogen at tempera-

tures below 1000C. to form red resins. The few investigations of

pyridine-borane to be found in the literature concern its cryoscopic

molecular weight and differential thermal analysis,13 its dipole

moment,7 its infrared spectra,5'6 its synthesis and typical reac-

tions,1 and its B-11 magnetic resonance spectra.9 No further ex-

perimental work has been reported for the alkyl-substituted pyridine-

boranes.

Since the heats of formation were the only data common to all

of the tertiary amine-boranes, and the sole data besides melting

points for the substituted pyridine-boranes, it was proposed to

undertake a series of quantitative studies relating the reactivities

of these compounds to their structures. The results of the first

set of experiments in the series, which pertain to trimethylamine-

borane, have already been published, Comprising the second part

of the series, the experiments to be described in this dissertation

will be devoted to a kinetic investigation of the solvolysis reaction

of the pyridine-boranes in n-propyl alcohol

C9Hgmrl3 + 3C3HH ----> 3H2 + CAsN + B(oC3H7)3
The purpose of these experiments is to determine if there is a

correlation between the mechanism and rates of solvolysis of the

pyridine-boranes, and their structures and heats of formation.














CHAP'ITFR II


FYXPJ1RTEITTAL


Chemicals .-A sample of Callery Chemical Company's pyridine-

borane used for preliminary experiments was a clear, colorless liquid.16

Purification was easily accomplished by dissolving the amine-borane

in ether, extracting with dilute sodium hydroxide, discarding the

aqueous phase, and removing the ether from the ethereal phase by

vacuum distillation, The purified material was analyzed for active

hydrogen by a slight modification of the iodate method.17 Because

of the limited miscibility of pyridine-borane in water, it was

necessary to dissolve the weighed sample of amine-borane in a few ml,

of n-propanol, and then add the excess standard potassium iodato

immediately to this solution. No difference in iodate titer was

found when standard potassium iodato solution containing 10 ml, of

n-propanol was converted to iodine with excess acid and potassium

iodide, neutralized with sodium bicarbonate, and titrated to the

starch end point with standard sodium arsenite.

A boron determination was also made on the purified pyridine-

borane. To the methanol solution of the amine-borane, concentrated

hydrochloric acid was added crop:wio until there was no further

evolution of hydrogen. An additional 5 ml. of concentrated acid and

10 grams of anhydrous calcium chloride were added, and the resultant

solution heated to distill methyl borate into a delivery tube immersed

in a beaker of water. After neutralization of the hydrochloric acid

3











which distilled along with the methyl borate, the determination was

completed by titration of the boric acidmannitol complex with

standard base. Table 1 contained the results of these analyses.

When it became necessary to obtain additional pyridine-borane,

a pound was purchased from the Callery Chemical Company. This

material turned out to be a somewhat turbid, distinctly yellow liquid.

Application of the above purification procedure in conjunction with a

filtration through alumina or magnesia did not produce a stable, color-

less product. The method of purification finally adopted consisted of

putting the pyridine-borane over calcium hydride, slowly heating it to

700C. with continual pumping in a vacuum system, and vacuum distil-

ling it at 70-800C. onto a cold finger. Samples purified in this

manner were clear, colorless liquids which typically gave active

hydrogen analyses within 97.3-983% of the theoretical values. No

difference in kinetic behavior was observed between the several

purified batches of pyridine-borane.

All of the amines used in this study were purified by 8-12

hours of stirring over calcium hydride, and distillation from the

calcium hydride on a Todd column at a 0:1t reflux ration. Table 2

lists the pertinent data. Synthesis of the alkyl-substituted pyridine-

boranes was accomplished by bubbling diborane, generated by the method

of Brown and Tierney,18 from a delivery tube immersed in mercury into

an excess of the respective free amines. 'Host of the excess amine

was removed from the solid pyridine-boranes by decantation; the re-

mainder was taken off by continual pumping in a vacuum system for

several hours. The liquid pyridine-boranes were subjected to high










TABLE 1

ANALYSIS OF PURIFIED PYRIDINE-BORANE


Weight of pyridine-borane Calc. Value Exptl., Value

0.1149 gram 3.71 mr.mole H" 3.67 m.mole IH"

0.1332 gram 4.30 m.mole H" 4.26 m.mole H"

0.2882 gram 3.10 m.mole B 3.06 m.mole B
0.2636 gram 2.84 rm.mole B 2.81 m.mole B



TABLE 2

AIMIES


Amine Manufacturer Distilling range oC.

4-picolino Matheson, Coleman, &. Bell 144.0-144.7

3-picoline Matheson, Coleman, &* Bell 143.5-143.8
2-picoline Matheson, Coleman, &. Bell 128.0-128.0

2-ethylpyridine Reilly Tar &, Chemical Corp. 148.8-148.8

4-ethylpyridine Reilly Tar &. Chemical Corp. 167.0-167.3

2,h-lutidine Reilly Tar &. Chemical Corp. 156.9-157.1

2,6-lutidine Reilly Tar &. Chemical Corp. 142.0-143.0

2,4,6-collidine Reilly Tar &. Chemical Crop. 170.0-170.0











vacuum overnight, or until no further reduction in volume was observed.

4-Picoline-borane was recrystallized from n-nonane as a white,

crystalline solid melting at 72-73oC. It was analyzed by the

Schwartzkopf licroanalytical Laboratory with the following results:

% C- calc*, 67.37; found, 67.20

% N- calc., 13,10; found, 13.32

% B- calc., 10.12; found, 10.31

% H, calc., 9.42; found, 9.66

A liquid, 3-picoline-borane was purified by vacuum distil-

lation onto a cold finger. Active hydrogen analyses gave 98.3 and

97.1% of the theoretical values.

Repeated recrystallizations of 2-picoline-borane from n-nonane,

1:1 benzene-2,3,5-trirethylhexane, or benzene-n-hexane mixtures failed

to yield a product melting above 46-08C., (literature value 50-510C.)

or analyzing higher than 95.5% of the calculated active hydrogen.

Following a recrystallization from n-nonane, 2- thylpyridine-

borane melted at 49.8-51.20C., and gave an active hydrogen content

97.4 and 97.3'" of the theoretical values.

4-Ethylpyridine-borane, a liquid at room temperature, was

not further purified, characterized, or used for Kiinetic studies.

2,4-Lutidine-borane is a white, crystalline solid. Both the

material synthesized as above and a sample from the Gallery Chemical

Company were recrystallized from n-nonane. The melting points were

76.5-77.70C. for the fonrer, and 76.5-77.5C. for the latter. The

Callcry product was analyzed by the Schwarzkopf Microanalytical

Laboratory as follows:











% C- calc., 69.49; found, 69.69

% H- calc., 10.0 ; found, 9.98

% N- calc., 11.58; found, 11.17

In addition to repeated recrystallizations, a vacuum sublima-

tion was tried to purify 2,6-lutidine-borane. The highest purity

obtained was 96% of the theoretical active hydrogen content.

A white, crystalline solid from n-nonane, 2,h,6-collidine-

borane melted at 99.2,100,30C. The Schwarzkopf Microanalytical

Laboratory also analyzed this material:

% C- calc., 71,16; found, 71.32

% H- calc., 10.45; found, 10.39

% N- calc., 10.37; found, 10.12

Fisher Scientific Company's rcagent grade normal propyl

alcohol was purified by refluxing it over anhydrous barium oxide for

8-12 hours, and distilling it on a Todd column at a reflux ratio of

50Sl. About the first and last 10% would be discarded, while the

middle fraction, distilling between 97.0 and 90.00C, would be collected.

Purified grade p-dioxane uas also obtained from the Fisher

Scientific Company. It was refluxed over calcium hydride for several

days and distilled from the calcium hydride just before use. Distil-

ling at 1010C., this material gave a negative test for peroxide upon

addition of potassium iodide, dilute hydrochloric acid, and starch.

Apparatus.--An E. H. Sargent &. Co. "Thermonitor" control and

heater-circulator were used with a Raytheon voltage-regulating trans-

former to provide the constant temperature bath required for kinetic
studies. In addition, the rlass tank supplied with the Sar;,nt unit











was fitted with an aluminum top and surrounded with a 2-3 inch layer

of venniculite. As measured by a Beckmann thernometor, the bath

could be maintained to within t 0.0050C at any temperature between

30 and 6500. For temperatures below 30C., the Sargent "Thermonitor"

unit was used in conjunction with a temperature-controlled, refrig-

erated water bath. In this case the Sargent glass tank containing

the "Thermonitor" heater-circulator was placed directly into the

reservoir of the cooling bath, and the thermostat of the latter was

set a few degrees below the desired te::iperature. The "Thermonitor"

control was then adjusted to bring the temperature inside the Sargent

bath up to, and maintained at, the desired value. The constancy of
this arrangement was t 0.030C. maximum at 1000., and less than t 0.01C.

for tem-peratures between 18 and 300C.

Procedure.--All of the kinetic experiments to be reported in

this dissertation involved determining the "reducing strength," i.e.,

the number of milli-equivalents of reducible species per 5 ml. sample,

of a reaction mixture as a function of time by means of the iodate

method.17 For convenience the "reducing strength" will henceforth be

designated by the abbreviation RS. The detailed kinetic procedure, in

order of operation, consisted of the following: (1) direct weighing

of the anine-borane into a 100 r.l volumetric flask used as a reac-

tion vessel; (2) diluting to the mark with n-propanol at room tempera-

ture; (3) through mixing of the solution and discarding of about 10 ml.

so that there would be no liquid in the neck of the flask; (4) immersing

the volumetric flask in the constant temperature bath to within about

an inch of the standard taper and waiting a minimum of 30 minutes for











therntal equilibrium to be established; (5) pipetting 5$ ml of the

reaction mixture directly into an excess of standard potassium iodate;

(6) adding solid reagent grade potassium iodide to the iodate solu-

tion; (7) acidifying the iodate-iodide solution with an excess of 2 I

sulfuric or hydrochloric acid; (C) neutralizing the excess acid with

excess reagent grade sodium bicarbonate; and (9) titrating the lib-

erated iodine to the starch end point with standard sodium arsenite.

The amount of 0,01 IT standard potassium iodate solution taken

for step (5), as measured by a 50 ml. burette, was regulated to give

an approximately constant back titration. Tb minimize experimental

error, the excess iodate--and hence the arsenite titration--was kept

as small as possible consistent with obtaining a quantitative oxida-

tion of the kinetic sample. Typically the arsenite back titration

amounted to 1 ml. of 0.025 LN solution. Pipetting presented a problem

at high reaction rates due to the copious evolution of hydrogen bubbles.

It was found that shaking the volumetric flask prior to taking a

sample eliminated most of this trouble. A further precaution taken

in connection with pipetting was to rinse the pipette with methanol

and dry it on an aspirator after each sampling.

lTmperaturo differences were measured by using three Beckmann

thermometers which had been set so that the bottom range of one over-

lapped with the top range of another. Once a temperatare on a O.10,

thermometer was taken as an arbitrary standard, temperature intervals ,

of 12-140C, could be determined to within a few thousandths of one

dcgrce centigrade. Time intervals ranging from 15 minutes to several








10


hours were obtained to within $ seconds by means of a stopwatch.

With the exceptions of 2,6-lutidine-borane and 2,h,6-collidine-

borane, at least tLo kinetic "runs" of widely differing concentration

were made at each temperature, giving an estimate of the experimental

precision. Tnose concentrations were usually between 0.06 and

0.1 N.














CHAP TER III


RESULT


Very similar kinetic results were obtained for all of the

pyridine-boranes. Those for the alkyl-substituted pyridine-boranes

will be given in tabular form; while those for pyridine-borane itself

will be presented graphically as a representative example. In either

treatment the time scale for many of the experiments has been shifted

so that zero time refers to the time the first sample was taken, and

not to the time the experiment was begun. This change will not affect

the interpretation of the data, because as will be considered subse-

quently, the rate constant for a first-order reaction is independent

of both initial time and concentration.

Since propyl alcohol was used as the solvent for these studies,

its concentration rained essentially constant during the course of

a kinetic experiment, making the solvolysis reaction ajpiarcntly zero-

order in alcohol. Graphs of log (RS), log (milli-equivalents of

reducible species/5 ml. sample), against time will be utilized to

present the results of the observed first-order reaction in amine-

borane, These results for pyridine-borane are shown in figures 1 to

10, 11-a, and 12 to 16. The size of the circles representing the

data is an estimate of the relative experimental error. Included on

each graph is the least squares slope, as determined from those points

which fall on the line.









12
























Lea
-0.
.600




60
0.560
+
CMJ
CM
r--t

.520-



69 309
Time (Min.)
Figure 1. Pyridine-borane, '3.120C.


st Squares Slope
000218 (min.)-l


I9
549











.820.


Least Squares Slope
-0.000211 (min.)


.7800
.O



* .7)40





I I I I
131 251 371 491
Time (Min.)
Figure 2. Pyridine-borane, h3.120C.





















Least Squares Slope 30
-0.000292 (min.) 300




.240




180 I




120





60



.600 .580 .560 .54o
1.222 + log (RS)
Figure 3. Pyridine-borane, 45.90C.


























Least Squares Slope
-0.000290 (min.)'-
.800





.780


+
CM
CM

S.760






.7h0- -



60 120 180 240 300
Time (Min.)
Figure 4. Pyridine-borane, 454.90C.









Least Squares Slo e
-0.000378 (min.)"


.680





.660

0

+

c .6o0 o





.620






120 180 2h0 300
Time (Min.)
Figure 5. Pyridine-borane, 47.630C.























Least Squares Slope
-0.000398 (min.) -


.660




.6ho





.620




.600


0o 100 160 220
Time (Miin.)
Figure 6. Pyridine-borane, 47.630C.











Least Squares Slope
0-.000353 (min.)-1


.760


0o

+
0H



.700










16b 320b 80
Time (Min.)
Figure 7. Pyridine-borane, 47.630C.


.82C














Least Squares Slope
-0.000491 (min.)-1


Time (Hin.)
Pyridine-borane, 50.C00C.


.300.






. 280

0


C"
+
c\J
j .260
r-





.240.


Figure 8.


























Least Squares Slope
-0.000o93 (min.)-1


180


Pyridine-borane,


.980





.960


.900


Time (Min.)
50.000.


Figure 9.




























Least Squares Slope
-0.0oo093 (min.)--


Figure 10.


60 120
Time (Min.)
Pyridine-borane, 50.00C.


.640


o .620
r-I
+
C\1
(14
(04
*
r-6
.600













Least Squares Slope 0.000490 (min.)-1
Sun of squares of deviations from Least
Squares Line 0.000018.
.48o0




0 .h60

CM
C\M


.o44





S10hho
51 100 150


Figure 11-a.


Time (Min.)
Pyridine-borane, 50.000C. Behavior over several half-lives.











Least Squares Slope
-0.000489 (min.)--


.300





. 200




C%
S.100






Jo 5o0 700 o9b
Time (Min.)


Figure 11-b.


Pyridine-borane, 50.00C.


Behavior over several half-lives.












Least Squares Slope
-0.000530 (min.)-


Figure 11-c. Pyridine-borane,
50.00C. Behavior over several
half-lives.


1600
Time (Min.)


100


2000


0

.650
CM
CM
(M
CM
oJ
oJ


1200


1400


1 I _4I I I












Least Squares Slope
-0.000631 (min.)"


Pyridine-borane, 52.h30C.


Time (Min.)


0.


(\j
cuj
cM
Hu


Figure 12.
















Least Squares Slope
-0.000632 (min.)-


Time (Min.)


Figure 13. Pyridine-borane, 52.h30C.


.640

0


Cr-
CM
C\M
, .620.






.600.












Least Squares Slope
-0.000617 (min.)-1


Figure 14.


60 90
Time (Min.)
Pyridine-borane, 52.430C.


.94o


, .900
bf
0
r-1
+
C\J
cli
CM
8 .880






.860


180












Least Squares Slope
-0.000788 (min.)-


Figure 15.


Time (Min.)
Pyridine-borane, 5h.26C.


.560





S.520
0



.480





.h4o













Least Squares Sloe
-0.C00770 (min.)-


120


Figure 16. Pyridine-borane, 5h.260C.


_ II I I _ _


180 2h6 30b 360
Time (iin.)


.800.


.7Q0









S.680
b0
0
O
+
C(M
CM




.620.










Pseudo-first-order rate constants for pyridine-borane were
calculated from the average slope of two or more experiments with
different initial concentrations at each temperature. A swumary of

this rate constant data is given in table 3. Figure 17 is the Arrhenius

activation energy plot, i.e., log k against l/T, for pyridine-borane.

From a least square analysis (see statistical appendix), the equation

of the Arrhenius plot was detcri-ained ast

loc k 11.05 -S100/T (sec.)"

with a standard error of 1 0.10 for the intercept and t 31 for the

slope, iMultiplication of the slope by -2.303 R yielded an experi.

mental activation energy of 23.34 t 0.1l kcal./mole.

The results of a lacutica1y stirred pyridine-borane "run"

at 54.300C, are given in figure 1G. Showing the continued linearity

of the log (RS) vs. tine plots, figures 11-b and 11-c extend the o50C,

experiment of figure 11-a to almost 3" half-lives. For the combined

data of these three figures, the overall rate constant was found to

be 0.00117 (min.)-1. Added pyridine 0.26 and 0.h2 M, figures 19 and 20,

yielded respective rate constants of 0.00193 and 0.00188 (min.)"l at

54.20c.

Table h .ivcs log (PS) and time values together with the least

squares slopes for a number of pyridine-borane experiments using

propanol-water mixtures as solvents, and for one experiment using a 50% by

volume p-dioxane-propanol .mixture, all at 500.00C. As typical of the

results from these mixed-solvent experiments, figure 21 contains a

plot of the data for the 307 water-70% propanol by vei-ht solvent

mixture,









TABLE 3

TFPYITIE-:3n'?A;.E rPJ-T CONSTANT DATA


Temp* 0C. -Avg. Slope k = -2.303 Slope -log k (min.)-1

43.12 0.000216 o0.oo497 3.304
45.*9 0.000291 0.000670 3.174
47.63 0.000376 o0.oo866 3.063
50.00 0.000492 0.00113 2.947
52.43 0.000627 0.00144 2.842
54.26 0.000779 0.00179 2.747







TABLE 4
IilXFE-SOLVFJiT EXPERIME1 S 50.0000C,


10% water 20% water 3c0 water 4 water
90% PrOH by wt. 80C PrOH by wt. 70% PrOH by wt. 60% PrOl by wt.


1 + log (RS) Tine (Hr.)

0.596 0
0.513 5.5
0.470 8
0.434 10.5

0.387 13.5

0.339 16.5
0.284 20

Least Squares Slope

-0.000260 (min.)-1


1 + log (RS) Time (Hr.)

0.618 0

0.603 1.5

0.524 8

0.494 10.5
0.458 13.5

0.423 16.5

0.383 20
Least Squares Slope

-0.000198 (min.)"-


1 + log (RS) Time (Hr.)

0.683 0

0.670 1.5

0.635 5.5
0.610 8

0.589 10.5
o.561 13.5

0.533 16.5
Least Squares Slope

-0.000152 (min.)-1


1 + log (RS) Time (Hr.)

0.693 0
0.665. 3

0.639 6
0.614 9

0.587 12
0.562 15

0.536 18
Least SquaresSlope

-o.oool45 (nin.)"l


~~~----







TABLE 4--Continued

50, water $55.% water* 60Z water 50% p-dioxane
50% PrOH by wt. 44.6% PrOH by wt. 40 % PrOH by ut. 50% Pr0H by volume


1 + log (PS) Time (Hir.) 1.222 + log (RS) Time (Min.)

0.633 0 0.658 4,
0.610 3 0.654 75

0.588 6 0.651 120
0,568 9 0.629 285
0.547 12 0.612 435
0.524 15 0.606 495
0.502 18 0.597 570
Least Squares Slope 0.591 630
-0.000120 (min.)"1 Least Squares Slope

-0.000116 (min.)"'


1 + log (RS) Time (ir.) 1.222 + log (RS) Time (Min.)
0.656 0 0,572 37
0,639 3 0.546 53
0.620 6 0.523 70
0.601 9 0.499 85
0.582 12 0.479 100
o.566 15 0.453 115
0.542 18 0.43 130
Least Squares Slope O.412 143
-0,0001.o (min.' 0.596 156

0.370 172
Least Squares Slope
-0.001U9 (min.)1


*Calculated for a 50% by volume solution at 2000.








-2.741Q 34






Least Square







-2.940








I
.r--



0
-3.14o















-3.340

305 310
1/T x 105


is Slope -5100


Figure 17. Arrhenius Plot for Pyridine-borane.











Least Squares Slope
-0.00788 (min.)"


180



120




60



I I 1
.550 .520 .490 .460 .430
1 + log (RS)
Figure 18. Pyridine-borane, 54.300C. Magnetically stirred.












Least Squares Slope
-0.000837 (min.)-


Figure 19. Pyridine-borane,


300 450
Time (Min.)
54.240C. Added Pyridine 0.26M


0h
45


CM
CM
C(




20
r-2








.250












Least Squares Slope
-0.000818 (min.) -


Figure 20.


1-u 300 450
Time (Min.)
Pyridine-borane, 5$.24C. Added Pyridine 0.42M


.650o











0


CMj
CMj
CM










.2501















Least Squares Slope
-0.000152 (min.)


Pyridine-borane, 50.000C.


Time (Min.)
Mixed Solvent 30% Water, 70% Propanol by weight.


H..


Figure 21.








39

The kinetic results for the alkyl-substituted pyridine-

boranes are given in tables 5 to 11 and figures 22 to 27* Table 12

summarizes the derived thermodynamic quantities (see discussion ) for

all of the pyridine-boranes,







TABLE 5
SUI2IAY OF KINETIC TDTA FOR 4-PICOLINKE-DORAJE


58.2o00. 58.200C. 60.100C. 60.10C,

1 + log (iRS) Time (-in.) 1 + log (RS Time (Min.) 1 + log (RS) Time (Min.) 1 + log (RS) Tine (Min.)

0.376 0 0.131 0 0.456 0 0.349 0
0.349 60 0.099 60 0.422 60 0.305 60

0.321 120 0.072 120 0.382 120 0.268 120
0.274 210 0.028 210 0.343 180 0.233 180
0.246 270 -0.015 270* 0.304 240 0,189 240
0.215 330 -0.039 330 0.267 300 0.151 300
0.184 390 -0.071 390* 0.227 360 0.112 360
Least Squares Slope Least Squares Slope Least Squares Slope Least Squares Slope
-0.000494 (rain.)"' -0.000509 (rin.)-1 -0.000640 (min.)"1 -0.000654 (min.)~1


*Not included in least squares calculation.








TABLE 5-Continued


62.00c. 62.00oc. 62.00oc. 64.90c.


1 + log (RS) Time (in.)

o.455 0
0.423 45

0.386 90

0.350 135

0.306 180

0.256 240

0.208 300

Least Squares Slope

-0.000837 (min.)"-


1 + log (RS) Time (MIn.)

0.277 0
0.243 45

0,205 90

0.180 120

0,151 150

0.125 180

0.100 211

Least Squares Slope

-0.00oo081 (min.)1


1 + log (RS) Timne (MiLn.)

0.581 0

0.558 30

0.533 60

0.507 90

0.483 120
0.456 150

o.434 180
Least Squares Slope

-0.000827 (min.)"1


1 + log (RS) Time (Uin.)

0.303 0
0.108 180

0.077 210

0.043 240
0.008 270

-0.037* 30oo*

-o.o37" 330*
-.0.067* 330*

-0.109 360

Least Squares Slope
-0.00109 (min.)-1


*Not included in least squares calculation.


---







TABLE 5--Continued

64.94tC. 66.43oC. 66.643C.


1 + log (RS) Time (Hin.)

0.526 0

0.493 30

0.461 60

0.430 90

0.397 120

0.364 150

0.332 180

Least Squares Slope

-0.00108 (min.)"'


1 + log (RS)Time (Min.)

0.504 o

0.471 30

0.433 60

0.393 90

0.356 120

0.277 180

0,213 225

Least Squares Slope

-0.00130 (rin,)"1


1 + log (PS) Time (Min.)

0.357 0

0.320 30

0.282 60

0.240 90

0.200 120

0.158 150

0.123 180

Least Squares Slope

-0.00131 (in.)-1


- -~-







TABLE 6
SLUP2ARY OF KINEEIC DATA FOR 3-PICOLINE-BORANE

52.70ca 52.70c.a 527.70oca 55.2000.

1 + log (RS) Time (Min.) 1 + log (RS) Time (lin.) 1 + log (RS) Time (Min.) 1 + log (RS) Time (Min.)
0.613 0 0.425 0 0.513b 0b 0.73 0

0.589 60 o.4oo 60 0.492 60 o0.36 75
0.564 120 0.362 150 0.460 150 0.ol5 135
0.537 180 0.338 210 0.434 210 0.374 195
Least Squares Slope 0.312 270 0.390b 270b 0.338 225
-0.00022 (min.)"1 0.289 330 0.383 330 0.302 315
0.265 390 0.358 390 0.273 375
Least Squares Slope 0.338 l50 Least Squares Slope
-0.00O11 (rmin.)" Least Squares Slope -.0o000o 2 (min,)-
-o.oOo5 (min.)-"


bNot included in least squares calculation.


aA weighted average was used to calculate the average slope for this temperature.







TABLE 6-Continued

55.20C. 57.470,C 57.47 C. 59.97C.


1 + log (RS) Time (Min.) 1
0.376 0

0.333 75
0.306 135

0.276 195

0.240 255

0.206 315
0.178 375
Least Squares Slope
-o.c0o530 (min.)"1


+ log (RS) Time ( in.)
0.532 0
0.501 45

0.465 90

0.434 135
0.405 180

0.359 240
0.332 275
Least Squares Slope
-0.000723 (rain.)-1


1 + log (RS) Time (MIin.)

o.449 0
o.417 45
0.385 90

0.353 135
0.314* 180*
0.281 240
0.248 275
Least Squares Slope
-0.000720 (min.)"-


1 + log (RS) Time (Min.)

0.517 0
0.484 30

0.443 75
0.401 120
0.358 165
0.321 210
0.278 255

Least Squares Slope
-0.000928 (min.)"1


,.ot included.in least squares calculation.








TABLE 6--Continued


59.97C. 62.070c. 62.07oC.


1 + log (RS) hime (Min.)

0.o38 0

O.409 30

0.368 75

0.325 120

0.283 165

0.239 210

0.200 255

Least Squares Slope

-0.c00937 (rin.)-1


1 + log (RS) Time (;in.)

0.560 0

0,522 30

0.685 60

0.455 90

0.419 120

0.383 150

0.323* 180*

Least Squares Slope

-0.00117 (min.).1


1 + log (RS)Time (HMin.)

0.462 0

C0.27 30

0.388 60

0.359 90

0.321 120

0.288 150

0.254 180

Least Squares Slope

-o.o0015 (min.)-1


*Not included in least squares calculation.


-







TABLE 7
SU1IIAEY OF KINETIC DATA FOR 2-PICOLINE-BORANE


39.6700. 39.670C. b1.800C. h1,800o ,


1 + log (RS) Time (Min.) 1

0.554 o
0.517 105
o.687 180

o.459 255
0.431 330

0.4065 405
0.376 80
Least Squares Slope
-0.000371 (min.)-1


+ log (RS) Time (Min.)

0.452 0

0.410 105
0.385 180
0,361 255

0.333 330
0.306 4o5
0.270 495
Least Squares Slope

-0.000361 (min.)-'


1 + log (RS) Time (Min.)

0.572 0

0,536 75
o.516 120

c0.87 180
0.460 240
0.429 300
0.395 375
Least Squares Slope

-0.000473 (rmin.)"-


1 + log (RS) Time (Min.)

0.470 o

0.433 75
0,o13 120
0.382 180

0.354 240
0.319 315
0.289 375
Least Squares Slope

-0.000481 (min,)"1








TABLE 7--Continued

43.950. 43.950. 46.1ic. 46.h44c.


1 + log (RS) Time (iIin.)

0.577 0
0.526 45

0.499 90

o.474 135
o.446 180
0.413 225
0.388 270
Least Squares Slope
-0.000624 (min.)Y1


1 + log (RS) Time (1ian.) 1

0.151 0

o0.424 15
0.391* 90*
0.376 120
0.362 150
0.344 180
0.311 225
0.286 270
0.25! 330

Least Squares Slope
-0.000609 (min.)"1


+ log (R3) Time (Min.) 1

0.569 0
0.533 45
0.457 135
0.421 180
0.386 225
0.347 270
0.311 315
Least Squares Slope

-0.000820 (rnin.)-


+ log (RS) Time (Min.)

0.474 0
0.438 45
0.l01 90

0.363 135
0.329 180
0.293 225
0.256 270
Least Squares Slope
-o.0oo806 (min.)-


*Not included in least squares calculation.








TA LLE 7--Continued
- -,-


1 + log (RS) Time (iin.)

0,598 0

o.566 30

0.536 60

0.5o0 90

.t470 120

0.437 150

oto6 100

Least Squares Slope

-0.00107 (in.)"


1 + log (RS) Time (;in.)

0.471 0

0.439 30

o.h4o 6o

0.373 90

0.337 120

0.308 150

0.274 180

Least Squares Slope

-0.00110 (n.)"


4h0.8h0c.


h8.8hc.


- I' --- -~-- ~--







TABLE 8
SU L:::A- OF :I::EzIC DATA FOR 2-ETYLPYrrLJ.^-LO...E


30.0000. 30.0000. 32.S0. 32.0Co
32- 0. 32. 0


1 + log (3S) Tioe (Min.)

0.543 0
0.525 60

0.510 120

0.479* 2o40
0.459 300

0,441 360
0.426 420
0.388 $40

Least Squares Slope

-0.000284 (rin.)-1


1 + log (0S) IThe (MIin.)

0.156 0
0.o45* 60*

0.424 120

0.389 240
0.372 300

0.39* 360*

0.336 420
0.320 400
0.302 $540
Least Squares Slope

-0.000287 (min.)"1


1 + log (PRS) Time (Hin.)

0.497* o*
0.467 60

0.446 120

0.420 180

0.396 240

0.375 300

0.348 360
0.325 420
Least Squares Slope

-0.000397 (min.)-l


1 + log (i1S) Time (Uin.)

0.389 o
0.363 60

0.341 120

0.319 180
0.298 240

0.272 300
0.246 360
Least Squares Slope

-0.000389 (min.)"1


*Not included in least squares calculation.







TABLE 8-Continued


34.70C. 34.780C. 37.270C. 37.270C.


1 + log (RS) Time (Min.) 1

0.537 0
o0.5G 60

0.471 120

O.039 180
0.408 2140

0.374 300

0,343 360
Least Squares Slope
-0.0005o 4 (~in.)1)


+ log (RS) Time (Min.) 1

O.!l1 0

0.l14 60

0.384 120
0.352 180

0.312 240
0.282 300

0.253 360
Least Squares Slope

-.OOCC055 (rin.)-1


+ log (RS) Time (CIin.)

o.560 0

0.529 45

0.494 90
0.6h4 135

0.).32 180

0.397 225

0.364 270
Least Squares Slope

-0.000725 (nin.)"1


1 + log (PS) Time (Min.)

0.&39 0

0.403 45

0.374 90

0.337 135
0.302 180

0.273 225

0.236 270

Least Squares Slope

-0.000747 (min.)"


~ ~-~ --~--~-i~-------~---








TAiLE 8--Continued


1 + log (RS) Tiro (1iin,)

0.534 0

o.490 o4

0.U59 90

o.hoo 135

o.356 ISO

0.310 225

0.267 270

Least Squares Slope

-0.0oC992 (min.)-1


1 + log (RS) Tine (.in)

0..1.6 0

0.368 45

0.341 75

0.310 105

0.285 135

0.256 165

0.226 195

0.196 225

Least Squares Slope

-0.000965 (min.)"1


39.770c


39.770C


~~ ---


~~- ~ ---~~~~----








TALE 9
Su:LAR*i OF KIIETIC DATA FOR 2s,-LUTIU.I:;E-OJA;;E


48.49c. 48.490c. 48.h90c. 50.950c.


1 + log (RS) Time (ain.)

0.55,9" 0

0.528 60

0.500 120

0.477 10,
C.l51 240
0.431* 300*

0.397 360

0.373 U20

0.315 180
Least Squares Slope

-0o.00033 (min.)-1


1 + log (ES) Time (Min.)

0.610 0

0.586 60

0.553 120*
0.530 130
c.So4 24o
0.473 300

0o.43* 360
o.419 h20
Least Squares Slope

-0.000458 (rain.)-1


1 + log (RS) Time (Iki,.)

O.410* 0
0.389 60
0.362 120

0.334 180
0.308 24o
0.282 300
0.260* 360*

0.226 h20

Least Squares Slope

-0.0ooo0051 (min.)"1


1 + log (RS) Time (Uin.)

0.531 o
0.502 60

0.h64 120

o.l31 180
o0.oo 240

0.357* 300*
0.328 360
0.294 420

Least Squares Slope

-.ooo0568 (min.)-1


*Not included in least squares calculation.


--







TABLE 9--Continued


50.950. 53.350C. 53.3$c. 55.870C.
~~- J Ii 1 I


1 + log (rS) Time (Min.)
O.462 0

0.434 60

0.394 120

0.363 180

0.334 2Loe

0.293 300
0.263 360
0.227 420
Least Squares Slope

-0.000563 (min.)-1


1 + log (RS) Time (min.)

0.563 0

0.525 45

0,95 90

0.457 135

0.424 180

0,391 225
0.355 270
Least Squares Slope

-0.000764 (min,.)1


1 + log (RS) Time (Iin.)

o.489* o*

0.463 45
0.!26 90

0.393 135

0.358 180
0.321 225

0.293 270
Least Squares Slope
-O.000762 (min.)-1


1 + log (is) TiMe (Min.)

0.519 0
0.468 45
0.428 90

0.382 135

0.349 165

0.323 195
0.295 225
Least Squares Slope

-0.000993 (min.)"1


*Not included in least squares calculation.


----I








TABLE 9--Continued


55.?870. 58.37C. 8.370
... ,, = , ,


1 + log (RS) Time (llin.)

o ,!5o o
0..50 0
O.J01 45

0.356 90

0.308 135

0.281 165

0.249 195
0.220 225
Least Squares Slope

-00.0102 (min.)-1


1 + log (RS) Time (iin.) 1

0.581 0

0.544 30

o0.oh 60

0.l60 90

0.19 120

0.373 150
0.312 180

Least Squares Slope

-0.00136 (nin.)"1


+ log (0S) Time (lin.)

0.482 0

0.432 30

0.395 60

0.35h 90
0.310 120

0.279 350

0.246* 180*

0.156 20o

Least Squarcs S1~pe
-0.00o13 (min.)


*Not included in least squares calculation.


- -







TA- LT 10
SUI'MAE OF KIC2TIC DATA FOR 2,6-LU1 Tl .-0iAvE


10.000c, 10.00C. 1.oo0%. 1.oo00c


1 + log (RS) Thme (Uin.)

0.596 0
o.589 60
o.570 165
0.561* 2,5"
0.532* 435
O.4h9 615
0.394 12P5
0.367 1?.25
0.331 1605
Least Squares Slope
.00o016o (min.)"1


1 + log (RS) Time (Miin.)
0.488 0
0.481 75

0.462 165
0.W 2 255
0.425* 435*
0.392 615
0.287 1245
0.254 125
0.227 1605
Least Squares Slope
-0.000165 (min,.)"1


1 + log (RS) Thie (MIin.)

0.577 0
0.558 90
0.535 165
0.510 240
0.492 315
O.466 390

o.h8 465
0.u11 570*

0.390 660
Least Squares Slope
-0.000288 (min.)1"


1 + log (PS) Time (;.in.)

o.471* o*
o.450 90
o.432 165
o.410 240
0.388 315
0.368 390
0.3t44 65
0.315 570
Least Squres Slope
-0.000285 (rain.)")


*Not included in least squares calculation.


--~-----~








T.'LL? 10--Contirnued


17,0800, 17.080c. 19.600c. 22,10o0.
.- .,, ,- ,


1 + log (RS) Tine (Lin.)

0.538 o

0.520 45
o.490 105

0.470 150

0.453 195
0.428 255

O.a08* 315*

0.365 375*

Least Squares Slope

-0.000436 (min.)'"


1 + log (;-s) Tine (Iin.)

0.413 0

0.1422 15

0.393 105

0.372 150
0.356 195

0.329 255

0.302 315
Least Squares Slope

-..cco'n4i (min,)-1


1 + log (RS) Tihe (Min.) 1

0.538 0

0.524 30

O.505 60

0.o189 90

o0.65 135

o.434 1830

0,419 210
O.402 240

0.386 270
Least Squares Slope
-0,000573 (min.)"1


+ log (RS) TiMe (rin.)

0.560 0

0.537 30
0.513 60

0.494 90

o0.69 120

0.-37* 15*
0,420 130

0.398 210

0.370 240
Least Squares Slope

-0.000785 (min,.)-


L*ot included in least squares calculation.











TABLE 11
KIi:ETITC DATA FOR 2,4,6-COLLIDIiE-BORANE


30.00CG.


1 + log (RS)

0.51 1
0.526

0.517

0.499
0.497

0.479
0.468

0.461

Leat Sqares Slope
Least Squares Slope


Time (Min.)

0

15

30

45
60

75

90

105
120
- 0.000752 (min.)"I


I"-- I~ ~` I ~- -- ""~~I










TABLE 12


FIT OF RATE CONSTANTS TO ARRHEU~IS EQUATION,
log k A + B/T (sec.)"
r 5 ; ~- A.
Amine-borane A -B I S e.u.
kcal../raole at 65$C.

pyridino 11.0+o0.10 5100+31 23.34+0.14 -10.9

4-picoline 12.13+0o.8 5576+282 25.52+1.29 -5.98

3-picoline 11.51+0.36 5316+117 24.33L+0.5 -8.79
2-picoline 11.59+0.12 5143+39 23.54+0.18 -8.43

2-ethlyl-
pyridine 12.32+0.39 5240+119 23.98+0.54 -5.07
2,4-lutidine 11.43+O.42 5213+139 23.86+0.64 -9.14
2,,6-utidine 11.h2+,0.40 704+114 21.52+0.52 -9.21















ji












-2.600





H -2.700





o -2.800





-2.900


Least Squares Slope -5576


259

Figure 22. Arrhenius


297/T x 10299 30
Plot for -Picoline-borane
Plot for h-Picoline-borane.


I I I i









60
















-2.600





-2.700





S-2.800


S4
0
-2.900





-3.00


I I


Least Squares Slope -5316


I I


306


298 300 302 30o
1/T x 105
Figure 23. Arrhenius Plot for 3-Picoline-borane.


I I I I


I






















Least Squares Slope -5143


I I I
312 311T x 16 318
l/T x -P l -
Figure 24. Arrhenius Plot for 2-Picoline-borane.


-2.65o




-2.75o




bfl

0--~~


-3.5o0
























-2-700 Least Squares Slope -5240





-2.800










bfl
-2.900




0
-3.000o






-3.100





-3.200



320 322 32h 326 32
1/T x 105
Figure 2$. Arrhenius Plot for 2-Ethylpyridine-borane.
























Least Squares Slope -5213


I I I I I


310


-2.650



I

S-2.750


hO
0

-2.850





-2.950


302 304 306 308
1/T x 10i
Figure 26. Arrhenius Plot for 2,4-Lutidine-borane.



















Least Squares Slope -h70h


1/T x 105
Figure 27. Arrhenius Plot for 2,6-Lutidine-borane.


-2.850


-2.95o.





-3.o0o





--3.15o
bfl




-3.250





-3.350














CHAPTER IV


DISCUSSION


During the course of these experiments the conversion of the

excess iodate to iodine and the arsenite back-titration were carried

out at times ranging from less than a minute to several hours after

the kinetic sample had been pipetted into the iodate. Since no

difference in kinetic results was observed, it would appear that the

oxidation of a borane group by iodate is an instantaneous reaction.

This is especially interesting considering the minimal excess of

iodate used here, and the large excess required for a quantitative

oxidation of borohydride.

Failure to obtain a constant 5 ml. siple because of hydrogen

evolution in the pipette was the largest single source of experimental

error. The magnitude of the scatter caused by hydrogen bubbles may

be seen by comparing figures 15, 16, and 18, While the interval

between sampling is about the same, 20-30 minutes, the experiment

which was magnetically stirred to purge the reaction mixture of hydrogen

shows considerably less scatter.

Since the iodate method gives the total amount of oxidizable

species, the RS value would be identical with the amine-borane

normality only if there were no siCnificant build-up of oxidizable

intermediates. Considering the many variations in these experiments,

i.e., change of amine-borane change of amine-borane concentration,

65










addition of amine, change of solvent composition, change of tempera-

ture, and investigation of differing extents of reaction, the excellent

fit of the data to the first order log (RS) versus time plots ar,-ucs

against the existence of any oxidizable ini.cr.ediales. ',;c PS values

will therefore be taken as being equivalent to arine-borane normalities.

The general rate equation for the solvolysis reaction is civcn

by

-d(amine-borane) = k1(iaine-borane)m (alcohol)n (1)
dt
Since the concentration of alcohol remained essentially constant rela-

tive to that of the amine-borane, (1) may be re-formulated ast

-d(amine-bor k( n-borane) k ine-borane (2)
dt
where k' = k (alcohol)n If "m" is set equal to 1, and the amine-

borane concentration exi.re
a, and x, the number of roles per liter which have reacted after time

t, equation (2) becomes

dx/dt = kI (a-x) (3)

A'trLr separation : the variables, (3) is readily inI.c rated to .ive

-ln(a-x) = k't + Consto (4)

When t O0, x = 0, so that the constant of (4) is simply -In a.

Whence (4) becomes

-ln(a-x) = k't -ln a (5)

A plot of lo:; (a-x) a:.ainst ti ie should thus be linear with a

slope -k'/2.303, and an intercept lo-; a. Fro.i- equations 3-5 it can be

seen that the rate of this pseudo-first-order reaction is independent

of both initial concentration and initial time, and is determined only

by !he concentration of unrcacted amine-borane. The excellent linearity











of the log (PS) versus time plots attests to the solvolysis being first

order in amine-borane.

When it was found that vigorous stirring would cause several

hundred milligrams of pyridine-borane to dissolve in 100 ml. of

water, both quantitative and qualitative investigations were made on

the rate of reaction of pyridine-borane with distilled water. In a

quantitative experiment at 5oC., only about 12% of the active hydrogen
had been lost in 19 hours, corres.ondirn; to a half-life of about 90 hr.

The half-life for the propanol solvolysis at this temperature is about

6 hours. After having remained at room temperature for a month, an

aqueous pyridine-borane solution readily reduced silver nitrate and

evolved hydrogen upon addition of dilute acid, On the basis of these

results it was assumed that within experimental error, water could be

used an an inert diluent for the propanol.

Returning to the relationship that k' = k(alcohol)n ,

log k' = log k + n log (alcohol) (6)

A plot of Ic-, k' against log molall concentration of alcohol) should

be linear with slope "n" and intercept log k. Figure 28 is such a

plot for pyridine-borane with pure p opanol and propanol-water mixtures.

'Tr- curvilinear relationship obtained in figure 28 is attributed to

strong electrostatic interactions of the amine-borane with the solvent,

causing log k to vary. This result i.- :it have been anticipated from

the vastly different slopes obtained with pyridine-borane for 50C

propanol-50C water and 50% propanol-5 C. dioxane by volume (table h).

Uhile the former had a slope several times smaller than the value for










-2.950








-3.150









0




O0


-3.550.




I I I I
.900 1.00 1.10 1.20
log molall concentration of alcohol)
Figure 28. Rate dependence on alcohol concentration
50.00C.










pure alcohol at the same temperature, the latter had a slope several

times larger than the pure alcohol slope. Further evidence of solvent

electrostatic interactions with pyridine-borane comes from the work of

il:hecva and Fedneva.13 These investigators found pyridine-borane to

be highly associated in benzene solutions, and only slightly associated

in the more polar nitrobenzene solutions of the same concentration.

Considering only strong electrostatic interactions, Kirkwood,19a

has derived a relationship which gives the effect of the dielectric

constant of the medium on the free ener;;:- of a polar noleculo. For a

rate process this relation becomes


-1 -o --+-- (7)
RT(aD+I) L 3 ri rJ
where k is the rate constant in medium of dielectric constant D; ko is

the rate constant in a condensed rcdi,.- of dielectric constant unity;

the / %3 refer to the dipole moments of reactants, A and B, and of the

transition state; and the r's refer to the effective radii of these
species. While the dielectric constant of the iKirkwood equation is

actually the value for the solution, it has been found that for a dilute

solution the dielectric constant of tile pure solvent nay be used.

For various mixtures of two inert solvents, a plot of log k

against (D-1)/(21+1) should be linear. In addition, if as a first

approximation the volume of the activated complex is taken to be equal

to the sum of the volumes of the reactants, a negative slope should be

obtained when the ground state has a larger dipole moment than the

transition state.










Dielectric constants for propanol-water mixtures by weight have
20
been detenained at 500C. by Akerlof.2 The dielectric constant for the

to,6% alcohol mixture was obtained from a plot of electricc constant

versus weight percent alcohol, which was found to be linear in the range

20-60W alcohol. If the rate data for propanol-water mixtures were used

with AKcrlof's dielectric constants, a linear iirkiood plot could be

realized only if the solvolysis were zero order in alcohol. The Kirkwood

plot is shown in figure 29. In addition to visual linearity, this plot

was analyzed statistically (see statistical appendix) for a linear

relationship between the two variables. The correlation coefficient

proved to be 0.99, giving a highly significant linear relationship

between log k and (D-1)/(21)+l). It is thus concluded that the solvolysis

reaction is zero order in alcohol and that the dipole moment of the

ground state is larger than that of the transition state. kt of equa-

tion (2) must then be identical to kr,

The calculation of Arrhenius activation energies was given in

chapter 3. Tntrcipies of activation were calculated as follows: The

"thermodynanic" treatment of reaction rates9b gives


kr = (kT/h)edS*/R e-H/RT (8)

where,4H" is given by

dH = EArr. RT (9)
for a reaction in solution in the liquid state. Taking the common

logaritlh of both sides of (8)

log kr = log(kT/h) + d S*/2.303 R -AI1/2.303 RT' (10)

Substitution of a particular temperature, in this case 338.20i., into



























-3.30





S-3.-




o -3.50.





-3.60.



-.63 .h69 .h75 .481
D-1
( 2D+1 )
Figure 29. Kirkwood plot for propanol-water mixtures at 500C.










the least squares equations of the Arrhenius plots gave the values of

log kr in (sec.)"1 for that temperature. If (kT/h) is taken as approxi-

mately 1013, and equation (9) is substituted into (10),

log kr 338,20i. = 13 +S*/2.303 R -(EArr.-RT)/2.303 RT (11)

Equation (11) nay thus be solved for the entropy of activation.

Consistent with this work would be a mechanism involving a slow

dissociation of the amine-borane into one or more reactive fragments

which would then undergo a fast reaction with the solvent to yield the

observed products. Two different dissociative mechanisms will be con-

sidered. The first of these would involve essentially an ionization of

the amine-borane to give hydride ion:


araine'BH3l )- (anine'BlH2) + H" fast
3 slow \ RH Products (12)

Equation (12) predicts a large generation of charge in the transition

state, and hence an increased rate with increasing dielectric constant

of the solvent, Exactly the opposite effect was found experimentally.

lBrg21has suggested that as the Lewis base stren-:th of the donor atom

is increased, the hydridic character of the borane complexes also is

increased. If (12) were the correct mechanism the Arrhenius activa-

tion energies should decrease with increasing base strengths for steri-

cally si.il.r pyridine-boranes. "The pKa values for the pyridine-boranes

are listed in table 13 along with a number of other physicochemical

quantities. Thus both 3-picoline-borane and I-picoline-borane should

have Arrhenius activation e.ncr, ics less than that for pyridine-borano,

while 2,t-lutidine-borane would be expected to have an Arrhenius acti-

vation .ner,~- less than that for 2-picoline-borane. Again just the

opposite result was obtained experimentally. This dissociation










mechanism is hence eliminated as being inconsistent with experiment.

The second dissociative mechanism involves a slow dissociation

of amine-borane into the amine and a borane group, followed by a fast

reaction of the latter with the solvent:

a~inb*3 --l--- aiiino + BII3 fast
sI ROH > Products (13)
The transition state in (13) would have a lesser separation of charge

than the ground state because of the lesser extent of dative bonding

between boron and nitrogen.

Hence a decreased rate with increasing dielectric constant of

the solvent would result. Looked at from a different viewpoint, the

polar amine-borane ground state (dipole moment of pyridine-borane

5.86 D.) would be stabilized relative to the non-polar transition state

by an increase in solvent dielectric constant, thus causing the rate to

decrease. Fxperimentally just such a decrease in rate with increasing

dielectric constant of the solvent was observed, If (13) were the

correct mechanism, the Arrhenius activation energies should increase

with increasing base strengths for sterically similar pyridino-boranes.

That such an increase was found experimentally has been mentioned

above, and is evident from the data in table 13.

Since the dissociative mechanism of equation (13) is in good

accord with experiment, it will be examined further. A steady-state

assumption for the concentration of borane yields the following rate

law t

d(products) = klk2 (,-'ine-borane) (alcohol)
dt k-. (amine) + k2 (alcohol) (14)











Depending upon the relative nagnitudes of the kl(azinine) and k2(alcohol)

terms, a variety of rate dependencies can be predicted. If they are

both of about equal magnitude, the rate would be first order in amine-

borane, and of c. .,A-,x order in alcohol and arine. If the k_-(amine)

toer isneglibible in comparison to the k2(alcohol) term, a shaple

first order reaction in amine-borane would be observed. If, however,

the opposite were true, the rate would be found to be first order in

aaine-'oorane, first order in alcohol, and inversely proportional to the

amine concentration.

It is apparent from the fonm of (lh) that either k.2 would need

to be very large or k2 very snall for the concentration of amine pro-

duced by the reaction to have a detectable effect on the rate in alcohol

solvent. Tn.5 the log (iS) plots in figures ll-b and 11-c are still

linear to almost 33 half-lives. While the slope has actually in-

creased in the latter plot, the overall least squares slope for figure

11 is only about le larger than the initial slope. Consideration of

the experimental error involved in determining the RS for times greater

than 2 half-lives has led to the conclusion that this difference in

slope is not significant.

The further attempts to detect a retardation by adding amine

0.26M and 0.14Z figures 19 and 20 respectively, were also fruitless

as the rates in these e;:,,ri c,-ts agreed with each other and with the

rate obtained in the absence of additional arinc, to within the

coti.-ited experimental error. It was not possible to use a higher

concentration of amine, because this substance interfered with the

analytical procedure by completing the iodine.











If, as these results indicate, the k.l(amine) term is negligible,
then equation (lh) becomes

d(products) = kl(anine-borane)
dt (15)

which is exactly the rate dependence obtained previously from a

consideration of the individual rate dependencies of amine-borane and

alcohol. Hence it is concluded that the solvolysis mechanism of

equation (13) is entirely consistent with the experimental results if

the recombination rate is assumed to be insignificant with respect to

the rate of the borane-alcohol reaction. The validity of this assump-

tion will be considered later. In nccepAi.t: equation (13) as the most

probable mechanism consistent witi experiment, it should again be noted

that any hypothetical mechanism involving the hydridic character of

the amine-borane would necessarily result in the fallacious prediction

of decreasing experimental activation ener':ics with increasing base

strength s for sterically similar pyridine-boranes.

Information as to the nature of the transition state for the

solvolysis reaction can be obtained by ai.plication of an in;genious

technique developed by Brown et al.22 These investigators found that

for 1"nar Lewis acid-base reactions of the pyridines an estimate of the

relative steric requirements of the acids could be determined from

plots of an energy function for the reaction of each acid with the

pyridines against an encr,-y function for some other reaction of the

pyridines. The ener:'- functions may be activation energies, heats of

association, rate constants, and even pKa values.

Three plots of this ty! e, the data for which nay be found in











table 13, are sho~m in figure 30-32. Following Broom's Cexaple,

energy differences between the values for the respective substituted

pyridines and that for pyridine itself are plotted, rather than Vhe

actual nr.rr:: va.l--,. Figure 30 is a plot of the Arrhenius activa-

tion cnr 'ic; for the solvolysis reaction :;:..rt the heats of reac-

tion of the pyridines with methane sulfonic acid. "Toto that the

ortho-substituted pyridines fall below the line formed by pyrid'L;ic,

3-picoline, and h-picoline.

A stxilar plot of the Arrhenius activation energies for the

solvolysis reaction .;-ain-t the heats of reaction of the amines with

diborane is shown in fi -'re 31. In this case, however, the ortho-

substituted pyridines lie above the line formed by pyridine, 3-picoline,

and h-picoline, Finally figure 32 is a plot of the Arrhenius activa-

tion ener-ics for the solvolysis reaction against the heats of associa-

tion of the i;-rL:V'i s with boron trifluoride. Here the ortho-sub-

stituted pyridines a-ain lie above the line formed by pyridine,

3-picolinnr and h-piicol:nin, but are much further displaced from the

line.

It has been shown by Brown et al.22 that the steric require-

ments for the acids considered in figures 30-32 increase in the Cx-

pected order proton boranleboron trifluoride. The generalization

may thus be made that reactions of the pyridines with acids having

a smaller steric requirement than the solvolysis transition state

will result in the ortho-substituted pyridines falling below the

pyridine, 3-picoline, and h-picoline line when the inr, ,'ii data are

plotted as above. Sin.larly acids with a greater steric rcquirer:ment










TABLE
T17SICOCI!FTJCAL DATA


13
FOR 7l r IP'YP.!I!.TS


Amine pKaa Heats of reaction, -kcal./molea Arrhonius b, c
Activation
CH1S3OH -1(BHl3) BF3 Energes


pyridine 5.17 17.1 17.9 32.9 23.34

4-picoline 6.02 18,. 18.5 33.4 25.52

3-picoline 5.68 17.8 18.2 33.2 24.33

2-picoline 5.97 18.3 17.2 31.2 23.54

2-ethyl-
pyridino 5.92 18.2 16.9 30.6 23.98

2,4-lutidine 6.79d 23.86

2,6-lutidine 6.75 19.5 16.3 25.4 21.53


a reference 22

b this work

c kcal./mole
d Detennination of Organic Structures by Physical Methods, E. A.
Braude and F. C. Nachod Academic Press Inc. iic' York, U. Y.,
1955, p,,o 59h.
















3-Pic


O 2-Etpy


O 2-Pic


2,6-Col

O


dHB -
Figure 30. Steric requirement
of methanesulfonic acid, for


1.6 2.1
AH kcal./mole
of transition state relative to that
the pyridine bases.


1.6


0)

-4
0


C)


0.0


-1.6







h-Pic


3-Pic


O 2-Etpy


O 2-Pic


D 2,6-Col


I I


-.81 o.o0
dHB Hpy kcal./mole
Figure 31. Steric requirement of transition state relative to that
of a borane group, for the pyridine bases.


.8


1.6


0.0


-1.6
'"


i







4-Pic


O 2-Etpy


02-Pic


O 2,6-Col
I


-6.0'
SHB
Figure 32. Steric
to that of boron


-4.o -2.0 0.0
- A Hpy kcal./mole
requirement of transition state relative
trifluoride, for the pyridine bases.


-1.6


| I


|


~
X

I











than the solvolysis transition state will give plots with the ortho..

substituted pyridines lying above the line formed by yrridine,

3-r'icoliic, and 4-picoline.. INeglecting specific sterio interactions,

a statistical analysis (sec statistical appendix) of figure 31 for a

linear relationship between the Arrhcnius activation energies and the

heats of association of all the pyridines with diborane yielded a

correlation coefficient of 0.83. iTis value was statistically signi-

ficant at the 0.05 level, but not significant at the 0.01 level. When

steric interactions are considered, however, it is seen by inspection

of figure 31 that the correlation is much better and 1iore closely

related to chemical behavior if the results are grouped into steri-

cally similar compounds. Thus pyridine, 3-picoline and 4-picoline

would form one group, while 2-picoline, 2-ethylpyridine, and

2,6-lutidine would form another.

On this basis it is concluded that the transition state for

the solvolysis reaction has a smaller steric requirement than the

association reaction of the pyridines with dioborane. Since the

equilibrium

B2116 = 2IH3
is a comraon factor in the association of the pyridines with diborane,

it is also cr --,lu.cld that the steric requirem-ent for the solvolysis

transition state is less than that for the formation of amine-borane

from the awine and a borane group. This is of course in agreement

with the proposed solvolysis mechlaniir, as a stretching of the N-B

bond in a dissociation would relieve steric repulsions between BH3

and an ortho substituent.











Returning then to figure 31, if the line formed by pyridine,

3-picoline, and 4-picoline is taken as the relationship between these

two reactions in the absence of steric effects, then deviations from

this line may be taken as an estimate of relative steric strain. Thus

transition-state steric strains for the solvolysis reaction were

estimated to be about 0.8 kcal./mole for 2-picoline-borane, and 1.1

kcal./mole for both 2-ethylpyridine-borane and 2,6-1utidine-borane.

These values may be compared with the respective ground state strains

estimated by Brown and Domash,12 1.3, 1.5, and 2.7 kcal./mole.

If a value of 28.5 kcal./mole23 is taken for the gas-phase

enthalpy of dissociation of diborane into 2 borane groups, and if

it is also assumed that heats of reaction in nitrobenzene do not

differ significantly from those in the gas-phase, the gas-phase

enthalpy change for the reaction

CoSH NBH3 -C5H + BH3 (16)
may be estimated to be 32.2 kcal./mole. While it would thus appear that

the Arrhenius activation energies were far below the energetic require-

ments of the proposed mechanism, it should be remembered that the 32.2

kcal./mole value was for the gas-phase, where there are no solvent

effects. A strong coordination of either or both of,the products of (16)

relative to that of the ai.dne-borane would supply sufficient energy to

overcome the apparent deficiencies. Ihe situation here is quite similar

to the simple reaction of dissolving an ionic solid in water, for which

reaction the energy required to overcome the ionic lattice is supplied

by coordination of the ions with water. A very strong coordination of

the boron atom of the electron-deficient borane group with the oxygen











atom of an alcohol molecule will thus be proposed not only to account

for the aparent energy differences, but also to explain the negligible

rate of the reverse of equation (16) in propanol solution.

Evidence for this tight coordination comes from the reactions

of diborane with alcohols. With methanol, for example, more than enough

coordination energy is available to dissociate the diborane even in the

gas-phase.24 If a limited amount of alcohol is used, intermediates

such as (CH30BH2)x and (CH30)2BH may actually be obtained. The avail.

ability of a coordination energy of this magnitude for a borane group

in propanol solution would thus be more than sufficient to make the

proposed solvolysis mechanism energetically feasible. Concomitantly,
once a borane group were tightly coordinated, the association energy

released by the formation of amine-borane would not be sufficient to
overcome the coordination energy, and hence the rate of the associa-

tion reaction would be predicted to be insignificant with respect to

the dissociation reaction.

Hawthorne and Lewis25 have proposed that the hydrolysis of

pyridine diphenylborane involves the electrophilic attack of a water

proton on the electrons of the B-H bond. Under the same conditions

the hydrolysis of pyridine phenylborane was too slow to measure accu-

rately. In view of the considerable steric and electronic differences

between pyridine diphenylborane and the pyridine-boranes, it is thus

not unreasonable for diverse solvolysis mechanisms to hold.

As final support for the solvolysis mechanism proposed in

this dissertation, Davis and Kirby26 have recently reported a
negligible isotope effect, k/kD 1.05+0.02, for the reaction of








8h

trimethylamine-borane in 1.401~ hydrochloric acid. Since this reaction

has also been found to be first order in amine-borane,15 a dissociation

into the amine and a borne Zroup similar to that for the propanol

solvolysis is indicated for the protolysis reaction. It is indeed

gratifying to report that the protolysis reaction for pyridine-borane

has been found to be very much faster than that for trimethylamine-

borane under the same conditions.27












STATISTICAL AiTE:DIX


The least squares "normal" equations are well known and have

been used extensively in the physical sciences to fit a straight line
to experimental data. How to obtain standard errors for sample esti-

mates of the parameters in the least squares equation, on the other

hand, is little known and not always to be found in mathematical hand-

books. The calculation of these quantities, the calculation of a cor-

relation coefficient, and a statistical test for the homogeneity of

several least squares slopes will be given in this appendix.28 29

Both the mathematical expressions and the arithmetical calcula-

tions will be considerably simplified if the following symbols are

defined:

S?2 _Y; (E y)2/n
x2 FX2 ( x)2/n

Sxy where n is the number of points to be fitted by the least squares

line, and X and Y refer respectively to their Cartesian coordinates.
If the equation of the line to be determined is written as

Y' + a + hb
it can be shown that

b =fxy/'x2 and a *= Y bX

where X and Y are the corresponding averages of the X and Y variables.

When the individual deviations of the points are of no particular

interest, the sum of squares of deviations from the least squares line

85










may be obtained from
2- y2 )2/ 2.

Since a and b are estimators, computed from the sample, two degrees of

freedom have been lost, and the mean square deviation from the line

is given by

1.x d.d/(n-2).
The standard errors for the slope and intercept are respec-

tively determined from

sb 0 sy.X/f ) 2 and

sa = sY.x(/n + X21 )
The preceding equations have been derived from the assumptions that

the X values are measured without error, and that for each X, the Y

values will be normally and independently distributed with the same

variance. They are however, often applied even though the X's are

not exact if the error in X is negligible with respect to that of Y.

Data with a considerable amount of scatter may be tested for

a linear relationship by computing the correlation coefficient.

The correlation coefficient is a measure of the fraction of the total

variation explainable by a linear relationship between the variables.

It is defined by

r g371y/(f x2y2)t
The calculated value of the correlation coefficient is then examined

for statistical significance.

Pour kinetic experiments made with h-picoline-borane at 620C.

gave least squares slopes of -0,000827, -0.000837, -0.0008~1, and

-0.000926 (min.)"l. While the last slope would sceea inordinately








87

high, there was no justification for arbitrarily discarding it. Con-

cequcntly a statistical test based on the hypothesis that the slopes

were identical was applied to the data from these experiments, and the

hypothesis was rejected. Elimination of the -0.000926 (min.)-" value

resulted in acceptance of the hypothesis that the slopes were identi-

cal when the test was repeated. Hence the fourth slope was not used

in determining the avera e slope for this temperature. The test which

resulted in rejection is illustrated in table 14.









TEST FOR :0C.. MF!ETTY.


TABEI 14

OF SEVERAL


LAST SQUSTTArr SiLPES


Experiment Degrees M22 :w2 $272 2 Degrees Mean
of -( xy) /x of Square
Freedom Freedom

1 6 25200 -20.850 .017258 .000007 5

2 6 68336 -57.221 .47964 .000050 5

3 6 33429 -28.433 .02W206 .000022 5
S?5 39000 -.36.100 .033h59 .000043 4

23 165965 -142.60o .122887 .000122 19 0.00000642
Totals
.000356 22

Difference for testing Ho b, = b2 b3 = b4 .000234 3 0.0000780

3 0.0000780 12.15
19 0.00000642

For the 0.05 level, F3 is 3.13 therefore reject Ho
19












SFrT.': CS


1. A. !. 7. ur: and HI, I. ,c-lQcoin cr, J. An. Smhc:. soc., 29, 70o (1937)
2, S :-. auer, J. A. fI:cn. Soc1.02, 9 lOO1 (1937)
3. S, Gellor, 1, E. Hushies, and J, L. Holard, Acta. Cr st. 3 0 (1.1)
4h* 4. C, Price, R. D, Be Fraser, To. 'obinson, and H, C, Longuet.
uilc;lins3. ice. 1araday Soc., 1950 ('9 1231.
5. Bi. mico, n, J, Galiano, and 9,' J. Lchm.-~n, J, Ph:4s. Ch i, 61,
1222 (1957)
6. A, RIt :atrit53::.-, J* Chem, SOC., 2049 (1959)
7. C*. ;, *ax., A. R. .atritzcrl, and L. E.* utton, J. Chem. Soc, 1258
(1?58)
6, R. E. I'cCny and S. 1. lDauer, J. An. Chec. Soc., 70, 2061 (19?6)

9. UT 1, r'hill1ip, I. C, Irller, and '"-, L, t'uottrties, J. Am Chli. .
So 8c. 1, 81 96 (1959)
10. H. H. Jones, J. Am. Chl. Soc., 82, 2528 (1960)
U11 HI C. 3rown, H. I, SchlcinAcr, and 1. 2. Cordon, J. A.7 Cho. *oc.
66 325 (1942)
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77, 106 (1955)
15. G B. WE. aschiritsch, J. A i, Che. Soc., 02, 3290 (1960)
16, ve would like to tian.. the Callry C rheIical Copany for sawplca of
^'rldinc-borano and 2,L.l1utidine-borane.
17 1.. Jrnsen, !. Little, and U, ltruc,:, Anal. Chin., 24, l8h3 (i.52)
18. Ii. C. irown ind P. A. Tiernf, J. Anm. Chem. Soc., 00, 1552 (156)










19a. ;inctica an ::ccnnionr by A. A* Front andi RG* O Pearson, John lilcy
aid cons, njc.3, :c:u Iori, Y. (1953) p. 130.
19b. Ibid., p. 95.
20. o. Ai:,rlo", J. An. C:5n. Soc,, 2, 125 (1932)
21. A. B, ur rc..orc of Chrcical Pro-reso, .35 / 159 (19h)
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5387 (1956)
23, S. H. Bauer, J. Am. ChQa. Soc.. 78, 5575 (1956)
2h* A. B, DBurg and II. 1. cSchLesin.c-r, J. A.i. Cho:.. Soc.* $*, 4020
(1933)
25* ':. F, i.aut-h2or:e anid nr, S, 1".D, J, C;c-., coo., G., 4296
(193;)
26, R, F' Pavis and, C. L. Kirby, J. Ar. C;]n. Soc., 82, 5950 (196n)

27, G0 E, TrschlewIitsch, unpublished woirk
28. Statistics in Research by Bernard Ostle, Iowa State College Press,
Amcs, Iowa (195)1 p. 133-138.
29. Statistical Methods by George W. Snedecor, Iowa State College Press,
Ames, Iowa (1956) p. 122-159.














BIOGRAPHY


Ernest Rodman Birnbaur was born on October 4, 1933, in Newark,

New Jersey. His undergraduate work was started at U.C.L.A. and com-

pleted at the University of California in June, 1955 with the attain-

ment of a Bachelor of Arts Degree in Chemistry.

He began his graduate study under Dr. Anton Burg at the

University of Southern California where he was a teaching assistant

for two years. Following graduation from tlis institution with a

Master of Science degree in Chemistry he came to the University of

Florida in August of 1958 to pursue work leading toward the degree

of Doctor of Philosophy.

While at the University of Florida he has held a research

fellowship. He is a member of Kappa Nu, the American Chemical Society,

and the Chemical Society of England.










Thio dis sert.tion o~s prepared runcr tho direction of t.he
chairman of the ccndidato's .up:crisory coii,.tte.. and has been approved

by all .Tebers of that co~iittee. It was shbteitted to 'the Poan of the
CollcLc of Arts aznd Ecicnccs cnd tIo the Gracuhate Covncil, and was
ali- ro.'(_ as partial fulfllT'reInt of the rrc-irecen- for the degree of
Doctor of Fhilosophy,


January, 1961



Dean, College of Arts and Sciences


Dean, Graduate School


Supervisory Cor mittee:



fl/-Ai


L2 jZ-^










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AUTHOR: Birnbaum, Ernest
TITLE: Solvolysis kinetics ofpyridine-boranes. (record number: 424005)
PUBLICATION DATE: 1961


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