Title: Solvolysis kinetics of the methylamineboranes
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Title: Solvolysis kinetics of the methylamineboranes
Physical Description: vii, 111 l. : illus. ; 28 cm.
Language: English
Creator: Levy, Newton, 1935-
Publication Date: 1964
Copyright Date: 1964
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Subject: Dynamics   ( lcsh )
Boranes   ( lcsh )
Chemistry thesis Ph. D
Dissertations, Academic -- Chemistry -- UF
Genre: bibliography   ( marcgt )
non-fiction   ( marcgt )
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Thesis: Thesis - University of Florida.
Bibliography: Bibliography: l. 107-110.
General Note: Manuscript copy.
General Note: Vita.
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Bibliographic ID: UF00098226
Volume ID: VID00001
Source Institution: University of Florida
Holding Location: University of Florida
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Resource Identifier: alephbibnum - 000427114
oclc - 11084936
notis - ACH5856

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SOLVOLYSIS KINETICS OF THE

METHYLAMINEBORANES














By
NEWTON LEVY, JR.











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












UNIVERSITY OF FLORIDA


August, 1964














ACKNOWLEDGMENTS


The author wishes to express his sincere appreciation to Dr. G.

E. Ryschkewitsch, Chairman of the author's Supervisory Committee for

his guidance. His ideas, encouragement, and genuine concern have been

vital factors in the preparation and completion of this work.

The author wishes to thank his wife, Barbara, for her faith and

understanding. The completion of this work is due in no small part

to her constant encouragement and infinite patience.

To his parents, the author is indebted for their unselfish help

and sacrifices during his entire education.

To all of the supporting departments, and especially to Mr. Morris

Mixson and Mr. Richard Logsdon, the author is indebted for their help

in obtaining chemicals and apparatus.

The author also wishes to thank the Petroleum Research Fund,

whose financial support helped make this research possible; and the

faculty and staff of the Chemistry Department, who have spent valuable

hours teaching and assisting the author.















TABLE OF CONTENTS


ACKNOWLEDGMENTS

LIST OF TABLES

LIST OF FIGURES

CHAPTERS


I.

II.


III.

IV.


IINTODUCTION

EXPERIMENTAL

A. Materials

B. Procedures

EXPERIMENTAL RESULTS

DISCUSSION OF RESULTS


A. Aqueous kinetics

B. Kinetics in the 50 per cent 1-propanol-
water u.i;xed solvent

C. The distribution coefficient of mono-
methylamineborane between water and
benleiie

D. The cryoscopy of iLnouethlyla~ -ineborane
in benzene

V. SUMMARY

BIBLIOGRAPHY

BIOGRAPHICAL SKETCH


iii


Page

ii

iv

vi














LIST OF TABLES


LIST OF SYMBOLS AND ABBREVIATIONS


Table

1

2

3

4

5

6

7

8

9

10

11

12


13 KINETIC


DATA FOR MMAB IN WATER AT 27.550C.

DATA FOR MMAB IN WATER AT 30.700C.

DATA FOR MMAB IN WATER AT 31.940C.

DATA FOR iMAB IN WATER AT 35.940C.

OF KINETIC DATA FOR MMAB IN WATER

DATA FOR DIIAB IN WATER AT 36.080C.

DATA FOR DMAB IN WATER AT 40.340C.

DATA FOR DIAB IN WATER AT 42.790C.

DATA FOR DMAB IN WATER AT 44.950C.

DATA FOR EMWB IN WATER AT 50.270C.

OF KINETIC DATA FOR DNAB IN WATER

DATA FOR TMAB IN 50 PER CENT 1-PROPANOL-


WATER AT 31.830C.

14 KINETIC DATA FOR TMAB IN 50
WATER AT 35.850C.

15 KINETIC DATA FOR TMAB IN 50
WATER AT 40.230C.

16 KINETIC DATA FOR TMAB IN 50
WATER AT 44.380C.

17 SUMMARY OF KINETIC DATA FOR
1- PROPANOL- WATER

18 KINETIC DATA FOR DIAB IN 50
WATER AT 6.50C.


PER CENIT 1-PROPANOL-


PER CENT 1-PROPANOL-


PER CENT 1-PROPANOL-


TMAB IN 50 PER CENT


PER CENT 1-PROPANOL-


Page

22


KINETIC

KINETIC

KINETIC

KINETIC

SUMMARY

KINETIC

KINETIC

KINETIC

KINETIC

KINETIC

SUMMARY









LIST OF TABLES Continued


Table Page

19 KINETIC DATA FOIR DFB Il 50 PM. CENfT 1-
PROPANOL-WATER AT 10.50C. 55

20 KINETIC DATA FOR DMAB IN 50 PER CENT 1-
PLC?^AO',-' CI:' AT 14.6C. 56

21 SUMMARY OF KINETIC DATA FOR DMAB IN 50
PER CENT 1-PROPANOL-WATER 58

22 THE ACID HYDROLYSIS OF THE METHYLAHINE-
BORANES 72

23 SELECTED PHYSICAL VALUES FOR THE METITYL-
AMIIIEBORANES AND TIE M4ETHYLANdINIUIl IONS 80

24 COMPARISON OF THE SECOND ORDER RATE CON-
STIANTS OF DI- AND TRIE~ETHYLAfImTEBOR1PUE IN
WATER AND IN 50 PER CENT 1-PROPANOL-WATER 82

25 SUMMARY OF THE EFFECTS PRODUCED BY VARYING
CONDITIONS IN THE ACID HYDROLYSIS OF TIAB AT
LOW CONCENTRATIONS IN 50 PER CENT 1-PROPANOL-
.-ATER 87

26 DISTRIBUTION COEFFICIENTS FOR 10MAB IN
BENZENE AND WATER 97

27 CRYOSCOPIC DATA FOR I1ONOMETHYLAMINEBORANE
IN BENZENE 101













LIST OF FIGURES


Figure Page

1 Cryoscopic Apparatus 17

2 Temperature Versus Time Plot for Monomethylamine-
borane in Benzene with Extrapolation to Correct
for Supercooling 18

3 pH Versus Time Plot for 1bnomethylamineborane
(IMAB) in Water 26

4 Concentration Versus Time Plot for Monomethyl-
amineborane (MIMB) in Water 27

5 Second Order Rate Plot for Honomethylamineborane
(MMAB) in Water 28

6 Arrhenius Plot for the Acid Hydrolysis of Mono-
methylamineborane in Water 32

7 pH Versus Time Plot for Dimethylamineborane
(DMAB) in Water 42

8 Concentration Versus Time Plot for Dimethylamine-
borane (DNAB) in Water 43

9 Second Order Rate Plot for DImethyiaminaborane
(DMAB) in Water 45

10 Arrhenius Plot for the Acid Hydrolysis of
Dimethylamineborane in Water 46

11 Second Order Rate Plot for Trimethylamineborane
(TMAB) in 50 Per Cent 1-Propanol-Water 49

12 Arrhcnius Plot for the Acid Hydrolysis of Tri-
methylamineborane in 50 Per Cent 1-Propanol-
Water 53

13 Second Order Rate Plot for Dimethylamineborane
(DFAB) in 50 Per Cent 1-Propanol-Water 57

14 Arrhenius Plot for the Acid Hydrolysis of
Dimethylamineborane in 50 Per Cent 1-Propanol-
Water 59










LIST OF FIGURES Continued


Figure


15 Solubility of lMthylamineborane (MMAB) in Water

16 Solubility of Dimethylamineborane (DMAB in Water

17 Solubility of Trimcthylamineborane (TMAB) in Watce

18 Solubility of Monomethylanineborane (MMAB) in
Benzene

19 Reaction Path in the Aqueous (Lower Curve) and in
the Gas Phase (Upper Curve)

20 Pseudo-First Order Rate Plot for Tricethylanine-
borane in 50 Per Cent 1-Propanol-Water for 0.02 IM
TMAB
1
21 Plot of Log Distribution Coefficient Versus r for
SMAB in Benzene and Water

22 Plot of i Values Versus Molality for MMAB in
Benzene


vii


Page

60

61

62


63


76



88


98


102


r













CHAPTER I

INTRODUCTION


By virtue of an available empty p-orbital on the boron atom, tri-

covalent boron compounds can accept an electron pair from various donor

species. Addition compounds between boranes and a variety of Lewis

bases have been know for many years, but there is little known about the

change in activity with the change in the donor species.

Perhaps the best characterized of these adducts are the aminc-

boranes, and there are several general review articles on the B-N

bond.6,25,28,63 The first methylamineborane to be characterized was tri-

methylamineborane, which was synthesized in 1937 by Burg and Schlesinger.13

The compound was originally prepared by a displacement reaction between

borine carbonyl and the amine, and subsequently by the direct reaction of

the amine and diborane. Trimethylamineborane has also been prepared

from hydrogen and trimethylaminetrialkylboranes.41 The variety of

melting points assigned to monomethylamineborane was attributed to small

amounts of impurity incurred in the preparation, but Parry et al.50

prepared the pure compound by condensing the amine onto a diborane-

tetrahydrofuran solution at -75C. Mothylamineboranes have also been pre-

pared from the reaction of trimethylamine borontrifluoride and lithium

borohydride, from triphenoxyborate, aluminum metal, and hydrogen; and

from the electrolysis of sodium borohydride in liquid armorLia and in

liquid amines.64 However, the most convenient route to the methylamine-

boranes results from the reaction of the methylamine hydrochlorides with





2


sodium or lithium borohydride in diethyl ether or diglyne42246'60

according to the overall equation


R2CH3N.HC1 + MBI4 -- 20 > R2CII31UBH3 + H2 + MC1.


All of the aothylanrincborcnce are white crystalline solidC and

chou an increasing resistance to water hydrolysis from monomethylamine-

borane to trimethylamineborane. Mono-, di-, and trimethylamineborane

melt at 56, 370, and 940C.,46 respectively. Trimethylamineborane can

be heated for several hours above 1000C. without a detectable change in

its physical properties. However, dimcthylamineborane forms the amino-

borane (RCII31BII2) when heated,12 nnd monouethylarineborane can eventually

forrl N ,H,N-trimethylbora ine.68

Somc of the physical-chelical properties of the ncthylemincborancs

Kh' ch have been detrcnincd include the dipole ronents;4,46,50 electron

d.iraction3 and ;:-ray patterns;26 molecular weight necaurements in

liquid armonia,49 benzene, water, and dio:.ne;46 vapor pressure reacure-

monts and ha-atc o. vaporization;50 heats of formnation from the gaseous

mines and diborane;44 B11 chemical shifts53 and other nuclear mragnctic

resonance studies;23'50 and infrared spectra, Paman spectra, and B-N
31,32,38,50,55,56
bond force ccnstants.3132 505556

All of the methylamincboranes react with halogen acids, except HF,

to form the monohalogen c-ineborane according to the overall equation


R2CH3NBI13 + HX --- R2CH3NBH2X + H2,


where X = C1, Br, or I, and R = CH3 or H.48 The methylamineboranes can

be used to prepare lithium borohydride by reaction with lithiun hydride.41

The only methylamineborane which has been utilized for hydroboration is









trimethylamineborane,1,37 and an extensive study of the hydroborating

powers of amineboranes has not been undertaken. The borohydride ion

and boranes will reduce the iodate ion in aqueous solutions, and this

reaction has been made the basis of an analytical procedure to determine

the concentrations of amineborane solutions.42 Monomethylamineborane

and other monoalkylamineboranes have been widely used to prepare trialkyl-

N-substituted borazines.5'22,60,66

Relatively little research has been done, however, on the solution

kinetics of the methylamineboranes. Kelly, Marchelli, and Giusto39

calculated the rate constants for the acid hydrolysis of several amine-

boranes at room temperature but did not determine the activation energies.

Ryschkewitsch57 made an extensive study of the acid hydrolysis of tri-

methylamineborane, in which the author determined the activation energy,

the ionic strength dependence, and a logical mechanism for the reaction.

A report of some of the results of this work has appeared in the litera-

ture.30 In a study of the hydrolysis by DC1, Davis et al.17 found a

rapid and quantitative exchange of the boron hydrogens in trimethylamine-

borane with D20. A large amount of reserach has also been done on the

aqueous hydrolysis of sodium borohydride.9,16,18,19,20,36,52,62 Kinetics

in the gas phase have been studied for the combination of amines and

borontrifluoride,10,24,27,40 but the aqueous hydrolysis has received

relatively little attention.59 The gas phase kinetic studies of other

B-N compounds include Brumberger's11 investigation of the relative

reaction rates of diborane with mono- and trimethylamine, and a de-

tailed work29 on the mechanism of the reaction between diborane and tri-

methylamine.









Intuitively, one could propose that the solvolysis kinetics of

the methylamineboranes involve a dissociation of the adduct into an

amine and a borane fragment, attack of an acidic species on a B-H bond,

or displacement of an amine from the adduct molecule by the attacking

species. As previously noted,solvolysis mechanisms have been proposed,

and one of the purposes of this investigation was to test these

mechanisms, and possibly elucidate in some detail, a mechanism for the

hydrolysis of the methylamineboranes. At the same time, the effect of

substituting a methyl group for a hydrogen on the nitrogen atom of the

adduct could be observed from a study of the activation energies. It

was also proposed to determine the heats of solution of the methylamine-

boranes in order to calculate the heats of hydration of the compounds.

It was thought that a correlation between the trend in activation energies

and heats of hydration would be evident. In the hydrolysis of diborane,

intermediates of the type H3B.H20 have been proposed,61'67 and this inter-

mediate was also suspected to be present in an acid hydrolysis of the

methylamineboranes. However, the reaction of diborane with water is very

rapid and did not lend itself to study by our analytical methods. It was

therefore proposed to undertake a study of the acid hydrolysis of tri-

methylamineborane in 1-propanol-water solutions to see if intermediates

of the type R3NBH2(OC3H7) and R3NBH(OC3H7)2 could be detected, and to

extend the work of Ryschkewitsch and Birnbaum58 on the kinetics of sub-

stituted pyridineboranes in 1-propanol and water. We also wanted to

observe the differences in activation energies in the mixed solvent and

in water. It was thought that from the observed trends in activation

energies in these two media, we could predict the effects of various





5


solvent changes on the activation energies. Finally, a brief study of

the association of monomethylamineborane in benzene was undertaken in

order to investigate a possible error in the literature.














CHAPTER II

EXPERIMENTAL


A. Materials

Cylinder gases. All of the cylinder gases used were obtained from

the Matheson Company.

Benzene. Reagent grade benzene was distilled from CaH2 and stored

over "dry-Na" or CaH2.

1-Propanol. In order to rid the alcohol of any impurities which

might react with the amineboranes, it was treated with trimethylamineborane

before distillation. Reagent grade 1-propanol was made approximately

0.02 M in trimethylamineborane and approximately 1 M in HC1, and this

solution was stored until gas evolution ceased--about 24 hours. The solu-

tion was then neutralized with NaOHl, and enough distilled water was added

so that the 1-propanol to water ratio did not exceed four. The solution

was then distilled, and the fraction boiling at 87.70C. was collected.

The fraction collected at this temperature was the 1-propanol-water azeo-

trope, containing 28.9 per cent water and 71.7 per cent 1-propanol by

weight.43

The alcohol was also distilled from reagent grade 1-propanol which

was pre-dried with CaSO4. The fraction collected boiled at 97-98C.

Trimethylamineborane. Samples of trimechylamineborane from Callery

Chemical Company were sublimed at 0C. onto a cold finger at -780C. The

pure compound melts at 940C.,46 and the purity exceeded 98 per cent as

determined by reducing equivalents.









Dimethylamineborane. Samples of dlmethylamineborane from Callery

Chemical Company and from Chemical Procurement Laboratories were re-

crystallized from n-hexane and diethyl ether. The pure compound melts

at 370C.,46 and the purity exceeded 95 per cent as determined by re-

ducing equivalents.

Monomethylamineborane. Several methods of preparing monomethyl-

amineborane were attempted. No product could be obtained by bubbling

diborane through methylamine at -780C. The method employed by Parry et

al.50 of condensing the amine onto a solution containing diborane and

tetrahydrofuran at -780C. proved inadequate for the preparation of 1 to

10 gm. lots. Larger, pure samples were obtained by a modified method of

Noth and Beyer,46 using lithium borohydride and monomethylamine hydro-

chloride.

To prepare the hydrochloride, a 500 ml. three-necked round bottom

flask, containing 300 ml. of diethyl ether and equipped with a teflon-

coated magnetic stirring bar, was surrounded by an ice bath. The flask

was fitted with two g&i delivery tubes which extended below the surface

of the ether, and with a simple mercury bubbler to observe any pressure

differential. A cylinder of anhydrous HC1 was connected to one of the

delivery tubes through a trap containing concentrated sulfuric acid. A

cylinder of anhydrous monomethylamine was connected to the other delivery

tube through a mercury bubbler and a trap containing solid potassium

hydroxide. Stirring was initiated, and the gases were admitted, adjusting

their flow rates so that no pressure differential existed. The product

formed immediately, and the reaction was continued until, by visible in-

spection, the desired amount had been prepared. The product was allowed

to settle, and most of the ether was decanted; the remaining ether was










removed on the vacuum line by heating the product at 100C. for two hours.

The purity of the monomethylamine hydrochloride was determined by the

adsorption-indicator method for chloride ion54 and exceeded 99 per cent.

After removal from the vacuum line, the product was stored in an inert

atmosphere box. Samples could be exposed to the atmosphere for only short

periods of time before water absorption affected the weight.

To prepare the monomethylamineborane, an approximately 0.4 M solution

of lithium borohydride in diethyl ether was prepared from reagent grade

ether and 85 per cent lithium borohydride obtained from Metal Hydrides,

Inc. The solution was prepared in the dry box and allowed to stand until

gas evolution ceased. Approximately 250 ml. of the solution was decanted

into a flask fitted with an airtight rubber syringe stopper. Aliquots

were removed by means of a graduated hypodermic syringe, and the nor-

mality Kas determined by the method of Lyttle, Jensen, and Struck,42

using arsenite instead of thiosulfate to titrate the iodine.

In a typical run, 6.5 gm. (97 rmoles) of monomethylamine hydro-

chloride were added to a 250 ml. two-necked round bottom flask fitted

with a teflon coated magnetic stirring bar. This operation was carried

out in the dry box. The flask was placed on the vacuurw line, and approx-

imately 75 ml. of ether were distilled from a lithium borohydride-ether

solution onto the monomethylamine hydrochloride at -780C. The flask was

then surrounded by an ice bath and brought to atmospheric pressure with

dry nitrogen. A rubber syringe stopper nnd a mercury bubbler were fitted

to the two necks of the flask, and a magnetic stirrer was placed beneath

the flask and the ice bath. A graduated hypodermic syringe was used to

add 120 ml. of 0.36 1 lithium borohydride (43 mmoles) in ether, and









stirring was initiated. Reaction occurred according to the overall equa-

tion


LiBH4 + CH3NH2.HC1 CH3{H2BH3 + LiCI + H2.


The gas evolution, as indicated by the mercury bubbler, ceased after

approximately two hours. The cold solution was then filtered in the

atmosphere, the filtrate placed on the vacuum line, and the ether removed

at 0C. The product obtained had a yellowish tint. The impure compound

was dissolved in 10 to 15 ml. of benzene and warmed to 400C. An equal

volume of n-heptane was added, and the solution was cooled in an ice

bath. This solution was then filtered, and a pure, white crystalline

solid was obtained. The product was placed on the vacuum line and was

pumped on for several hours at 0C. It was then sublimed at 250C. onto a

cold finger at -780C. The pure product melted at 55-56oc., in agree-

ment with the literature value.46 The purity was checked by adding a

weighed sample to an excess of iodate, acidifying, and titrating the

excess oxidizing equivalents present as iodine with arsenite to a starch

end point. The following set of equations describe this analytical

method:


CH3NH2BH3 + 103" B(OH)3 + CI3NH2 + I"


6H + 103- + 51 -> 312 + 3H20


12 + As033" + H20 AsO4" + 2H+ + 21".


An alternate analysis was performed by hydrolyzing a weighed sample with

a known amount of excess acid and back titrating with sodium hydroxide.

The hydrolysis proceeds according to the equation:









CH3NH2BH3 + 3H20 +- -- 3H2 + B(OH)3 + CH3IH3


The purity as calculated by both methods exceeded 98 per cent. These

results and the findings of a commercial analysis are listed below.

B-IH as re- Acid
7%C %H oN during eq./g. eq./P,.

Found: 27.03 18.17 30.99 \ 0.193 0.190

Calculated: 26.75 17.96 31.20 0.193 0.193

The yield of pure product was 56 per cent.

B. Procedures

Calibration of the pH apparatus. The pH was followed by a Model SR

Sargent Recorder equipped with a resistance-matching adapter, S-72172

Sargent pH adapter, used in conjunction with Beckman glass and calomel

electrodes. In the preliminary experiments, the electrodes behaved

erratically when placed into the reactant solutions. Several steps had

to be taken to produce consistent results. The electrodes were thermo-

stated in the constant temperature bath in order to insure rapid tempera-

ture equilibration. Enough solid potassium chloride was added to the

calomel electrode, so that a saturated solution was maintained over the

temperature range studied. Further stabilization was attained by con-

structing a barrier junction to fit over the calomel electrode. The

junction consisted of a tube with a capillary to allow solution contact

while preventing dilution and contamination in the reference cell. Before

each run the barrier was filled with a saturated potassium chloride

solution, which was also thermostated at all times in the constant tem-

perature bath. The whole apparatus was recalibrated with prepared buffer

solutions daily and before runs at a new temperature. The recorder speed









was known, and a pH versus time plot was obtained directly from the

recorder chart paper.

The acid hydrolysis of dimethylamineborane in water. The hydrolysis

proceeds according to the overall equation


(CH3)2NHBH3 + HI + 31120 > B(O0H)3 + (CH3)2NH2 + 3H2.


The progress of the reaction can be followed by determining the loss of

reducing power of the solution or the change in the hydrogen ion concen-

tration. Trimethylamineborane solutions have been analyzed by the former

method by Ryschkewitsch.57 In this investigation a weighed amount of di-

methylamineborane was added to a known volume and concentration of

hydrochloric acid. The pH and time were read directly from the recorder

plots. The instantaneous dimethylamineborane concentration could then be

determined from the stoichiometry of the reaction.

The reaction was studied over the temperature range of 36-500C.

The initial dimethylamineborane concentration varied from 0.008 to

0.0275 M, and the initial hydrochloric acid concentration varied from

0.004 to 0.015 M. All of the reactions were run at a constant ionic

strength of 0.10 M by adding appropriate amounts of potassium chloride.

All solutions used in the kinetic runs were thermostated in the constant

temperature bath to + 0.010C. The concentrations were adjusted so that

the half-life of the component in the lower concentration was attained in

10-15 minutes. There were two half-lives in the longest run, and the

usual run was halted after one half-life.

An excellent fit of the data to the rate equation


-d[dimethylamincborane]/dt = k [dimethylamineborane] [I0]









was obtained over the range of half-life, temperature, and solvent com-

position studied. The second order rate constants, k, were determined

from the slope of the straight line relation of

log [dimethylamineborane]/ [(] and the time. The method of least squares

was used to optimize the fit. The activation energy was then determined

from the Arrhenius equation, In k = A(exp)-ai4E/RT, where AF is the

activation energy, A is the pre-exponential factor, k is the second order

rate constant, and T is the absolute temperature. The activation energy

was calculated by a least squares treatment of log k versus 1/T.

The decomposition of dinethylamineborane in pure water occurred at

a negligible rate. A detailed discussion of the proposed mechanism of

solvolysis appears in a later section.

The acid hydrolysis of monomethylamineborane in water. The acid

hydrolysis of monomethylamineborane in aqueous solutions proceeds to

completion analogous to the dimethylamineborane hydrolysis according to

the overall equation


CH3NH2BH3 + +tf 3H20 > B(OH)3 + CH3NIH3 + 3H2.


The method of analysis of the monomethylamineborane hydrolysis was the

same as for the dimethylamineborane hydrolysis.

The reaction was studied over the temperature range of 27.5 to 35.90C.

The initial monomethylamineborane concentration varied from 0.0025 to

0.0070 M, and the initial hydrochloric acid concentration varied from

0.0030 to 0.0062 M. The ionic strength was adjusted to 0.10 M with

potassium chloride. Each run covered at least one half-life o; the com-

ponent in the lower concentration.









Monomethylamineborane is more sensitive to moisture than the di-

methylamineborane, and consequently, samples of the former were stored

in the dry box. In a typical run, a sample of monomethylamineborane

was removed from the dry box, dissolved in 50 ml. of 0.10 M potassium

chloride, and the initial concentration of the amineborane was deter-

mined by measuring the reducing power of the solution by the iodate

method. A known volume and concentration of hydrochloric acid was then

added, and the initial concentrations of both components were determined

from dilution factors. No decomposition of monomethylamineborane in pure

water was detected over the approximately ten minute time interval

utilized in the measurements.

The data obtained were treated in the same manner as that in the

dimethylamineborane hydrolysis.

The acid hydrolysis of di- and trimethylamineborane in 1-propanol-

water solutions. The hydrolysis proceeds in the same manner as that in-

dicated for the preceding two amineboranes. The method of analysis

differed from that in pure water, in that the concentration of the amine-

borane was followed by measuring the loss in reducing power by the iodate

method.

The reaction with trimethylamineborane in the mixed 1-propanol-

water solvent was studied over the temperature range of 31.8 to 44.40C.

The initial trimethylamineborane concentration varied from 0.370 to

0.710 M, and the initial hydrochloric acid concentration varied from

0.818 to 1.02 M. The solvent contained 50.6 per cent of 1-propanol by

volume. The longest run covered approximately two half-lives, and the

usual run covered one half-life.









The reaction with dimethylamineborane in the mixed solvent was

studied over the temperature range of 4.5 to 14.60C. The initial di-

methylamineborane concentration varied from 0.425 to 0.454 M, and the

initial hydrochloric acid concentration was 0.818 M. The volume ratio

of 1-propanol to water was constant throughout the experiments. Each run

covered at least one half-life of the amineborane.

It was found that when the reaction was carried out in an open

vessel and with large reaction rates, the hydrogen gas evolved would

carry out considerable amounts of the solvent, and preferentially the

1-propanol. This proved to be a serious source of error. This effect

was minimized by running the reactions in a closed tube connected to a

mercury bubbler which maintained a slight positive pressure on the

solution, but which allowed the hydrogen to escape with much smaller

amounts of the solvent. At high concentrations of amineturane and hydro-

chloric acid, solvent evaporation restricted the reaction conditions to

a relatively narrow range of concentrations and temperatures.

In a typical run, a weighed amount of the amineborane was dissolved

in the 1-propanol-water azeotrope, which contained 75.85 per cent 1-

propanol by volume. The initial concentration of the amineborane was

found by measuring the reducing power of the solution by the iodate

method. A known amount of this solution was then pipetted into the re-

action tube, and a known volume and concentration of hydrochloric acid

was added. The volume ratio of the azeotrope to the hydrochloric acid

solution was two, so that the volume ratio of 1-propanol to water in the

final solution was 1.02. Aliquots of the reacting solution were removed

at different time intervals, and the instantaneous concentration of the

amineborane was found by the iodate method. The instantaneous acid









concentration could then be determined from the stoichiometry of the

reaction.

No ionic strength adjustments were made. The solutions used in

the trimethylamineborane hydrolysis were thermostated to + 0.010C., and

those used in the dimethylamineborane hydrolysis were thermostated to

+ 0.10C. The data obtained were treated in the same manner as that of

the hydrolysis in pure water. Vigorous stirring of the solution before

removing the aliquots reduced the interference of the hydrogen gas

evolution with pipetting.

Attempts were made to study the trimethylamineborane reaction as

a pseudo-first order reaction, in which the amineborane concentration

was approximately 0.02 M, and the hydrochloric acid concentration was

1.0 M. Consistent and reproducible data could not be obtained, and the

behavior of these reactions will be discussed in a later section.

The distribution coefficient of monomethylamineborane between

benzene and water. A weighed amount of monomethylamineborane was dis-

solved in a known volume of water, and this solution was shaken with an

equal volume of benzene. All of the solutions were thermostated to

+ 0.10C. Twenty to thirty minutes were allowed for equilibrium to be

reached, and then aliquots were removed from both the organic and aqueous

phases and titrated by the iodate method.

Precision of the data was relatively poor, and the experiment was

abandoned in favor of direct solubility measurements in water. The

reasons for the initial undertaking of this experiment and for the poor

results are outlined in the discussion section.

The solubility of monomethylamineborane in benzene. The solubility

was studied over the temperature range of 5.9 to 21.50C. A saturated









solution with a large excess of solid was prepared at room temperature.

The temperature was held constant to + 0.10C. in a fJcwr flask. Aliquots

were pipetted out at several temperatures through a small glass fritted

filter tube, which was attached to the pipet by means of a lubricated

loop of tygon tubing. The concentration of the auineborane was then de-

termined by measuring the reducing power of an aliquot by the iodate

method. A straight line was obtained from a plot of log (monomethyl-

arineborane concentration) versus temperature, and the heat of solution

was calculated from the least squares slope of this line. The solubility

of the monomethylamineborane at the freezing point of the benzene solution

was determined by extrapolation from this plot, and this particular con-

centration was then utilized in calculations involved in the cryoscopic

study.

The cryoscopy of monomethylamineborane-benzene solutions. The

freezing point depressions in benzene were determined at various concen-

trations of monomethylamineborane. The apparatus consisted of a double-

walled, glass wool insulated vessel, a Beclkan thermometer, and a mag-

netic stirring mechanism operated by a relay timer, as shown in the

figure on the following page.

The thermometer was calibrated at the freezing point of pure benzene.

In order to correct for the effect of supercooling, the temperature ver-

sus time curve was extra:nolated through the supercooling region in all

instances. The intercept of the extrapolated line and the initial portion

of the curve was taken as the freezing point of the solution. Temperature

differences were measured to + 0.0050C.

In a typical run, 8-10 ml. of a monomethylanineborane-benzcne solu-

tion were placed in the freezing point apparatus. The concentration of






17













u

CI





0
o







4J4
Oo














-,4 -4 C) co

Q) 0 E-I 0
-i C E


I I I II I 11 Ii1 )













4o o
0 0
0 *d

(0

COM


0





r-4
























0
*,
i(-




C)
94

CO







18

3.32-







3.28







3.24-




00
(U




a, 3.20
Ca





> 3.16-
r4
4j
u-



C)



P 3.12-

ca





3.08-







3.04-







3.00
2 4 6 8 10

Time in Minutes

Fig. 2. Temperature vs. Time Plot for Monomethylamineborane in
Benzene with Extrapolation to Correct for Supercooling









the solution was then determined by the iodate method from the reducing

equivalents of an aliquot. The molality of the solution was then approx-

imated by dividing the molarity by the density of benzene at 250C. The

density of benzene at 250C. was calculated to be 0.874 g./ml. from a linear

interpolation of the densities at 200C. and 300C.21 The van't Hoff

factor, i, was determined from the ratio of the expected freezing point

depression, ATe, to the observed freezing point depression, ATo; i.e.,

i = ATe/ATo. The expected freezing point depression can be calculated

from the expression ATe = K morality, where K: is the cryoscopic con-

stant for benzene and has the value of 5.12.15 These i values were then

compared to those in the literature.46

The solubility of trimethylamineborane in water. A saturated,

aqueous solution of trimethylamineborane was prepared at room temperature.

Aliquots were removed over the temperature range of 0-33 C. by means of

a glass fritted filter tube attached to a 1 ml. pipet. The temperature

was held constant to + 0.10C. The reducing equivalents contained in the

aliquots were then determined by the iodate method. The heat of solution

was calculated from the least squares slope of a plot of log (trimethyl-

amineborane concentration) versus I/temperature.

The solubility of mono- and dimethylamineborane in water. Since

these compounds were vastly more soluble in water than was the trimethyl-

amineborane, a slightly different procedure was used in order to econo-

mize on reagents.

Saturated, aqueous solutions of the amineboranes were prepared at

temperatures below 200C. A 0.10 ml. capacity Hamilton microliter syringe

with a small styrofoam ball attached to the orifice of the needle tip






20


was used to remove 0.05 to 0.085 ml. aliquots from the solutions. Di-

methylamineborane was studied over the temperature range of 0-200C.,

and monomethylamineborane was studied over the temperature range of 0-

130C. The reducing equivalents contained in the aliquots were then

determined by the iodate method. The heats of solution of these amine-

boranes were determined in the same manner as the heat of solution of

trimethylamineborane.















CHAPTER III

EXPERIMENTAL RESULTS


The data compiled from the acid hydrolysis of the methylamine-

boranes are listed in Table 2 through Table 21. Following Table 3 for

monomethylamineborane in water, Table 11 for dimethylamineborane in

water, Table 14 for trimethylamineborane in 50 per cent l-propanol-water,

and Table 20 for dimethylamincborane in 50 per cent 1-propanol-water, are

plots indicative of the treatment given to each kinetic run. At the end

of the pertinent tables, a summary of the kinetic data for each amine-

borane hydrolysis is given, along with the Arrhenius plots and parameters.

The solubility curves for each of the methylamineboranes are pre-

sented at the end of the kinetic data.

The results of the cryoscopic work on monomethylanineborane in

benzene and the distribution coefficient data appear in the discussion

section.

Tables of comparison of data, or tables containing specific data

from the literature, are presented in the discussion to which they

refer. A table of the list of symbols and abbreviations used throughout

the thesis is presented on the following page.









TABLE 1

LIST CF SYMBOLS AND ABBREVIATIONS


Symbols Definitions


MMAB Monomethylamineborane

DMAB Dimethylamincborane

TMAB Trimethylamineborane

aTe Expected freezing point depression

ATo Experimentally observed freezing point depres-
sion

i van't Hoff factor

OK. Degrees Kelvin

AE* Activation energy in kcal./mole

T Absolute temperature

k Second order rate constant in liter/mole/time

t Time

P Ionic strength in moles/liter

[ ] Concentration in moles/liter

[ J] Initial concentration in moles/liter

M Molarity in moles/liter

m Molality in moles of solute/1000 g. of solvent

R Gas constant in cal./mole/deg.

A Defined by In k = A B/T

B Defined by In k = A B/T

Aki Heat of hydration in kcal./mole

AHs Heat of solution in kcal./molc









TABLE 1 Continued


Symbols


Definitions


AHV Heat of vaporization in kcal./mole

K An equilibrium constant

50% 1-propanol-water A solution containing 50.6% of 1-propanol by
volume












KINETIC DATA


TABLE 2

FOR MMAB IN


WATER AT 27.550C.


[] x 103 IMMAB] x 103 log eMMAB]/[1e] Time, min.


6.026

3.793

2.692

2.079

1.687


6.132

3.899

2.798

2.185

1.793

k =


5.129

3.970

3.153

2.584

2.165

k =


6.673

5.652

4.927

4.417

4.047

k =


6.012

4.853

4.036

3.467

3.048


0.0078

0.0120

0.0166

0.0216

0.0265

25.57 liter/mole/min.


-0.0690

-0.0872

-0.1072

-0.1277

-0.1576

26.33 liter/mole/min.


0.2074

0.2582

0.3135

0.3700

0.4270

25.20 liter/mole/min.


4.410

3.119

2.394

1.884

1.514


--








TABLE 3

KINETIC DATA FOR MIAB IN WATER AT 30.700C.


[H] x 103 24MAB] x 103 log [MMAB]/ [I] Time, rain.


6.000 7.896 0.1193 0

4.853 6.749 0.1433 1

3.990 5.886 0.1688 2

3.304 5.200 0.1970 3

2.786 4.682 0.2256 4

2.399 4.295 0.2529 5

2.070 3.966 0.2824 6

1.815 3.711 0.3107 7

1.603 3.499 0.3391 8

1.419 3.315 0.3685 9

1k 33.91 liter/mole/min.


4.083 6.830 0.2235 0

3.357 6.104 0.2596 1

2.780 5.527 0.2984 2

2.317 5.064 0.3397 3

1.954 4.701 0.3813 4

1.671 4.418 0.4223 5

1.429 4.176 0.4657 6

k = 33.98 liter/mole/min.








2.85





2.77


(4IMAo = 0.008 M
T = 30.700C.
i = 0.10 M


2.69





2.61






2.53





2.45





2.37






2.29





2.21
2 4 6 8
Time in Minutes

Fig. 3. pH vs. Time Plot for Monomethylamineborane (MMAB)
in Water








8.0


7.4


iHJ = 0.006 M
T = 30.700C.
p = 0.10 M


6.8






6.2






5.6
o\
0




L 5.0






4.4






3.8






3.2 ---
2 4 6 8

Time in Minutes

Fig. 4. Concentration vs. Time Plot for Monomethylpmineborane
(MMAB) in Water







0.37 r


0.34 --


0.31





0.28





0.25


CM1AB] = 0.008 M
C[H+o= 0.006 M
T = 30.700C.
p = 0.10 M


0.22 --


0.19





0.16




0.13


/I I I I I


Time in Minutes

Fig. 5. Second Order Rate Plot for Monomethylamineborane (1iAB)
in Water









TABLE 4

KINETIC DATA FOR MMAB IN WATER AT 31.940C.


[H+] x 103 [IMAB] x 103 log [In4AB]/ [1] Time, min.


3.140

2.673

2.286

1.954

1.698

1.493

1.318

1.159

1.040


4.971

4.504

4.117

3.785

3.529

3.324

3.149

2.990

2.871


0.1995

0.2266

0.2555

0.2871

0.3:77

0.3475

0.3782

0.4116

0.4411


k = 38.30 liter/mole/min.


0.1884

0.2127

0.2398

0.2691

0.2978

0.3265

0.3564

0.3853

0.4151

38,36 liter/mole/min.


3.162

2.716

2.328

2.000

1.742

1.531

1.349

1,202

1.072


4.878

4.432

4.044

3.716

3.458

3.247

3.065

2.918

2.788

k =










TABLE 5

KINETIC DATA FOR MAB IN WATER AT 35.940C.


[1+] 2 103 [MMAB] x 103 log [ MAB]//[Ir] Time, min.


6.166 2.799 -0.3430 0

5.495 2.128 -0.4119 1

5.012 1.645 -0.4839 2

4.624 1.257 -0.5659 3

4.364 0.997 -0.6411 4

4.140 0.773 -0.7289 5

3.990 0.623 -0.8066 6

3.855 0.488 -0.8976 7

3.767 0.400 -0.9739 8

3.698 0.331 -1.0482 9

k = 54.55 liter/mole/min.


4.169 2.506 -0.2210 0

3.715 2.052 -0.2577 1

3.365 1.702 -0.2960 2

3.076 1.413 -0.3378 3

2.851 1.188 -0.3802 4

2.692 1.029 -0.4177 5

2.547 0.884 -0.4595 6

2.432 0.769 -0.5000 7

2.328 0.665 -0.5441 8

2.259 0.605 -0.5722 9

k = 55.29 liter/mole/min.









TABLE 6

SU~IARY OF KINETIC DATA FOR


MMAB IN WATER


Standard
k, l.molc-lin."1 deviation log k 1/T, 103 x K.-l


25.70 0.47 1.410 3.326

33.95 0.05 1.531 3.292

38.33 0.04 1.584 3.278

54.92 0.05 1.740 3.236

AE* = 15.81 kcal./mole


Arrhcnius parameters from log k = A B/T: k in l.mole'lGec.1I

A = 11.14

B = 3,457








1.74





1.70





1.66





1.62






2 1.58
0




1.54





1.50





1.46





1.42
3.31 3.29 3.27 3.25 3.23


1 x 103 oK.-l
T

Fig. 6. Arrhenius Plot for the Acid Hydrolysis of Monomethylamine-
borane in Water











KINETIC DATA FOR


TABLE 7

DMAB IN WATER AT 36.080C.


[( 1] 102 [DMAB] x 102 log [DI-AB]/[i ( Time, min.


1.589 2.759 0.2396 0*

1.400 2.570 0.2639 2*

1.268 2.438 0.2840 4

1.183 2.353 0.2986 6

1.109 2.279 0.3128 8

1.038 2.208 0.3278 10

0.975 2.145 0.3424 12

0.918 2.088 0.3568 14

k = 1.439 liter/mole/min.


"'These points omitted from the least squares fit.












KINETIC DATA FOR


TABLE 8

DMAB IN WATER AT 40.340C.


[H+] x 103 [DMAB] x 102 log [DMAB]/ [1] Time, min.


9.772

8.974

8.260

7.621

7.063

6.546

6.109


1.903

1.823

1.752

1.688

1.632

1.580

1.537

k =


1.879

1.799

1.727

1.662

1.606

1.552

1.503

k =


1.873

1.797

1.728

1.664


0.2894

0.3077

0.3263

0.3452

0.3636

0.3827

0.4005

2.309 liter/mole/min.


0.2790

0.2960

0.3143

0.3326

0.3506

0.3698

0.3890

2.390 liter/mole/min.


0.2767

0.2934

0.3105

0.3282


9.885

9.099

8.375

7.727

7.161

6.622

6.138


9.908

9.141

8.453

7.816









TABLE 8 Continued


[IEr] 103 [DMAB] x 102 log [DMABJ/[He] Time, min.


7.228 1.605 0.3466 8

6.683 1.551 0.3655 10

6.180 1.500 0.3852 12

k = 2.366 liter/mole/min.












KINETIC DATA FOR


TABLE 9

DMAB IN WATER AT 42.790C.


[IJ] x 103 [DMAB] x 102 log [DIAB]/ [IM Time, min.


10.28

9.506

8.770

8.110

7.482

6.966

6.457




10.47

9.660

8.912

8.222

7.568

6.998

6.501


1.434

1.357

1.283

1.217

1.154

1.103

1.052

k =


1.457

1.376

1.301

1.232

1.167

1.110

1.060

k =


1.642

1.549

1.464

1.386


0.1446

0.1544

0.1652

0.1764

0.1884

0.1995

0.2119

3.244 liter/mole/min.


0.1436

0.1535

0.1644

0.1758

0.1881

0.2003

0.2124

3.289 liter/mole/min.


0.1903

0.2049

0.2206

0.2368


10.59

9.660

8.810

8.035








TABLE 9 Continued


[1&] x 103 [DMAB] x 102 log [DMAB]/ [H+] Time, min.


Y.345 1.317 0.2538 8

6.730 1.256 0.2709 10

6.209 1.204 0.2874 12

k = 3.250 liter/mole/min.









TABLE 10

KINETIC DATA FOR DMAB IN WATER AT 44.950C.


['] x 103 [mAB] 3i 103 log [DMAB]/[K1] Time, min.


4.027 9.897 0.3906 0

3.767 9.637 0.4079 2

3.507 9.377 0.4272 4

3.273 9.143 0.4461 6

3.048 8.928 0.4667 8

2.851 8.721 0.4856 10

2.673 8.543 0.5046 12

k = 3.771 liter/mole/min.


4.027 9.124 0.3553 0

3.758 8.855 0.3722 2

3.057 8.604 0.3897 4

3.311 8.408 0.4047 6

3.112 8.209 0.4213 8

2.917 8.014 0.4389 10

2.748 7.845 0.4556 12

k = 3.769 liter/mole/min.


4.027 8.058 0.3012 0

3.793 7.824 0.3145 2

3.597 7.628 0.3265 4

3.404 7.435 0.3390 6









TABLE 10 Continued


[1+] x 103 [DMAB] x 103 log [DMAB]/[H+] Time, min.


3.221 7.252 0.3524 8

3.041 7.072 0.3665 10

2.877 6.203 0.3804 12

k = 3.756 liter/mole/min.












KINETIC DATA FOR


TABLE 11

DMAB IN WATER AT 50.270C.


[(1] x 103 [DMAB] x 103 log [DIAB]/ [1+] Time, min.


4.808

4.345

3.954

3.614

3.296

3.013

2.754

2.547


8.403

7.940

7.549

7.209

6.891

6.608

6.349

6.142

k =


9.314

8.802

8.343

7.931

7.583

7.290

7.023

6.785

k =


8.027

7.591

7.184


0.2425

0.2617

0.2808

0.2999

0.3204

0.3410

0.3637

0.3822

6.419 liter/mole/min.


0.2711

0.2936

0.3172

0.3422

0.3699

0.3908

0.4155

0.4406

6.474 liter/mole/min.


0.2095

0.2253

0.2423


4.989

4.477

4.018

3.606

3.258

2.965

2.698

2.460


4.955

4.519

4.112








TABLE 11 Continued


[It] x 103 [DMAB] x 103 log [DIAB]/ [t] Time, min.


3.741 6.813 0.2603 6

3.444 6.516 0.2769 8

3.170 6.242 0.2942 10

2.917 5.989 0.3126 12

k = 6.459 liter/mole/min.






42
2.61

[DMAB]o = 0.01 M
T = 50.260C.
S= 0.10 M

2.57





2.53





2.49






2.45





2.41






2.37





2.33





2.29
3 6 9 12 15

Time in Minutes

Fig. 7. pH vs. Time Plot for Dimethylamineborane (DMAB)
in Water








9.7





9.3


C[Ho = 0.005 M
T = 50.260C.
p = 0.10 M


8.9






8.5





8.1


LiS


7.7





7.3





6.9





6.5
3 6 9 12 15

Time in Minutes

Fig. 8. Concentration vs. Time Plot for Dimethylamineborane (DMAB)
in Water










TABLE 12

SUMMARY OF KINETIC DATA FOR


DMAB IN WATER


Standard
k, 1.mole-lmin.-1 deviation log k 1/T, 103 x K.-1


1.439 0.1581 3.234

2.355 0.07 0.3700 3.190

3.261 0.03 0.5132 3.165

3.765 0.01 0.5758 3.143

6.451 0.04 0.8096 3.092

aE = 21.21 kcal./mole


Arrhenius parameters from log k = A B/T: k in l.mole'isec.-1

A = 13.38

B = 4,637








0.44 --


0.42 --


0.40 --


0.38 -


0.36





0.34


CDMAB~o = 0.01 M
CH'o = 0.005 M
T = 50.260C.
p = 0.10 M


0.32 --


0.30 --


d I I I 1 1
3 6 9 12 15


Time in Minutes

Fig. 9. Second Order Rate Plot for Dimethylamineborane (DMAB)
in Water


0
0
M-


0.28








0.84






0.75






0.66






0.57






0.48




0
o
0O
0.39






0.30






0.21






0.12
3.21 3.18 3.15 3.12 3.09

1 x 103 oK.-l
T

Fig. 10. Arrhenius Plot for the Acid Hydrolysis of Dimethylamine-
borane in Water









TABLE 13

KINETIC DATA FOR TMAB IN 50 PER CENT 1-PROPANOL-WATER
AT 31.83oC.


[l,] [TMAB] log []H+]/ AB] Time, min.


1.022 0.706 04161 0

0.930 0.614 0.180 60

0.840 0.524 0.205 120

0.744 0.428 0.240 255

0.705 0.389 0.258 315

0.699 0.353 0.278 375

0.647 0.331 0.291 426

k = 2.2 x 10-3 liter/mole/min.


1.022 0.574 0.250 0

0.946 0.498 0.279 60

0.879 0.431 0.309 117

0.811 0.363 0.349 210

0.772 0.324 0.377 267

0.741 0.293 0.403 325

0.711 0.263 0.432 382

k = 2.42 x 10-3 liter/mole/min.












KINETIC DATA FOR TMAB


TABLE 14

IN 50 PER CENT 1-PROPANOL-WATER AT 35.850C.


[f] [TMAB] log [H']/[TMAB] Time, min.


4.05 x 10-3


0.298

0.355

0.413

0.517

0.572

0.663

liter/mole/min.


0.514

0.402

0.320

0.222

0.186

0.141

k =


0.659

0.588

0.466

0.392

0.305

0.240

k =


4.20 x 10-3


0.340

0.400

liter/mole/min.


1.022

0.910

0.828

0.730

0.694

0.649


0.191

0.209

0.250

0.285


1.022

0.951

0.827

0.755

0.668

0.604


0

60

121

233

301

407


0

30

90

150

240

308






0.42






0.39


CTMABD, = 0.65 M
[HjO = 1.02 M
T = 40.230C.


0.36 -


0.33 -


0.30 -


0.27 -


0.24 -


0.21




0.18
0.18


120


240


300


Time in Minutes

Fig. 11. Second Order Rate Plot for Trimethylamineborane (TMAB)
in 50 Per Cent 1-Propanol-Water


0
P-4
0
O












KINETIC DATA FOR TMAB


TABLE 15

IN 50 PER CENT 1-PROPANOL-WATER AT 40.230C.


[( ([TMAB] log [ / [(TMAB] Time, min.


0.195

0.228

0.257

0.289

0.328

0.358

0.397

6.77 x 10-3 liter/mole/min.


1.022

0.914

0.825

0.760

0.696

0.657

0.616


0.653

0.541

0.456

0.391

0.327

0.288

0.247

k =


0.371

0.309

0.272

0.234

0.204

0.179

0.158

k =


0.583

liter/mole/min.


6.81 x 10-3


0.343

0.389

0.422

0.464

0.504

0.544


0

30

57

86

123

150

186


0

35

61

92

124

151

182


0.818

0.756

0.719

0.681

0.651

0.626

0.605








TABLE 16

KINETIC DATA FOR TAMB II 50 PER CENT 1-PROPANOL-WATER
AT 44.380C.


I[f] [TMAB] log [e(l/ [TMAB] Time, min.


0.818 0.422 0.287 0

0.719 0.323 0.348 31

0.659 0.236 0.399 58

0.613 0.217 0.451 84

0.568 0.172 0.519 120

0.540 0.144 0.574 151

0.522 0.126 0.617 174

k = 1.10 x 10-2 liter/mole/min.










TABLE 17

SUMMARY OF KINETIC DATA FOR ThAB IN 50 PER CENT 1-PROPANOL-WATER


1, 1.mole imi.-


Standard
deviation


log k


1/T, 103 x K.-l


2.31 x 10-3

4.13 x 10-3

6.79 x 10-3

1.10 x 10-2


0.11 -2.635

0.08 -2.384

0.02 -2.168

-1.959

AE+ = 23.4 kcal./mole


Arrhcnius parameters from log k A B/T. k in l.mole-1scc.-'

A = 12.30

B = 5,110


3.279

3.236

3.191

3.149








-2.7






-2.6






-2.5






-2.4






o -2.3






-2.2






-2.1






-2.0






-1.9
1.9 3.17 3.20 3.23 3.26 3.29

1 x 103 K.-l
T


Fig. 12. Arrhenius Plot for the Acid Hydrolysis of Trimethylamineborane
in 50 Per Cent 1-Propanol-Water









TABLE 18

KINETIC DATA FOR IDIAB IN 50 PER CENT 1-PROPANOL-WATER
AT 6.50C.


(Oi'] [DMAB] log [H+]/ [DMAB] Time, min.


0.818 0.425 0.284 0.0

0.769 0.376 0.311 20.0

0.729 0.336 0.336 41.5

0.685 0.292 0.370 63.5

0.649 0.256 0.404 85.0

0.595 0.202 0.469 122.5

k = 8.81 x 10-3 liter/mole/min.









TABLE 19


KINETIC DATA FOR DMAB IN 50 PER CENT 1-PROPANOL-WATER
AT 10.5C.


[I] [DMAB] log [I+ ]/[DMAB] Time, min.


0.818 0.425 0.284 0.0

0.764 0.371 0.314 15.5

0.705 0.312 0.354 31.0

0.656 0.263 0.397 49.0

0.599 0.206 0.463 70.0

0.571 0.178 0.506 85.0

k = 1.55 x 10-2 liter/mole/min.











KINETIC DATA FOR


TABLE 20

DMAB IN 50 PER CEIIT 1-PROPANOL-WATER
AT 14.6C.


[I] [DMAB] log [rft / [DMAB Time, min.


0.818 0.454 0.256 0.0

0.744 0.380 0.292 11.0

0.661 0.297 0.348 25.0

0.589 0.225 0.418 40.0

0.549 0.185 0.472 50.0

0.518 0.154 0.527 60.0

k = 2.87 x 10-2 liter/mole/min.






0.51
O
[DABQ, = 0.42 M
CH o= 0.82 M
T = 10.50C.
0.48


O

0.45





0.42





S 0.39
0
00


0.36





0.33


O


0.30





0.27
20 40 60 80 100

Time in Minutes

Fig. 13. Second Order Rate Plot for Dimethylamineborane (DMAB)
in 50 Per Cent 1-Propanol-Water









TABLE 21

SUMMARY OP THE KINETIC DATA FOR DEAB IN 50 PER CENT 1-PROPANOL-
WATER


k, l.mole-'min.-l log k 1/T, 103 x OK.-1


8.81 x 10-3 -2.055 3.575

1.55 x 10-2 -1.810 3.525

2.87 x 10-2 -1.542 3.475

AE- = 22.0 kcal./mole


Arrheniuc parameters from log k = A B/T: k in l.mole'1sec.-1

A = 13.64

B = 4,890







-2.2






-2.1


-2.0 --


-1.9 --


-1.8 --


-1.7






-1.6






-1.5






-1.4


3.49


3.51


3.53


3.55


3.57


1 x 103 oK.-1


Fig. 14. Arrhenius Plot for the Acid Hydrolysis of Dimethylamineborane
in 50 Per Cent 1-Propanol-Water







0.52


0.49





0.46


0.43 --


0.40 --


0.37 --


0.34 --


0.31 --


0.28


d I I I I I


3.63


3.59


3.55


3.51


1 x 103 oK.-1
T


Fig. 15. Solubility of Methylamineborane (MMAB) in Water


3.47








0.32





0.28


0.24 --


0.20





0.16
0
0



0.12





0.08


0.04 --


0.00 1
3.61 3.56 3.51 3.46 3.41


1 x 103 K.-l
T


Fig. 16. Solubility of Dimethylamineborane (DMAB) in Water








0.60


0.63


0.66





0.69

O


0.72
o
1-(
00



0.75

O



0.78





0.81





0.84
3.42 3.38 3.34 3.30

1 x 103 K.-1
T


Fig. 17. Solubility of Trimethylamineborane (TMAB) in Water


3.26






63
-1.14-





-1.10





-1.06






-1.02





-0.98
0
to
0



-0.94

0


-0.90





-0.86





-0.82
3.42 3.46 3.50 3.54 3.58

1 x 103 K.-1
T

Fig. 18. Solubility of Methylamineborane (MMAB) in Benzene













CHAPTER IV

DISCUSSION OF RESULTS


A. Aqueous kinetics

The reaction of the methylamineboranes in acidic solution proceeds

to completion according to the overall equation


R2CH3NBH3 + II + 31120 > R2CH3NIt + B(OH)3 + 3H2,


where R is CH3 or H. The reaction can be studied by several modes of

analysis. The IH concentration can be followed by measuring the in-

stantaneous pH of the solution, or by quenching with base and back titra-

ting. The amount of hydrogen gas evolved can be followed by manometric

techniques, taking into account the solubility of the gas in solution.

The reducing power of the solution can be followed by quenching the

reaction with excess iodate, converting the iodate to iodine with excess

acid and potassium iodide, and then back titrating the iodine with

arsenite. The following set of reactions describes this method:


R2CH3~BH3 + I03- >- B(O)03 + R2CH3N + I-


6HW + 10 3 + 51" > 312 + 3H20


12 + As033 + 120 > As043" + 21 + 21".


This method of analysis of boranes was first reported by Lyttle, Jensen,

and Struck42 for the borohydride group.










The stoichiometry of the reaction was established by a combination

of two methods. The hydrogen ion, boric acid, and methylanmonium ion

concentrations were determined by a potentiometric titration with OH",

as aliquots were removed during different stages of the reaction. The

instantaneous reducing power of the solution ras simultaneously deter-

mined by the iodate method. The experimentally determined stoichiometry

agreed with the overall equation for the acid hydrolysis indicated above.

These results also agree with RyschIkewitsch's57 study of the acid

hydrolysis of trimethylamineborane, in which the author concludes

that there is no build-up of B-H intermediates.

In this work the aqueous hydrolyses were followed by measuring the

instantaneous pH of the solutions. The excellent fit of the data to

second order kinetics, the agreement with the proposed stoichiometry,

and the conclusions of other investigators34,57 appear to preclude the

formation of any oxidizable B-H intermediates that could be detected by

our methods.

The general rate equation is given by


-d[R2CH3NBH3]/dt = k [R2CH3IBH3] [H], (1)


where k is the second order rate constant, t is the time, and the

brackets indicate molar concentrations. Expressing (1) in terms of the

initial concentrations and the amount reacted after some time, t; a and

b, and x, respectively, we have


dx/dt = k(a-x)(b-x). (2)


Separating the variables, integrating, and imposing the boundary condi-

tion that x is zero when t is zero, the resulting equation is









1 b(a-x)
kt (a-b) In a(b-x;) (3)


A plot of log (a-x)/(b-x) versus time should thus be linear with a slope

equal to k(a-b)/2.303. The linear plots obtained from our data are con-

firmation that the reaction is first order in amineborane concentration

and first order in hydrogen ion concentration. This is the same con-

clusion reached by Kelly, Marchelli, and Giusto39 in their investigation

of the hydrolysis of several amineboranes.

In a somewhat analogous manner to the treatment of the hydrolysis

of trimethylamineborane by Ryschkewitsch,57 there appear to be several

possible mechanisms consistent with the data. Besides the solvated pro-

ton, the acids water, methylaimonium ion, and boric acid could conceivably

react with the amineboranes according to the following equations:


R2CH3NBII3 + It H2 + products (4a)


R2CH3NBHI3 + H20 > H2 + products (4b)


R2CH3NBH3 + R2CH3N1 > H2 + products (4c)


R2CH3NBH3 + B(OH)3 > H2 + products (4d)

The reactions of the anineboranes with water, methylammonium ion, or

boric acid can be eliminated from consideration because of their com-

paritively slow reaction rates with respect to the rate of the reaction

with hydrogen ion in the time and concentration ranges studied.

There are two possible equilibria that could produce a BH3 species;

a displacement,


R2CH3NBH3 +- 1" R C R2CH3NH + BH3,










or a dissociation,


R2CH3NBH3 B R2CI3N + BH3. (6)


The BH3 fragment could then be attacked by the acids listed in

equations (4d)-(4d):


B113 + I -I > H2 + products (7a)


BH3 + H20 --- H2 + products (7b)


BH3 + R2CH3NH+ > % + products (7c)

BH3 + B(OH)3 > H2 + products (7d)


Another alternative is the direct attack of the solvated proton

on the B-H bond of the amineborane,


R2CH3NBH3 + H+ -- R2CH3NBHI2+ + H2. (8)


Hauthorne and Lewis34 have reported this type of mechanism in their study

of the hydrolysis of pyridine diphenylborane in acetonitrile. The most

direct approach would be a study of the isotope effect in the hydrolysis

of the anineboranes and their deuterated analogs, i.e., R2CH3NBH3 and

R2CH3NBD3. However, Davis et al,17 found that trimethylamineborane under-

goes a rapid and quantitative exchange of the boron hydrides with heavy

water, D20. Therefore, nothing would be gained by running the deuterated

compound in aqueous solutions if the rate of deuteration greatly exceeded

the rate of hydrolysis. Ryschkewitsch and Birnbaum58 studied the kinetics

of pyridineborane in 1-propanol and in 1-propanol-water solutions, and

the authors conclude that the reaction probably proceeds by loss of the








amine rather than H in the slow step. Kelly, Marchelli, and Giusto39

studied the solvolysis of several anineboranes and proposed that the

mechanism involved a rate determining attack of the solvated proton on

the amineborane, in which a proton is being transferred to the nitrogen

atom concurrent with a loosening of the B-N bond, followed by a rapid

solvolysis of the borane fragment. A borane cation of the type

(R3N)2BH2+ has been reported by. Miller and Muertterties45 and found to be

very stable toward acid attack. We would suspect that the borane cation

in equation (8) would be coordinated to the solvent water to form the

analogous species (R2CH3N)(H20)BH2+. On the basic of the evidence in the

literature cited above, we would thus eliminate equation (8) from con-

sideration as a possible mechanism.

First order kinetics in the acid concentration could result from

a rapid displacement equilibrium, followed by the slow reaction of the

borane fragment with a protonic species, equations (7b)-(7d). IWe would

like to show how this process can be eliminated from consideration. If

we assume the reaction proceeds by


R2CH3NBH3 + H+- k R2CI13NH + BH3
k-1


BH3 + acid k2 IH2 + products,


where k1 and k-, are the forward and reverse rate constants for the rapid

displacement equilibria, and k2 is the rate constant of the slow step,

then the rate expression would be given by rate = k2 IBH3] [acid]. The

BH13 concentration can be related to the equilibrium constant, K, of the

displacement equilibrium by










S [BH33 [R2CH3N +]
K [R2CH3NBH3] [11]

Solving for the BH3 concentration in terms of K and substituting this in

the rate expression, we arrive at the new rate expression


rate =k2KR2CH3NBH3] [H+] [acid]
rate = (9)
[R2CH3mt+]


It can be seen immediately xhy (7a) is eliminated. If the acid is I1, and

this is substituted in (9) for the acid, then this will give a rate

expression that is second order in hydrogen ion, which is contrary to

our observed data. However, substituting water, methylameonium ion, or

boric acid for the acid in (9) would present a rate expression that is

first order in hydrogen ion. The rate would also be inversely propor-

tional to the methylanmncnium concentration, as readily seen from equation

(9). Therefore the addition of the methylanmonium ion to the solution

should retard the rate, if this is the correct mechanism. This would be

true of course only for (7b) or (7d), which contain water and boric acid

as the attacking acids. Actually, the rate of the trimethylamineborane

hydrolysis is increased by the addition of trimethylannonium ion accord-

ing to Ryschkewitsch's57 study, and we would assume an analogous effect

for the other two adducts. Thus, we have so far eliminated a rapid dis-

placement equilibrium followed by a slow solvolysis of the borane frag-

ment by hydrogen ion, water, or boric acid. We would also like to

eliminate from consideration the equilibrium followed by the slow reaction

of the borane fragment with methylammonium ion, for the following reasons.

If step (7c) is analogous to reaction with hydrogen ion, and if the








reactivity increases with the acid strength, then both steps would be

competing and (7a) should be faster. Therefore, as long as the equi-

librium is maintained, step (7a) would produce second order kinetics in a

hydrogen ion. This has already been eliminated as being contrary to our

results. Therefore, a rapid displacement equilibrium followed by the

slow reaction of a borane fragment in steps (7a) through (7d) has been

eliminated from consideration as a possible mechanism. The dissociation

equilibrium (6), can also be eliminated, because the displacement

equilibrium (5), is a combination of (6) and the rapid proton transfer

equilibrium to the methylamionium ion,


R2CH3NBH3 z R2CH3N + BI5

+ R2CH3N + H P RP2CI3NH

= ~CH3NBH3 + H R2CH3NH + BH3.


This, of course, has been eliminated as indicated above.

However, if step (7a) were so fast that equilibrium is not main-

tained, then step (4a) would represent the reaction, i.e.,


R2CH3NBII3 + WH k > H2 + products.


Thus, first order kinetics in amineborane concentration and first order

kinetics in hydrogen ion concentration would be observed from the pro-

tonation of the nitrogen atom of the amineborane, concurrent with B-N

bond breakage, followed by a rapid solvolysis of the borane intermediate.

The reaction, including the transition state, can be conceived as








+
R2CH3NBH3 + 11 slow > R2C H3 R2CH3NIH' + B13



H


BH3 + H20 rapid > H2 + products.


This type of transition state is also in agreement with the ionic

strength dependence of the trimethylamineborane hydrolysis studied by

Ryschkewitsch,57 and that proposed by Kelly, Marchelli, and Giusto39 in

their study. The latter authors base their hypothesis on the influence

of substituents on the aniline ring in C6HBNH2BH3 and on dielectric effects

from using mixed solvents.

The activation energies and the Arrhenius parameters calculated for

the methylamineboranes are listed in Table 22 on the next pcae. On the

basis of the proposed transition state and the rate determining pro-

tonation of the amineborane, there are several factors which may be help-

ful in explaining the trend in activation energies. These factors include

any influences which weaken or strengthen the B-N bond, steric and in-

ductive effects, and the stabilization of the transition state and re-

actants by the solvent.

A study of the heat of the reaction


R2CH3N(g) + 1/2-B2116(g) >- R2CH3NBH3 g)


could be an indication of the relative B-N bond strengths of the adducts.

McCoy and Bauer4 have calorimetrically determined the heat of reaction

for the gaseous reactants to give the solid adducts. At the time of their

study, the heats of vaporization of the methylamineboranes had not been

determined. Values for the gaseous heats of reaction were obtained by








TABLE 22

THE ACID HYDROLYSIS OF THE METHYLAMINEBORANES


Methylamineborane A B AE0, kcal./mole Temp. range, C.


Water


lMethyl

Dimethyl

Trimethyl62


11.14

13.38

14.74


3,457

4,637

5,623


15.81

21.21

25.73


26.5-55.0

36.0-50.0

35.0-44.0


Dimethyl

Trimethyl


Arrhenius

energies:

log k =

A = log

B c AE


50 Per

13.6

12.3


Cent 1-Propanol-Water

4,890 22.0

5,110 23.4


6.5-14.6

31.8-44.4


parameters for the second order rate constants and activation


A B/T; k in 1.mole- sec.-1

A'

(1000)/2.303 R









adding the heats of vaporization found by Parry et al.50 to the values for

the solid adducts obtained by McCoy and Bauer. These values are listed

in Table 24 on page 82. The heat of dissociation of diborane enters into

the heat of reaction, but this will be true for each methylamineborane

and will not affect the trend in these heaOL. It is seen that the tri-

methylamineborane has the highest heat of reaction and would, on this

basis, be expected to have the highest activation energy. Analogous

arguments apply to the two other methylamineboranes, and the trend in

activation energies should parallel the trends in the absolute values of

the heats of formation of the amineboranes.

This same trend would also bo predicted on the basis cf inductive

effects. The trimethylamineborane would be expected to have the nitrogen

atom with the highest electron density compared to the other two methyl-

amincboranes, and would be expected to have the strongest B-N bond. The

stronger the B-N bond in the transition state, the more difficult it would

be to break this bond, and hence the activation energy should increase

with bond strength. Since the heats of reaction of the methylamines with

diborane is also a measure of the relative bond strengths, McCoy and

Bauer's work also indicates an increase in bond strength from methyl- to

trimethylamineborane, when their data has been adjusted to the gas phase.

However, some data on the interaction of methylamines and Lewis acids do

not exhibit this trend. Instead the trend appears to be dimethylamine >

monomethylamine. trimethylamine. This is true of Broxm, Bartholomay,

and Taylor's8 study of the gaseous dissociation of the methylaminetrimethyl-

boranes; Bauer's27 work on the methylamines and borontrifluoride; and the

data quoted by Braude and Nachod6 on the heats of reaction of the mcthyl-

amines and borontrifluoride in nitrobenzene. However, these reactions do








not involve the same substituents on the boron atom, are subject to sol-

vent effects in the case of the nitrobenzene measurements, and con-

sequently make strict comparisons difficult. Also, when the solid

adduct is formed from gaseous reactants, different Lewis acids will be

subject to different lattice effects which influence the heats of vapor-

ization.

In all of these cases, and in our case as well, there is only a

difference of a few kilocalories per mole for the heats of reaction of

the methylamines with the boranes. One can conclude that an explanation

of the trends in activation energies based only on the B-N bond strengths

or on inductive effects is inconclusive, but that these factors neces-

sarily influence the activation energies. Bauer27 also agrees that heats

of association and the activation energies of the methylamineborenes

should be only roughly related.

Assuming the first step in the mechanism, the slow or rate deter-

mining step, is the protonation of the nitrogen atom of the amineborane,

then there appear to be two important factors to consider. Sterically,

it would be expected that trimethylamineborane would be least susceptible

to an attack by a solvated proton. Conversely, the monomethylamineborane

would be the most readily attacked of the three amineboranes. Therefore

the monomethylamineborane might be expected to react faster than the

dimethylamineborane, etc. This does agree with the observed rates of

hydrolysis for the amineboranes.

One would also expect the solvent to influence this first step, de-

pending on the effective acidity of the solvated proton in different

media.









On the basis of electronic effects, the nitrogen atom in the mono-

methylamineborane would be expected to have the lowest electron density

of the three amineboranes. Therefore on the basis of attraction for the

solvated proton, the monomethylamineborane would react the slowest. How-

ever, the steric, electronic, and inductive effects may be minor compared

to other factors such as hydrogen bonding and dielectric effects, which

are discussed below.

In considering solvent effects, a hypothetical energy diagram can

be constructed showing the energies involved in going from the solvent

phase to the gas phase. This diagram is illustrated on the following

page. Attention will be focused on the reactants and the transition

states; the transition state is assumed to be a loosely bound methyl-

amnonium ion and a borane fragment. Once the B-N bond is loosened, the

borane fragment undergoes rapid solvolysis to the final products.

In the solvation process, the energy terms involved are the heats

of hydrrticn, AHh, of each species of reactants and transition states

and the activation energies of the gaseous and solvated transition states.

In the series of the methylamineboranes, the heat of hydration of the pro-

ton will be the same in all cases and should not affect the overall

energy differences. The borane fragment is present in each transition

state and one would expect that the borane portion of the transition

state would have an interaction energy with the colvent environment which

is independent of the aaine portion, even though the energy of formation

of each transition state from the gaseous reactants would be related to

the B-N bond strength.

Briegleb7 has found the heats of hydration of methyl-, dimethyl-,

and trimethylamine to be -10.4,-11.0, and -11.4 kcal./mole, respectively.















41


react.


sol'n. Hh prod.













Reaction Coordinate

Fig. 19. Reaction Path in the Aqueous (Lower Curve) and in the
Gas Phase (Upper Curve)










Therefore, as a first approximation, one would not expect the heats of

hydration of the methylamineboranes to differ by much. This is fairly

ucll substantiated by the experimental data obtained from the heats of

solution of the amineboranes listed in Table 23 on page 80, and from the

heats of vaporization found by Parry Ct al.50 The heats of solution

were determined from the least squares slope of a plot of log [methyl-

amincborane] versus 1/temperature for the saturated solutions. The

heats of hydration were then calculated by subtracting the heats of

vaporization from the heats of solution. The temperature ranges of the

solubility measurements do not coincide with the temperature ranges of

the kinetic data, but only a rough estimate was desired. The curvatures

in the solubility curves of the di- and monomethylamineborane (Figures

15 and 16 on pages 60 and 61, respectively) may be due to changes in the

degree of association of the molecules. Such behavior has been reported

by Noth and Beyer46 for the methylamineboranes. The values calculated

for the heats of hydration of tri-, di-, and monomethylamineborane are

-9.1, -12.9, and -12.5 kcal./mole, respectively. Since the heats of

solution were calculated from the slope of a log (concentration) versus

temperature plot, and since the curvature observed in the di- and

monomethylamineborane curves is toward a more rapid increase in the

saturation concentration than expected, then one might predict even higher

values for the heats of solution of these two adducts in the temperature

range of the kinetics, i.e., 25-50C. This in turn would lead to lower

negative heats of hydration of these two amineboranes, since the heats

of hydration were calculated from the differences between the heats of

solution and the heats of vaporization. The-efore the differences in the








heats of hydration of the methylamineboranes may even be closer together

than our data indicate.

It seems that the factor affecting the differences in activation

energies to the greatest extent is the heat of hydration of the tran-

sition state. If protonation occurs with the concurrent loosening of

the B-N bond to form loosely bound methylamnonium ion and BH3, then it

seems logical to discuss the differences in the heats of hydration of the

transition state on the basis of the differences ir. the heats of hydra-

tion of the methylamnmouium ions, assuming a fairly constant contribution

from the BH3 part of the molecule. Trotman-Dickenson65 has postulated

that in the equilibrium


R3NHI + 120 k H30+ + R311,


the degree of stabilization depends on the number of hydrogen bonds

capable of being formed, and on the aqueous ionization constants of the

amines. The author states that the effect of methyl substitution, although

increasing the base strength inductively, is more than compensated for by

the loss of hydrogen bonding; and that the amines will be stabilized to

approximately the same extent owing to solvation, but that the amine

ions will be stabilized much more, and to different degrees owing to

hydrogen bonding. Hall's33 study of the correlation of the base strengths

of the amines confirms Trotman-Dickenson's theory. Pearson and

Vogelsong51 have estimated the heats of hydration of the methylanmonium

ions to be -71, -63, and -55 kcal/mole for the methyl-, dimethyl-, and

trimethylammonium ions, respectively. These values were calculated from

a modified Born equation and from the lattice energies of the hydro-

chlorides and tetramethylanmmonium chloride. It is seen that there is a










decrease of approximately 8 kilocalories per mole in the heat of hydration

when methyl is substituted for hydrogen on the nitrogen atom. Therefore,

one might expect that a transition state involving a loosely bound methyl-

ammonium ion to a borane fragment is subject to somewhat analogous

effects. The monomethylamincborane transition state, capable of forming

the greatest number of hydrogen bonds cf the three methylamineboranes,

would be the most stabilized. Thus, the expected trend in activation

energies would be the same as that experimentally observed, for the

differences in the stabilization of the transition states are much greater

than the differences in the stabilization of the amineboranes.

It is proposed then, that a major factor in the observed trend in

activation energies of the methylamineboranes is the stabilization of

the transition state by the solvent water. It would be interesting to

determine the activation energies of several series of amineboranes, such

as the ethyl- or n-propylamineboranes, to test the hypothesis, and to

evaluate the consequences of possible steric effect which have been pre-

viously discussed. A discussion of the activation energies in a mixed

solvent is undertaken in Part B of this section.

Sunnarizing the factors contributing to the observed trend in

activation energies, it is proposed that steric effects play a relatively

minor role, that the activation energies follow the trend expected from

the B-N bond strengths, but that in addition a most important effect is

the stabilization of the methylaimonium ion part of the transition state

by hydrogen bonding in the aqueous solvent.

B. Kinetics in the 50 per cent 1-propanol-water mixed solvent

The kinetics of trimethylamineborane in l-propanol-water were con-

siderably different at large trimethylanineborane concentrations than












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CO h









I...
*4 *
I (







vHI










those at low concentrations, and the kinetics for these two cases will

be discussed separately.

Trimethyl- and dimethylamineborane concentrations greater than

0.300 M. The kinetic data for the acid hydrolysis of tri- and dimethyl-

anineborane show a linear fit to the second order rate expression pre-

viously discussed. Dimethylamineborane kinetics were not investigated

at low concentrations, but the trimethylamineborane kinetics do not obey

the second order rate expression at concentrations around 0.02 M.

The 1-propanol-water azeotrope used in the experiments contains

71.7 per cent 1-propanol and 28.3 per cent water, by weight,43 as pre-

pared in Part A of the experimental section. The volume ratio of the

azeotrope contains 75.8 per cent 1-propanol by volume, as calculated

from the room temperature densities of water and 1-propanol. In each

experiment two volumes of the azeotrope were mixed with one volume of

the aqueous acid (giving a mixture of 50.6 per cent 1-propanol by volume)

and throughout this paper this mixture will be referred to as 50 per

cent 1-propanol-water.

A comparison of the rate constants found in the mixed solvent with

the rate constants found in the pure water is listed in Table 24 on the

following page, and shows the rates to be slower in the mixed solvent.

This is similar to the behavior found by Ryschkewitsch57 in the acid

hydrolysis of trimethylamineborane in 20 per cent dioxane-water solutions,

and that reported by Kelly, Marchelli, and Giusto39 in their investigation

of the hydrolysis rates of several amineboranes in 50 per cent dioxane-

water solutions. Ryschkewitsch and Birnbaum,58 however, found an increase

in the rate of solvolysis of pyridineborane in 1-propanol-water soltuions

compared to the rate in pure water. These authors used the solvents









TABLE 24


COMPARISON OF THE SECOND ORDER RATE CONSTANTS
TRIMETHYLAMINEBOANE IN WATER AND IN 50 PER CENT


OF DI- AND
1-PROPANOL-WATER


NIthylamine- k in 50 per cent
borane k in water* 1-propanol-water Temperature, OC.


Dimethyl 6.50 x 10-4 1.47 x 10-4 6.5

1.10 x 10-3 2.58 x 10-4 10.5

1.91 x 10-3 4.78 x 10-4 14.6



Trimethyl 1.87 x 104 3.87 x 10-5 31.83

3.24 : 10-4 6.88 x 10-5 35.85

5.83 x 10-4 1.13 x 10-4 40.23

1.00 x 10-3 1.84 x 10~4 44.38

k in liter/mole/sec.


*These values were calculated at the temperatures indicated from the
relation log k = A B/T, where the parameters A and B are given in
Table 22 on page 72.










themselves as the hydrolyzing acids, and this is probably why the behavior

is different from the other cases listed above. The pyridineborane

solvolysis involves the dissociation of the adduct into pyridine and

a borane fragment as the rate determining step, whereas our data indi-

cate the rate determining step in the acid hydrolysis is the protonation

of the nitrogen atom of the amineborane.

The decrease in rates in the mixed solvents may be caused by a

decrease in the pre-exponential factor, A. This factor can be related to

the entropy of activation, but definite conclusions based on entropy con-

siderations would be difficult from the present state or our knowledge of

the liquid state. The differences in the activation energies in 1-

propanol-water and in water only will be discussed on the basis of di-

electric and hydrogen bonding effects.

The effect of stabilizing the transition state more than the ground

state by hydrogen bonding has been discussed earlier in this section. It

was postulated that the trend in activation energies observed in the acid

hydrolysis of methylamineboranes can largely be ascribed to the differ-

ences in the heats of hydration of the methylannonium ion portion of the

transition state caused by hydrogen bonding, the heats of hydration of

the methylamine adducts contributing little. On this basis, the transi-

tion states of the amineboranes would be stabilized more than the amine-

bcranes themselves by hydrogen bonding with the solvent. Therefore as

the hydrogen bonding power of the solvent is reduced, one would predict

a higher activation energy for each methylamineborane due to the relative

destabilization of the transition states in the mixed solvent. Thus, in

the acid hydrolysis of di- and trimethylamineborane in 50 per cent 1-

propanol-water, we would expect a higher activation energy in the mixed









solvent for both compounds than in water only. Moreover, the dimethyl-

amineborane transition state should be affected to a relatively greater

degree than the trimethylamincborane transition state by changes in the

hydrogen bonding power of the solvent. The activation energies obtained

in this investigation are listed in Table 22 on page 72, and it is seen

that although the activation energy for dimethylamineborane in the mixed

solvent is higher than in pure water, the activation energy of the tri-

methylamineborane is lower in the mixed solvent than in water.

We interpret these data to be the result of two opposing trends.

Decreasing the hydrogen bonding power of the solvent tends to increase

the activation energy, with the sensitivity increasing from trimethyl-

to monomethylamineborane. However, decreasing the dielectric constant of

the solvent would tend to decrease the activation energy. This is due

to the relatively greater destabilization of the ground state compared

to the transition state when the dielectric constant of the solvent is

decreased. Now, the dilution of water with 1-propanol reduces both the

hydrogen bonding power of the medium and the dielectric constant. Re-

ducing the dielectric constant would then destabilize the ground state

relative to the transition state and lead to a lower activation energy.

The acid hydrolysis of the amineboranes can be imagined as the

reaction between a small, positively charged, solvated proton and the

solvated amineborane dipole to form the larger, positively charged, sol-

vated transition state,


R2CH3N. BH3


H










We can assume that the positive charge on the transition state is more

diffuse than that on the solvated proton, and consequently the solvated

proton will be destabilized more than the transition state by a reduction

of the dielectric constant of the solvent. Moreover, we might expect

this destabilization to be fairly constant throughout the methylamine-

borane series because of the relatively small differences in the sizes

of the methylamineboranes and their postulated transition states.

The observed trend in activation energies in the mixed solvent can

now be rationalized on the basis of whether the dielectric effect or the

hydrogen bonding effect is predominant. The trimethylamineborane tran-

sition state is affected less by hydrogen bonding changes in the solvent

than the dimethylamineborane transition state, and therefore the effect

of changing the dielectric constant of the solvent might be the control-

ling factor in determining the activation energy. Thus, in the 50 per

cent 1-propanol-water solution the hydrogen bonding power of the solvent

is reduced, but this is more than compensated for by the concurrent re-

duction in the dielectric constant of the solvent. The dimethylamine-

borane transition state, being affected to a relatively greater extent

by the decrease in hydrogen bonding power, could have a higher activation

energy in the mixed solvent than in pure water. In the case of the di-

methylamineborane, it is assumed that the dielectric effect is of the

same order as the trimethylamineborane, but that the larger influence of

the loss of hydrogen bonding power leads to an increase in the activation

energy in the mixed solvent over water. We would therefore predict an

even larger difference between the activation energies of monomethylamine-

borane in the mixed solvent and in water than the difference in activation

energies of dimethylamineborane in the same media.









Trimethylamineborane concentrations below 0.02 M. Originally, it

was decided to run the kinetics of trimethylamineborane in 50 per cent

1-propanol-water as a pseudo-first order reaction using a large excess

of hydrochloric acid and following the reaction by the loss of reducing

power of the solution. However, reproducible data could not be obtained.

The unexpected behavior of these solutions will be discussed along with

the implications of each experiment. A brief summary of these reactions

is listed in Table 25 on the following page.

The upper curve in Figure 20 on page 88 is the pseudo-first order

plot of reducing strength versus time for the hydrolysis of approximately

0.02 M trimethylamineborane and 1.0 M hydrochloric acid in 50 per cent

1-propanol-water at 31.40C. This type of deviation from linearity could

be caused by the build-up of intermediates of the type BH2OR and BH(OR)2,

or R2CH3NBH20R and R2CH31M'I(OR)2. If these intermediates were present,

then an analysis.of the reducing power of the solution would erroneously

attribute these reducing equivalents to unreactcd trimethylamineborane.

From the results of the experiments outlined below, we would like to

eliminate the possibility of intermediate build-up.

If we assume that the mechanism of the reaction involves the

loosening of a B-N bond in the transition state to form the ammonium ion

and a borane fragment, and that the borane fragment then forms inter-

mediates with the solvent, then the other methylamineboranes as well as

diborane would be expected to form the same intermediates. These BH2OR

and BH(OR)2 intermediates would also be expected to have the same reac-

tion rates, regardless of the Lewis base to which the borane fragment was

originally attached. Solutions of dimethylamineborane and diborane with

with approximately the same initial concentrations as those used in the










TABLE 25

SUMMARY OF TIE EFFECTS PRODUCED BY VARYING CONDITICUS IN THE
ACID HYDROLYSIS OF TMAB AT LOW CONCENTRATIONS IN
50 PER CENT 1-PROPANOL-WATER


Effect on rate


Added solid TMAB to spent solution

Added allyl alcohol tc fresh solution

Added allyl alcohol to spent solution

Added trimethylaimonium ion and boric
acid to fresh solution

Used alcohol distilled froi a spent
solution

Used alcohol solution in which tri-
methylammonium ion, boric acid,
and hydrochloric acid had sat
for twelve hours


Marked decrease

No change

No change


No change


Slightly accelerated


No change


Condition varied







88
-2.3






-2.1




Original

-1.9






-1.7
c


-4





0)
0o








1 -1.3 Spent














-0.9
-1.1










-0.7 -




1 2 3 4 5


Time in Hours

Fig. 20. Pseudo-First Order Rate Plot for Trimethylamineborane in
50 Per Cent 1-Propanol-Water for 0.02 M TMAB










trimethylamineborane hydrolysis were prepared. The reactions with these

two compounds were very rapid, and over 90 per cent of the reaction ex-

hibited no tendency to slow to a rate comparable to that of the trimethyl-

amineborano. We therefore concluded that if our mechanism is operative

in the acid hydrolysis of trimethylamineborane in 1-propanol-water solu-

tion, there was no build-up of oxidizable intermediates of the type

BH20R and BH(OR)2.

If the intermediates were of the type R2CH3NBH20R and R2CH3NBH(OR)2,

then the stoichiometry would be different from that originally proposed;

i.e.,


R2CH3NBH3 + HI + 3H20 > B(OH)3 + 3112 + R2CH 3NI.


Aliquots were removed from an acid hydrolysis in 1-propanol-water of

similar concentrations to the original experiment with trimethylamine-

borane, and these aliquots were simultaneou!Ay analyzed for reducing

power by the iodate method and for boric acid, hydrogen ion, and tri-

methylammonium ion by a potentiometric titration. This analysis showed

that for each six equivalents of amineborane used, one mole of hydrogen

ion was used, one mole of boric acid was produced, and one mole of tri-

methylanmonium ion was produced. This appears to confirm the stoichio-

metry originally proposed for the acid hydrolysis. We therefore elimin-

ated a build-up of this kind of intermediate as an explanation of the

behavior exhibited by dilute trimethylamineborane hydrolysis in alcohol

solutions.

The lower curve in Figure 20 on page 88 was obtained from the data

when solid trimethylamineborane was again added to the solution originally

used to obtain the upper curve, and in which reaction had gone to









completion. The rate was now markedly slower and fitted pseudo-first

order kinetics. This could indicate a retardation of the rate by one of

the products of the hydrolysis. A solution approximately 0.02 molar in

trimethylaeaonium ion and boric acid was prepared in 50 per cent 1-

propanol-water. Solid trimethylamineborane was added to this solution;

no loss in reducing power was observed over a five hour period. This

indicates that the triiaethylamineborane itself does not react with the

products of the hydrolysis. A 50 per cent 1-propanol-water solution was

made 1.0 M_ in hydrochloric acid and 0.02 1 in boric acid and trimethyl-

ammonium ion. Solid trimethylamineborane was added to the solution; the

kinetic behavior of this reaction was the same as the behavior of the

reaction in a fresh solution without addition of reaction products. This

evidence appears to eliminate the possibility of rate retardation by a

product of the hydrolysis.

Another possibility of rate retardation was that the hydrogen ion

concentration was being depleted by an equilibrium which was slow to be

established relative to the half-life of the reaction. A 50 per cent

1-propanol-water solution was made 1.0 1M in hydrogen ion and allowed to

stand overnight. Solid trilcet.ylamineborane was again introduced, but

the behavior did not change. No change in the behavior was observed when

solid trimethylamineborane was added to a 50 per cent 1-propanol-water

solution containing 1.0 I1 hydrochloric acid, and 0.02 M in trimethyl-

anraonium ion and boric acid. It was concluded that the hydrogen ion con-

centration was not being depleted in the initial stages of the reaction

by an equilibrium.

If there were an equilibrium between the acid and the alcohol

and a competition of these two acids for the amineborane, a second order









rate expression would result anyway. Consider the equilibrium constant


Keq ROHI0 ] [] [H20]/ IROH] 1130],

where R is the n-propyl group. Assuming the [H20]/EROH] = 1/K ratio

remains constant, then [ROH2+] equals K'[H30+]. The two competing rates

are given by the equations

TMAB + 1130+ k I2 +I products and


TMAB + ROH2+ k2 H + products.


The rate expression is then given by


rate = k [TMAB] [H30 ] + k2 I[T1A] (ROH12+]


Substituting the [ROH2+] in the rate expression, we arrive at

rate = k [TlAB] [113O] + k2K' [ThAB] [H30+],


or factoring,

rate = [TMAB] (H130 ](k1 + K'k2).


Since tk1 + K'k2) is also a constant, this type of equilibrium should

also give second order kinetics. Therefore this line of reasoning cannot

explain the type of behavior observed.

Attempts were made to determine the initial rates of the reaction

while varying one concentration and keeping the other constant. How-

ever, reproducible results could not be obtained.

The next possibility investigated was that an impurity in the alco-

hol might be accelerating or retarding the rate of the reaction. It was




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