I THE EFFECT OF ATOMS UPON THE SN2 REACTIVITY OF OTHER HALOGEN ATOMS ATTACHED TO THE SANE CARBON ATOM

II A STUDY OF THE RATE OF PROTON TRANSFER REACTIONS

A THESIS Presented to the Faculty of the Graduate Division

by

Cyrus Henry Thomas

In Partial Fulfillment of the Requirements for the Degree Doctor of Philosophy in the School of Chemistry

Georgia Institute of Technology June 1953 I lilt hiikECT OF HALOGEN ATOMS UPON THE SN2 REACTIVITY OF OTHER HALOGEN ATOMS ATTACHED TO THE SAME CARBON ATOM

II A STUDY OF THE RATE OF PROTON TRANSFER REACTIONS

Approved:

r

• V • Date approved by Chairman: /14 4,ld / 95-3

ii Acknowledgement

I wish to express my sincere appreciation to Dr. Jack Hine, not only for the suggestion of the problems and his valuable aid and guidance, but also for the confidence which he displayed in my abilities by granting me a Fellowship and a Research Assistantship.

Secondly, I wish to thank the Research Corporation of New York and the Atomic Energy Commission for the funds which they supplied for these investigations. Thirdly, I wish to thank my wife for the constant encouragement she has given me and the never-ending patience she has displayed, without which I may never have completed this work.

iii TABLE OF CONTENTS

ACKNOWLEDGEMENTS LIST OF TABLES LIST OF ILLUSTRATIONS SUMMARY

PART I Chapter I INTROIUCTION 2 II PROCEDURE 5 Experimental Preparation and Purification of Reagents III DISCUSSION OF RESULTS 21

IV CONCLUSIONS 33 V RECOMMENDATIONS 34

APPENDIX A - SAMPLE CALCULATIONS 35 APPENDIX B - TABLES 44

APPENDIX C - GRAPHS 88

BIBLIOGRAPHY 97

iv PART II

Page Chapter I INTRODUCTION 100 II INSTRUMENTATION AND EQUIPMENT io6 III PROCEDURE 118 Experimental Preparation and Purification of Reagents IV DISCUSSION OF RESULTS 126 V CONCLUSIONS 132 APPENDIX A - SAMPLE CALCULATIONS 133 BIBLIOGRAPHY 140 VITA 142

V LIST OF TABTRS PART I

Table Page

1. Determination of Interference 7 2. Summary of Results for the Reaction RX + Y- RY + X- 23 3. Alkyl Bromide Plus 26 4. Alkyl Iodide Plus Methoxide Ion 28 5. Alkyl Bromide Plus Methoxide Ion 30 6. of Compounds Used 45 7. cH3I + NaOCH3 in Absolute at 20.3°C. 46 8. m131 + Na0CH3 in Absolute Methanol at 50 °C. 47 9. cH2cli + NaOCH3 in Absolute Methanol at 20.3°C. 48 10. CH2C1I + NaOCH3 in Absolute Methanol at 50 °C. 49

11. CH2BrI + NaOCH3 in Absolute Methanol at 20.3 °C. 50 12. CH2BrI + NaOCH3 in Absolute Methanol at 50 °C. 52

13. CH2I2 + NaOCH3 in Absolute Methanol at 20.3 °C. 54 14. CH2I2 + NaOCH3 in Absolute Methanol at 50°C. 56

15. CH3Br + NaOCH3 in Absolute Methanol at 20.3 °C. 57 16. CH3Br + NaOCH3 in Absolute Methanol at 50°C. 58 17. CH3CH2Br + NaOCH3 in Absolute Methanol at 20.3 ° C. 59 18. CH3CH2Br + NaOCH3 in Absolute Methanol at 50 °C. 60 19. CH2FBr + NaOCH3 in Absolute Methanol at 20.3 °C. 61

20. CH2FBr + NaOCH3 in Absolute Methanol at 50 °C. 62

21. CH2C1Br + NaOCH3 in Absolute Methanol at 20.3 °C. 63

vi 22. CH2C1Br + NaOCH3 in Absolute Methanol at 36°C. 64

23. CH2C1Br + NaOCH3 in Absolute Methanol at 50°C. 65

24. CH2Br2 + NaOCH3 in Absolute Methanol at 20.3°C. 66 25. CH2Br2 + NaOCH3 in Absolute Methanol at 36 °C. 67 26. CH2Br2 + NaOCH3 in Absolute Methanol at 50 °C. 68

27. CH2C12 + NaOCH3 in Absolute Methanol at 50 °C. 70 28. CH3Br + NaI in Acetone at 0 °C. 71

29. CH3CH2Br + NaI in Acetone at 20.3 °C. 72 30. CH3CH2Br + NaI in Acetone at 50 °C. 73 31. CH2BrF + Nal in Acetone at 20.3 °C. 74

32. CB2BrF + NaI in Acetone at 50 °C. 75 33. CH2C1Br + NaI in Acetone at 20.3 °C. 76 34. CH2C1Br + NaI in Acetone at 36°C. 77 35. CH2C1Br + Nal in Acetone at 50 °C. 78 36. CH2Br2 + NaI in Acetone at 20.3°C. 79 37. CH2Br2 + NaI in Acetone at 36°C. 80

38. CH2Br2 + NaI in Acetone at 50°C. 81

39. CH2BrI + Nei in Acetone at 20.3 °C. 82 40. CH2BrI + NaI in Acetone at 50 °C. 83 41. CH2C12 + Nei in Acetone at 50 °C. 84

42. CH2C12 + NaI in Acetone at 60 °C. 85 43. CH2C1I + NaI in Acetone at 50 °C. 86 44. CH2C12 + NaOCH3 in Absolute Methanol 87

vii PART II

Table Page

1. Calibration Data for Micropipette 114

2. Data Obtained by Orr 126

3. Calculated Rate Constants 127

4. Calculated Equilibrium Percentages 128

5. Results of Exchange Between Heavy and 130

6. Results of Exchange Reactions 131

viii LIST OF ILLUSTRATIONS PART I

Figure Page 1. Densities of Halomethanes 89 2. Log of Rate Constant versus the Reciprocal of the Absolute Temperature 90

3. Relationship Between Rate of Reaction of Alkyl Bromides with NaI and NaOCH3 91

4. CH2BrI + NaOCH3 in Methanol at 20.3 °C. 92

5. cH2I2 + NaOCH3 in Methanol at 20.3 °C. 93 6. CH3CH2Br + NaI in Acetone at 20.3 °C. 91 .

7. CH2BrI + NaI in Acetone at 20.3 °C. 95 8. CH2BrI + NaI in Acetone at 50 °C. 96

ix PART II

Figure Page

1. Mercury-Toluene Thermoregulator 108

2. Micropipette 110

3. Micropipette Mounting Assembly 112

4. Calibration Curve for Micropipette 113 5. Infrared Absorption Spectra for Partially Deuterated t-Butylalcohol and t-Butylalcohol 115

6. Infrared Absorption Spectra for Ethylamine and Partially Deuterated Ethylamine 116 7. Calibration Curve for Determination of Ethanol 117 Summary

Part I

Since one of the most important reactions in organic chemistry is the nucleophilic displacement reaction upon carbon, of the type:

RX + Y --is RY + X , it is of considerable interest to determine the effect of various groups in R upon the reaction. It is known for instance, that methyl groups on the alpha carbon atom of R have an inhibiting effect upon the SN2 reactivity and an enhancing effect upon SN1 reactivity. It has been shown semi-quantitatively that a halogen atom on the alpha carbon decreases SN2 reactivity but it is not known which have the greatest effect nor is the extent of this effect known. Hence, we were interested in determining the effect of alpha halogen atoms upon the SN2 reactivity of other halogen atoms attached to the same carbon atom under various conditions and with various reagents.

One method of attack was through the well known reaction of iodide in acetone upon the dihalo methanes, for example:

C1CH2Br + NaI --is C1CH2I + NaBr.

In a case such as this the bromine atom is the one assumed to react since in all known cases bromine in an organic reacts many times faster than chlorine. The reaction was followed through the rate of

xi disappearance of the sodium iodide by titration of the iodide ion with potassium iodate in the presence of cold concentrated hydrochloric acid, according to the equation: 21 - + 103 + 6HC1 31C1 + 3H20 + 3C1 -

The second method of attack was through the reaction of sodium methoxide in absolute methanol with the dihalo methanes, for example:

ow NaOCH3 C1CH2Br + NaOCH3 Sl C1CH2OCH3 CH2(0CH3)2 Fast

Here again it was assumed that the bromine atom reacts first and since it was shown in the course of this investigation that chloromethyl methyl ether reacts essentially instantaneously with sodium methoxide, the first step must be the slow, rate determining step. The reaction was followed through the rate of disappearance of sodium methoxide by quenching the reaction mixture in excess standard hydrochloric acid and titrating the excess with standard .

For both methods the experimental technique was the same. For runs at the higher temperatures sealed tubes were used. Those at lower temperatures were run in one hundred milliliter volumetric flasks and aliquots taken at appropriate time intervals. Those organic compounds containing and another halogen were found to be light sensitive and made it necessary to run the sealed tube experiments in the dark and the others in low-actinic glassware.

The results are tabulated in the following tables.

xii Table 1. Alkyl Bromide Plus. Iodide Ion

Alkyl k x 105 1 mole -1 sec -1 4Ha ASa Group 20.3 °C. 50°C. Kcal e.u.

CH3cH2 123.1 + 12.3** 1748 + 69 15.6 -18.4 cH2F 55.58 + 4.55 1350 + 18 19.4 -7.3 CH2C1 7.212 + .247 220.7 + 2.4 20.8 -6.4

CH2Br* 2.073 + .060 69.08 + .92 21.4 -6.9

C1121 5.66 + .57** u6 18.5 -14.9 CH Too fast to measure with any accuracy even at 0 3 °C. The specific rate constant is of the order of 0.03 1 mole -1 sec -1 .

* Rate constants contain a statistical factor of one-half.

** Deviation estimated for this value, all others are average deviations. Table 2. Alkyl Iodide Plus Methoxide Ion

Alkyl k x 105 1 mole -1 min-1 AHa ASa Group °C. 50°C. 20.3 Kcal e.u.

CH3 15.01 + .39 433.0 + 8.7 20.5 -6.2

CH2C1 .0863 + .0038 4.417 + .167 22.3 -9.7 CH2Br .031 + .003** 1.023 + .027 21.3 -15.7

CH2I* .01225 + .0012** .536 + .007 23.0 -11.5

* Rate constants contain a statistical factor of one-half.

** Deviation estimated for this value, all others are average deviations.

xiv Table 3. Alkyl Bromide Plus Methoxide Ion

Alkyl k x 105 1 mole -1 min-1 AHa 41Sa Group °C. 50°C. 20.3 Kcal e.u.

CH3 17.07 + .58 479 + 24 20.3 -6.6

CH2F 7.613 + .170 218.5 + 9.7 20.5 -7.4

CH3CH2 1.33 + .02 47.2 + 1.6 22.8 -3.2

CH2C1 .0413 + .0010 2.355 + .022 24.7 -3.4

CH2Br* .00613 + .0002 .3703 + .017 25.1 -5.9

* Rate constants contain a statistical factor of one-half. From these data it is concluded that the order of decreasing effect upon the SN2 reactivity is,

F ,> Cl > I Br.

Part II

A Study of the Rate of Proton Transfer Reactions

Since it is rather commonly accepted that "all proton transfers between and atoms in the liquid phase are too fast to measure", it was deemed desirable to make an experimental test of this statement.

An examination of the literature revealed a claim that the transfer of a deuteron between and ethanol occurred at a rate slow enough to measure. These results had been obtained by mixing ethanol and heavy water, removing the water as the hydrate of calcium sulfate and analysis of the resulting water. Using approximately the same procedure and analyzing for deuterium content by the falling drop method it was found that this reaction does not proceed at a rate slow enough to determine under these conditions. This was further borne out when recalculation of the original work revealed that a very small error in weighing (4;0.4 mg) was sufficient to account for the apparent slow reaction. It was felt that compounds which have low acidity constants and comparable basicity constants or those which are sterically hindered would have the best chance of undergoing a slow proton transfer. As a result the exchange between ethyl and heptyl amines, ethyl amine and tertiary-amyl alcohol, and tertiary-amyl alcohol and di-tertiary-butyl iso-propyl carbinol was studied. In these runs the reactants were mixed and in one case separated as quickly as possible after mixing and in a second case they were allowed to stand together for periods over eight hours. The analyses were made using infra-red absorption techniques. It was found that the deuterium content was the same in the first case as in the second. From these data it was concluded that the transfer studied was too fast to measure by this technique.

xvii 1

PART I

'Tit; EFFECT OF HALOGEN ATOMS UPON THE SN2 REACTIVITY OF OTHER HALOGEN ATOMS ATTACHED TO THE SAME CARBON ATOM 2

CHAPTER I

INTRODUCTION

One of the most important types of reaction in all organic

chemistry is the nucleophilic displacement reaction on carbon.

RY + X ---31.• RX + Y

Because of the importance of this reaction a great deal of interest

has been aroused concerning the effect of various substituents upon

the rate of displacement of various groups. It has been found, for

example, in the case of the methyl halides that replacing one of the

atoms with an alkyl group causes the displacement of the

halide to be slowed considerably in reactions which are considered to

be of SN2 type. Also, if the hydrogen is replaced with a phenyl

group, the reactivity of the halogen is noticeably increased. Of

all the groups studied, the least is known about the effect of halogens.

It thus seems very desirable to study this effect with various reagents

and under various conditions.

This type of reaction is believed to proceed by way of two basic

reaction mechanisms which have been designated by Hughes and Ingold (1) as SN1 and SN2; i.e., substitution nucleophilic first order and substi- tution nucleophilic second order. Since recent work of Hine and Lee (2,3)

1. E. D. Hughes, Faraday Soc., 37, 603 (1941). 2. J. Hine and D. E. Lee, J. Am. Chem. Soc., 73, 22 (1951). 3. J. Hine and D. E. Lee, ibid, 74, 3128 (1952). 3

has been concerned with the effect of alpha halogens upon the SN1

reactivity of other halogens attached to the same carbon atom, the

present work is concerned with the effect of alpha halogens upon the SN2 reactivity of other halogens on this same carbon. The only work of this nature which has been reported in the literature is a semi-quantitative treatment carried out by Petrenko- Kritschenko and his co-workers (4-11) in which they found that the reaction of alkyl halides, which normally undergo SN2 type reactions,

was slowed down by the addition of a second halogen to the same carbon atom. However, this was not done with any attempt to correlate the

variations of SN2 reactivity with the variations of halogen substituents.

The proposed mechanism for an SN2 type reaction may be depicted as:

O O

4. P. Petrenko-Kritschenko et al, Z. Physik. Chem., 116, 313 (1925).

5. P. Petrenko-Kritschenko, J. Prakt. Chem., 111, 23 (1925). 6. P. Petrenko-Kritschenko and V. Opotsky, Ber., 59B, 2131 (1926).

7. P. Petrenko-Kritschenko et al., J. Russ. Phys. Chem. Soc., 60, 149 (1928).

8. P. Petrenko-Kritschenko et al., Ber., 612, 845 (1928). 9. P. Petrenko-Kritschenko et al., ibid., 62B, 582 (1929).

10. P. Petrenko-Kritschenko et al., J. Russ. Phys. Chem. Soc., 61, 1778 (1929).

11. P. Petrenko-Kritschenko et al., Ukrainskii Khemichinii Zhurnal, 1, 304 (1924). 14.

The attacking group Y approaches the molecule from the rear on a line

of attack approximately along the axis of the C-X bond. Since this is a nucleophilic attack by an electron rich species seeking an electron- poor center, it seems that the approach of the attacking group would be facilitated by any group which would tend to lower the electron at the reaction center and that those which tend to increase

the electron density at this center should do the opposite. At the same time, those groups which have a tendency to lower the electron density are effectively increasing the kernel charge on the center and

should in this way increase the strength of the C-X bond. Thus we have two effects upon the rate of reaction working in opposition to one another.

There are, of course, other effects which enter into the rate at which an SN2 reaction proceeds. Among these we must include steric factors, ion-dipole repulsions and polarizability of both the attacking species and the molecule being attacked. 5

CHAPTER II

PROCEDURE

Experimental

If data are to be obtained which can be used to evaluate the

effect of an alpha halogen atom on the SN2 reactivity of other halogens,

it is necessary that the effects of other groups in the molecule be

minimized. For this reason it was decided to study the reactivity of

the various halides towards various nucleophilic reagents

under conditions which promote SN2 reactions. Since solvents of low

solvating power tend to inhibit SN1 reactions and promote SN2 reactions

with anions, it was decided to use iodide ion in acetone (12-14) and

methoxide ion in absolute methanol as the nucleophilic reagents for

these studies.

I - + XCH2Y XCH2I + Y -

9C11 OCH3- + XCH2Y XCH2OCH3 + Y- Fast. CH2(OCH3)2

X and Y are both halogens and in some cases X and Y are the same.

12. J. B. Conant and W. R. Kirner, J. Am. Chem. Soc., 46, 232 (1924).

13. J. B. Conant and R. E. Hussey, ibid., 47, 476 (1925).

14. J. B. Conant, W. R. Kirner and R. E. Hussey, ibid., 47, 488 (1925). 6

The kinetics of the iodide ion in acetone reaction is followed

by titration of excess iodide ion with standard potassium iodate in

very concentrated hydrochloric acid, keeping the solution very cold,

according to the following equation:

I(1) + 21 - + 6HC1 31C1 + 3H20 + 3C1 -

The titration is that of Andrews (15) as modified by Conant and co-

workers (12-14). Twenty-five milliliters of cooled concentrated

hydrochloric acid, five milliliters of carbon tetrachloride and some

crushed ice are placed in a large-mouthed, glass-stoppered bottle.

The tube of reactants is broken into the bottle. Iodate solution is

then added in a fast stream until the original yellow color of the

solution just disappears. At this point, more crushed ice is added and the mixture is shaken thoroughly. The titration is continued with frequent shaking and additions of ice until the iodine color disappears from the carbon tetrachloride layer.

Since it had been found by Senior, Hetrick and Miller (16) that acetone interferes in the determination of iodide ion by the method of

Andrews when chloroform is used as indicator, a control experiment was run to determine if acetone gave the same interference when carbon tetrachloride was used as indicator. Five milliliters of a solution of

0.4122 grams of sodium iodide per 100 ml. of solution was titrated by

15. L. W. Andrews, J. Am. Chem. Soc., 25, 736 (1903).

16. K. I. Senior, R. R. Hetrick and J. G. Miller, J. Am. Chem. Soc., 66, 1987 (1944). 7

this method both in the presence and absence of 5 ml. acetone. The following data were obtained:

Table 1. Determination of Acetone Interference

Solvent Ml. of 0.0100 N Molarity KI03 NaI

No acetone 27.60 0.0276 No acetone 27.57 0.0276 No acetone 27.52 0.0275 Acetone added 27.51 0.0275 Acetone added 27.55 0.0275 Acetone added 27.50 0.0275 Calculated molarity of NaI = 0.0275

It can be seen that the added acetone has caused no error in this determination and thus carbon tetrachloride may be used as indicator without fear of inaccuracy from this source.

The reaction itself was run in a series of sealed tubes , when the temperature was near the boiling point of one of the liquids used, in a constant temperature bath which is regulated to hold within + 0.05 °C. Small, thin-walled ampoules of approximately 0.1 ml. capacity, which are very similar to miniature Dumas bulbs, were blown from soft glass tubing and weighed. The ampoules were then filled with 0.1 ml. of methylene halide by surrounding the bulb with crushed 8

dry ice and inserting the needle of a hypodermic syringe into the neck

of the ampoule and discharging the required volume of the halide into

the bulb. The neck was then sealed, being careful not to loose any

glass from the ampoule, by heating the tip in a very hot oxygen and

natural gas flame of a hand torch. These ampoules and five or six

glass beads of two to three millimeters diameter were placed in a soft

glass test tube having a diameter of fifteen millimeters and a con-

striction made in the test tube. Exactly 5.0 ml. of a freshly stand-

ardized solution of sodium iodide in dry acetone was placed in the

tube. The soft glass test tube was then placed in an ice bath and

sealed with a hand torch. A scratch was made around the tube. The

tubes were then allowed to equilibrate in the constant temperature bath

and after equilibration, the ampoules were broken by vigorous shaking

of the sealed tube. This gave instantaneous intimate mixing and was

taken as zero time for the subsequent calculations. When it was desired

to take a reading as the reaction progressed, the tubes were broken into

the ice, hydrochloric acid mixture by touching the scratch on the tube

with a hot pyrex rod. The iodide ion concentration was then determined

as described above.

When the reaction temperature was low as compared to the boiling

point of the liquids used, the procedure was changed in order to make

the manipulation easier. The halide was weighed into a 100 ml.

volumetric flask. The flask and one containing the sodium iodide in

acetone were placed in a constant temperature bath. After equilibration,

enough of the second reactant was added to bring the liquid to the mark.

At various times, 5 ml. of liquid were withdrawn and analyzed. 9

The kinetics of the reaction of methoxide ion with the methylene

halides was followed in all cases, except that of fluorobromomethane, by

quenching the reaction in excess standard hydrochloric acid and titrating

the excess acid with standard sodium hydroxide solution. It has been

reported (17) that a great many indicators are interfered with in

acid-base titrations in aqueous methanol solution. In order to determine

which indicator would serve the best in this determination, standard

hydrochloric acid was titrated both in the presence and absence of methanol,

using various indicators. It was found that rosolic acid gave the best

end point with no detectable interference by methanol, and it was adopted

for use in the titration.

Fluorobromomethane, when reacted with methoxide ion, was found

to proceed only to the point where the bromine had been replaced. When

the reaction was quenched and the acid titrated, no definite end point

could be reached until a great excess of sodium hydroxide had been used.

In one case, for instance, 1.4550 meq. of acid was added and the titra-

tion required 2.0057 meq. of sodium hydroxide to reach a fair end point.

Since the reaction mixture had contained 1.4291 mmoles of fluoro-

bromomethane and 1.3130 meq. of sodium methoxide, it can be seen that

if the reaction had gone to completion the most sodium hydroxide which

would be necessary to neutralize the hydrochloric acid is 1.4550 meq.

It appeared, then, that hydrolysis of the fluoromethyl methyl ether was taking place during the titration. In order to prevent this from

occurring the methoxide ion was titrated directly with p-toluenesulfonic

acid in methanol with bromphenol blue as indicator. The sample to be

17. J. Stastny, Chem. Zentr., (1942) II, 1039. 10

titrated was poured into 5 ml. of methanol, which had been cooled in a

dry ice-acetone bath, and titrated while still very cold. Near the end

point the solution went from blue to green but the titration was con-

tinued until a definite, long-lasting yellow color was obtained.

In all of the cases where sealed tubes were used , the reactants

were placed in the tube at room temperature and the reactions were run

at other temperatures, thus making it necessary to apply corrections to

account for the expansion of the liquids. For the most part the densi-

ties used were obtained at two temperatures using pycnometers of standard

type, and plotted in "Figure 1", assuming a straight line relationship,

to obtain values at other temperatures. It was necessary to use a

special technique for those compounds which are gaseous at the temper-

atures under consideration, i.e., methyl bromide and fluorobromomethane.

A one-milliliter graduated pipette was sealed at the bottom marker and

evened at this marker on the inside so that the calibrations were still

accurate. When various amounts of water were placed in the pipette

using a one milliliter hypodermic syringe, it was found that the volume

read on the pipette was within 0.02 of that added. This gives a

maximum error of two per cent in the densities to be obtained, and

since the volumes used in the calculations are rounded off to 0.01 ml.

this error is negligible. The empty pipette was weighed, a small

amount of the liquid to be tested was inserted with the aid of a hypo-

dermic syringe, the pipette sealed at the top and reweighed. The pipette was then placed in a constant temperature bath until the volume had reached a constant value. Using this value of the volume, and the 11

known weight and neglecting the volume which had become a gas, the

following densities were obtained:

Compound Wt. Volume at Density at 20.3 °C. 35°C. 50 °C. 20.3 °C. 35°C. 50°C.

0.9111 .558 .572 1.633 1.593 CH3Br .570 1.87 1.78 CH3FBr 1.0147 .542

The values used for the densities of acetone and methanol were taken

from Beilstein (18) and the best straight line plotted through them for purposes of extrapolation. The values for methylene chloride were

taken from the work of Timmermans and Hennault-Roland (19). The values for methyl iodide and ethyl bromide were calculated, using the equation from the "International Critical Tables" (20). A summary of the densities used in all subsequent calculations may be found in Table 1 of the Appendix.

In order to obtain the initial molar concentrations it was

assumed in all cases, that with the small concentrations of halide used, the halide and acetone or methanol added ideally with no change

in total volume. The volume of acetone at the reaction temperature was

18. Beilsteins Handbuch der Organischen Chemie, Band I Verlag von Julius Springer, Berlin, Germany, 1918, pp. 273 and 635. 19. J. Timmerman and Mme. Hennault-Roland, J. Chim. Phys., 29, 529 (1932). 20. International Critical Tables, Vol. III, 1st ed., McGraw-Hill Book Co., Inc., New York, N. Y., (1926), p. 27. 12

obtained by using the known densities to convert the known volume at room temperature, assumed to be 25 °C., to the desired temperature. The volume of the methylene halide was obtained from the known weight and density at the reaction temperature. The two volumes were then

added together to obtain the total volume of solution. The rate constants for the iodide ion in acetone reaction were obtained using the normal second order and those for the methoxide ion reaction were obtained using a modified version of the second order rate equation. Since the calculated rate constants

for bromoiodomethane, methylene iodide and ethyl bromide were not constant 1 , initial rate constants were obtained by plotting the concentration of halide used versus time and obtaining the initial slope. The rate is then calculated from the equation; k = dx dt Y-

All of the compounds containing iodine were found to be extremely light-sensitive. In order to prevent complications in the kinetics arising from decomposition it was necessary to make all runs involving these compounds in an absence of light. This was accomplished in two ways. Those runs which were made in sealed tubes were placed inside of a covered metal can which was baffled so as to allow free circulation of the water of the constant temperature bath but to prevent light from reaching the reactants. On the other hand, low-actinic glassware was used for those runs made in flasks.

1 See p. 54, 72, and 82. 13

Preparation and Purification of Reagents

Acetone.- Commercial acetone was dried by the method of Conant and

Kirner (12) by refluxing over calcium oxide and potassium permanganate for twenty-four hours. The acetone was fractionated through a vacuum-jacketed column of approximately 2 cm. inside diameter, one meter long, packed with three-eighths inch glass helices.

The fraction boiling at 56.0 °C. at a barometric pressure of 740.2 mm of mercury was collected for use. This method does not give a completely dry solvent, however, since the same procedure is always used it is suitable for our purposes.

Methanol.- Commercial methanol was first dried by shaking constantly with anhydrous calcium sulfate for one hour and then allowed to stand overnight. The methanol was then distilled from the calcium sulfate.

To the distillate was added 8 grams of magnesium metal turnings (21) and the mixture allowed to react, under reflux, until all of the magnesium had disappeared. The dry methanol was distilled from the and stored in two and a half liter bottles with rubber -gasketed tops and then a length of Gooch tubing was stretched over the cap and neck of the bottle to seal it.

21. L. F. Fieser, Experiments in Organic Chemistry, Part II, 2nd ed., D.C. Heath and Co., New York, N. Y., (1941), p. 359. 11+

Sodium Iodide.- Baker's Analyzed, C.P., sodium iodide was found to be

free of iodate and was dried under vacuum at 97°C. for two hours. A

solution of approximately 0.025 N sodium iodide in dry acetone was prepared and stored in a brown bottle. The solution was standardized

immediately before use in each subsequent run.

Potassium Iodate. - Baker's Analyzed, C.P., potassium iodate was used

as a primary standard to prepare a solution exactly 0.0100 N as an oxidizing agent.

Sodium Methoxide.- Freshly cut sodium was added to a flask of methanol in a dry box through which dry nitrogen was passing constantly. After the sodium had completely dissolved, the solution was stored in the

same manner described for the dry methanol. The solution was approxi- mately 0.25 M and was standardized before use.

Standard Sodium Hydroxide.- C. P. sodium hydroxide was dissolved in boiled, free, distilled water and stored in a polythene bottle, with the air vent protected by a soda lime drying tube. The solution was standardized against potassium acid phthalate using rosolic acid as indicator.

Standard Hydrochloric Acid.- Concentrated, C. P., hydrochloric acid was diluted with distilled water to give a solution approximately

0.05 N which was then accurately standardized with standard base using rosilic acid as indicator. 15

Methylene Chloride.- Eastman Kodak's White Label methylene chloride was fractionated through a Todd Column having an inside diameter of

1.2 cm., a length of three feet and packed with one-eighth inch, single turn, glass helices. The fractionation was carried out under nitrogen and the fraction boiling between 39.5° and 40.0 °C. at 738.7 mm of mercury was collected and stored in brown glass bottles under nitrogen.

Chlorobromomethane.- Dow Chemical Company's chlorobromomethane was fractionated under nitrogen and the fraction between 67.8° and 68.0°C. at 744.0 mm barometric pressure was used. The compound is stored in the same manner as the methylene chloride.

Methylene Bromide.- A sample of dibromomethane from Columbia Organic

Chemicals Company was fractionated under nitrogen and the fraction boiling at 97 °C. at a barometric pressure of 733.3 mm was collected and stored in the usual manner.

Bromoiodomethane.- Bromoiodomethane has been prepared by refluxing 30 ml. of dibromomethane and 22.5 grams of sodium iodide in 50 ml. of acetone for four hours. The acetone was removed by distillation and found to contain about ten milliliters of a methylene halide which separated upon addition of water. This halide did not change color upon being exposed to light and is therefore believed to be the unreacted dibromo- . The was removed from the crude product remaining in the pot, by filtration and the filtrate was washed with solution and dried over calcium sulfate. This product was then fractionated at 45 mm pressure. The fraction boiling 16

at 59° to 60 °C. at this pressure was found to be very light yellow, n25 = 1.6353. The molar refraction calculated from this data is 26.8 and that from the atomic refractions is 27.4

Chloroiodomethane.- A sample of chloroiodomethane prepared by the action of sodium iodide in acetone on , as described above for bromoiodomethane, was fractionated through the Todd still under nitrogen. The fraction boiling at 108 °C. and a barometric pressure of 740.6 mm was collected, and stored under nitrogen in a brown glass bottle. The sample still maintained a slight violet coloration, undoubtedly due to the presence of minute amounts of iodine. No attempt was made to remove this color since the quantity present was so small. It was found that the compound was very light and heat sensitive, decomposing to liberate more iodine, as evidenced by the rapid deepening of color upon exposure to sunlight or excessive heat. The at twenty degrees was found to be 1.5812 and gives a molar refraction of 23.6 as compared to 2L..5 from the tables. We have also succeeded in preparing the compound using chloro- bromomethane and sodium iodide.

Fluorobromomethane, Fluoroiodomethane, and Fluorochloromethane.- Various attempts have been made to prepare chlorofluoromethane, bromofluoro- methane and iodofluoromethane from the corresponding dihalo and iodohalo compounds using mercuric fluoride as the fluorinating agent. In all cases 1.5 mole of mercuric fluoride was added in small amounts to

1.0 mole of the methylene halides with constant stirring. In the case of the dibromo and dichloro compounds there was apparently no reaction 17

in the cold so the mixture was heated to reflux. There was no product

collected in the dry ice-acetone traps from the dichloromethane and

three to four milliliters of a colorless liquid was obtained from the dibromomethane. This was not a sufficient amount to analyze or to use

in the kinetic runs. The diiodo compound reacted violently when heated and after two attempts, it was possible to isolate, by extraction with ether from the mercuric iodide which was formed, what was hoped to be

iodofluoromethane. Since the iodo compound reacted so well, it was hoped that the iodohalo compounds would also react. The iodochloro

compound gave about six to eight milliliters of a colorless liquid when the mixture was allowed to react at reflux with constant stirring.

Under these same conditions the iodobromo compound gave none of the

desired product. In the latter case, the mercuric fluoride changed from yellow to an off-white color, clearly indicating that a reaction had taken place. The small amounts of material obtained above were lost

in handling and further attempts to prepare larger amounts by the same procedure were entirely unsuccessful.

One mole of methylene chloride was stirred with 0.33 mole of antimony trifluoride and 0.33 mole of bromine at reflux for approxi- mately forty hours without success. The same reaction was repeated in a bomb at 100 °C. for approximately fifty-five hours again without success. To this reaction mixture was added about 2.5 grams of aluminum

chloride in the hopes that it would act as the antimony pentahalide

catalyst and was heated in the bomb again for approximately fifty-five hours at 100 °C. Still no product was obtained. The above was repeated 18

in a bomb, which could be shaken and heated to 120 °C., without success.

Antimony pentachloride was prepared by passing chlorine through molten antimony trichloride, and vacuum distilled, b.p. 86-87°C. at 30 to 33 mm pressure. This was used as catalyst in the reaction between bromo- chloromethane and antimony trifluoride, still with no results.

An attempt was made to react bromoform and antimony trifluoride in the presence of bromine to obtain the dibromofluoromethane after the method of Swarts (22). The reaction apparently began immediately since a good deal of heat was liberated at the outset and this was removed with cold water to moderate the reaction. However, the purification gave a product which did not correspond to any of the desired products or reactants. It is thought that this is possibly due to a reaction between the product and the thiosulfate which was used to remove the excess bromine.

It was found possible to prepare dibromofluoromethane by heating

1.0 mole of bromoform with 0.5 mole of mercuric fluoride. Thirty five grams of crude product having a boiling point of approximately 64 °C. were obtained. The attempt to reduce this with zinc in 70 per cent aqueous ethyl alcohol gave extremely small yields of a colorless liquid. It was found possible to prepare diiodofluoromethane by the same method and the product was found to have a boiling range of 78.0-78.5 °C. at

65 mm of mercury pressure which gives a boiling range of 140°-150 ° through extrapolation to atmospheric pressure. The index of refraction

22. F. Swarts, Bull. Acad. Roy. Belg., 113 (1910). 19

at 25 °C. found was 1.6365, density was 2.986 and the calculated molar refraction using this data was 34.3 as compared to 33.4 from the tables. Ruff and co-workers (23) have prepared what they believe to be diiodo- fluoromethane from the reaction of iodoform 2 mercurous fluoride and calcium fluoride and found the boiling point to be 100 °C. Since the boiling point of Ruff's compound is not that which would be predicted and since our compound gives good agreement with the calculated molar refraction as well as a boiling point closer to that expected, it is felt that our compound is actually the compound sought. However, reduction of the product, which we have obtained, with sodium arsenite gave none of the desired product. This may be due to the fact that an excess of arsenite was used which may have reduced the fluoroiodomethane to methyl fluoride.

It was finally found possible to prepare fluorobromomethane by decomposing silver fluoroacetate in the presence of bromine but this decomposition in the presence of iodine, in the dry, proceeded explosively.

In carbon tetrachloride, if any reaction occurred, it was impossible to isolate the fluoroiodomethane.

One hundred grams of Monsanto's was dissolved in 100 ml. of water and to this was added 170 grams of silver nitrate in 150 ml. of water. The resulting mixture was digested by boiling for approximately fifteen minutes, cooled and filtered. The residue was dried under vacuum at 75 °C. for approximately fifteen hours

23. O. Ruff, O. Bretschneider, W. Luchsinger and G. Miltschitzky, Ber. 69, 299 (1936). 20

and the crude, dry silver fluoroacetate was obtained in 89 per cent yield. The silver fluoroacetate was mixed with 100 ml. of bromine in 200 ml. of carbon tetrachloride. The reaction began almost immediately,

without heating, as evidenced by a vigorous evolution of gas. The

gases were collected in a dry ice-acetone trap and fractionated through a glass helices packed column having an inside diameter of 8 mm, a length of two feet and equipped with a water jacket so that the column could be maintained at a temperature near the boiling point of the desired compound. About thirty-five grams (27% yield based upon sodium

fluoroacetate) of fluorobromomethane with a boiling range of 18-20 °C. was obtained and stored in the freezing compartment of the refrigerator in a glass stoppered Erlenmeyer flask.

Methylene Iodide.- Eastman Kodak White Label methylene iodide was washed with sodium thiosulfate, dried over calcium sulfate and fractionated at 6 to 7 mm pressure. The fraction boiling at 65-66°C. was collected and stored under nitrogen in a brown glass bottle.

Methyl Bromide.- Matheson's methyl bromide was used without further purification.

Ethyl Bromide.- Merck reagent grade ethyl bromide was twice fractionated by Mr. W. H. Brader. The fraction boiling at 38.4 °C. was collected and stored in a brown glass bottle in the freezing compartment of the refrigerator. This fraction was found to have a refractive index of 1.4239. 21

CHAPTER III

DISCUSSION OF RESULTS

In Chapter I, the effects of substituents upon SN2 reactivity were discussed in a rather general way. In order to understand these effects better, it might be advantageous to consider the energies involved in this type reaction. If we take the case of a methyl halide reacting with some nucleophilic reagent Y;

H H H H H \/ Y+ H C- X Y - C - X Y - C + X

as Y approaches the methyl halide molecule, energy is utilized to over- come the repulsions of the halide and the group Y. As a bond begins to form between Y and carbon, energy is obtained from the bond formation and at the same time the C-X bond begins stretching. The stretching of the C -X bond requires that energy be supplied. We have now reached the transition state and thus it becomes apparent that the activation energy is the algebraic sum of these energies.

From a consideration of the foregoing discussion we may conclude that the speed of an SN2 reaction may be increased by decreasing the energy required to bring the attacking reagent to the reaction center, by decreasing the bond strength of the displaced group or by increasing the energy of the bond being formed. 22

Since, with the exception of Iodine, all of the halogens are more electronegative than carbon it is reasonable that the inductive effect will decrease the electron density upon the central carbon atom. This, it seems, should decrease the energy required for the approach of the nucleophilic species. However, it is postulated in this thesis that this same inductive effect also causes an increase in the strength of the bond being broken and also of the bond being formed.

An examination of Table 2, shows that in every case the dihalo- methanes react slower than the halomethanes and have higher energies of activation. It would seem that the supposition which has been made concerning the increase in strength of the bond being broken is valid and is the predominant energy factor in the decreased reactivity of the methylene halides. 23

Table 2. Summary of Results for the Reaction

RX + Y - RY + X -

RX Y - k x 105 1 mole -1 sec -1 Abs. Rate Ea. L Ha ASa 20.3 °C. 36 °C. 50 °C. Kcal e.u.

CH2C12** I- *.842 .211 29.0 +5.7 +.017 +.003

CH2C1Br I- 7.12 47.33 220.7 +.247 +.50 +2.4 21.2 -5.3

olislI I- .155 -- +.005 CH2BrF I- 55.58 1350 +4.55 +18 19.6 -6.6

CH2Br2 ** I- 2.073 16.00 69.08 +.06o +.33 +.92 21.6 -6.4

CB2BrI I- 5.66 116 + .57*** 18.8 -13.9

CH3CH2Br I- 123.1 1748 +12.3*** + 69 15.9 -17.5 OH2o12** 0o113 .0285 +.0010_ CH2C1Br OCHS .0413 .378 2.355 +.0010 +.0017 +.022 25.1 -2.2 CH2C1I OCH3 .0863 4.417 +.0038 +.167 24.4 -3.1

CH2BrF 0CH5r 7.613 218.5 +.170 + 9.7 20.7 -6.8 CH2Br2** 0C H3 .00613 .06175 .3703 +.002 +.0012 +.017 25.4 -5.0 24

CH2BrI OCH3 .031 1.023 +.003*** +.027 21.6 -14.8

CH2I2* OCH3 .01225 .536 +.0012*** +.007 23.4 -15.0

479 CH3Br OCH 3 17.0710 +.58 +24 20.6 -5.6 clig oc 113 15.01 433.0 +.39 +8.7 20.7 -5.6 CH3CH2Br OCH3 1.33 47.2 +.02 +1.6 22.1 -5.5

* This value is for 60 °C. rather than for 36°C. ** Rate constants for these compounds contain a statistical factor of one-half. *** Deviation estimated for this value, all others are average deviations. 25

However, within the methylene halide series themselves, it is

found that other factors are operating. Considering the complete series

found in the reaction of the bromides with sodium iodide in acetone,

Table 3, it is found that the reactivities at 50 °C. relative to chloro- bromomethane are:

CH3CH2Br = 7.9, FCH2Br = 6.1, C1CH2Br = 1.00, BrCH2Br = .31,

ICH2Br = .54

It is seen that the effect of the second halogen reduces the reactivity

of the bromine atom in the order: Br> I > Cl > F This suggests that, although we have reduced the reactivity overall, the fact that the inductive effect of the halogens increases the ease

of approach of the nucleophilic reagent is of major importance. Thus, the fluorine atom which has by far the greatest inductive effect

increases the reactivity six fold with respect to chlorine. On the other hand, bromine, which is only slightly less electronegative than chlorine, changes the reactivity by a factor of only one-third with respect to chlorine.

Iodine offers an apparent anomaly, in that it would seem to have little or no inductive effect since its electronegativity is the same as that of carbon. Thus one might expect the energy required to break the bond to be lowered, giving a lower energy of activation and an increased rate of reaction. It is noted that the lowered energy of 26

Table 3. Alkyl Bromide Plus Iodide Ion

Alkyl k x 10 5 1 mole -1 sec -1 41S Group 20.3 °C. 50 °C. Kcal e.u.a

cH3cH2 123.1 + 12.3** 1748 + 69 15.6 -18.4

CH2F 55.58 + 4.55 1350 + 18 19.4 -7.3 eli2c1 7.212 + .247 220.7 + 2.4 20.8 -6.4

CH2Br* 2.073 + .060 69.08 + .92 21.4 -6.9

01121 5.66 + .57** 116 18.5 -14.9

CH3 too fast to measure with any accuracy even at 0 °C. The specific rate constant is of the order of .03 1 mole-1 sec -1

* Rate constants contain a statistical factor of one-half.

** Deviation estimated for this value, all others are average deviations. 27

activation is, indeed, realized and that the rate of reaction is greater than for the corresponding dibromomethane. It is also to be noted that the entropy of activation of this compound is twice that of any of the other halides. It can be reasoned that the entropies of activation of the other halogens are fairly close to the same value since even in the most favorable ) case, they can be determined with an accuracy of only

+ 2.0 entropy units. Thus it seems that in these cases the entropy does not play a large part in the rate of reaction but in the case of bromo- iodomethane, the extremely large negative entropy of activation plays a decided role in the rate of the reaction.

If we now consider the case of the iodohalomethanes, Table 4, we find the reactivities at 50 °C. relative to chloroiodomethane to be:

CH3I = 98, C1CH2I = 1.00, BrCH2I = .23, ICH2I = .12

It will be noticed that iodobromomethane has been included in this list even though it was found that methyl bromide reacts slightly faster than methyl iodide with sodium methoxide in methanol while chloroiodomethane reacts faster than chlorobromomethane by a factor of two. Thus, it is not known to what extent the rate controlling step of the reaction consists of displacement of bromide and to what extent it is displacement of iodide.

However, the data which are of value, support the conclusions drawn above for the bromocompounds; i.e., iodine slows the reaction more than chlorine and the second halogen causes a decided decrease in reaction rate as compared to the methyl halide.

1 See p. 43. 28

Table 4. Alkyl Iodide Plus Methoxide Ion

Alkyl k x 10 1 mole -1 sec -1 L Ha, iSa Group 20.3 -C. 50°C. Kcal e.u.

CH3 15.01 + .39 433.0 + 8.7 20.5 -6.2

CH2C1 .0863 + .0038 4.417 + .167 22.3 _9.7 CH2Br** .031 + .003*** 1.023 + .027 21.3 -15.7

CH2I* .01225 + .0012*** 0.536 + .007 23.0 -11.5

* Rate constants contain a statistical factor of one-half. ** Due to reasons given in the discussion, it is uncertain whether this compound belongs in this series or not.

*** Deviation estimated for this value, all others are average deviations. 29

Consideration of the data for the reaction of the bromohalo- methanes with sodium methoxide in methanol (Table 5) leads to the con- clusion that reactivities at 50°C. relative to chlorobromomethane is:

CH3Br = 203.0, FCH2Br = 92.7, CH3CH2Br = 20.0, C1CH2Br = 1.00,

BrCH2Br = .16.

Again it is seen that the second halogen decreases the reactivity in the order: Br> Cl> F , in agreement with the previous results.

It was thought that it would be. of interest to compare the rate of reaction of ethyl bromide to that of the bromohalomethanes since the size of the methyl group is reported to be between the size of the chlorine atom and the bromine atom (24). We should thus be able to obtain a semi-quantitative idea of the importance of steric hindrance in this type reaction. However, not only is there a steric effect but an inductive effect; i.e., electron donation by the methyl group, which is working in exactly the opposite direction to the inductive effect already described for the halogens. The ethyl bromide was found to react at a rate one-tenth that of the methyl bromide, a fact which might be due either to the higher electron density at the reaction

24. H. Gilman, Organic Chemistry, Vol. I, 2nd ed., New York, John Wiley and Sons, Inc., (1948), p. 362. 30

Table 5. Alkyl Bromide Plus Methoxide Ion

Alkyl 6. 1Ia 6,Sa Group k x 105 1 mole -1 sec -1 Kcal e.u. 20.3 °C. 50°C.

0113 17.07 + .58 479 + 2!. 20.3 -6.6

01-12F 7.613 + .170 218.5 + 9.7 20.5 -7.4

01130112 1.33 + .02 47.2 + 1.6 22.8 -3.2

CH2C1 .0413 + .0010 2.355 + .022 24.7 -3.4

CH2Br* .00613 + .0002 .3703 + .017 25.1 -5.9

* Rate constants contain a statistical factor of one-half. 31

center or to the steric hindrance of the methyl group. The difference

in the two entropies of activation for this compound does not lend

itself to easy interpretation. However, it seems that, since the ethyl

bromide reacts at a rate twenty times that of chlorobromomethane, it

cannot be the steric factors which are of major importance even though

they are undoubtedly exerting some influence.

It has been called to our attention by Dr. L. D. Frashierl that

he has found lithium bromide to be a weak electrolyte in acetone and

therefore we might expect the same of sodium iodide. This would lead to

a rate equation of the type:

dx/dt = + k2 iNa.13} (RX)

rather than the type in which the sodium iodide was considered to be entirely ionic. It has been found by Kraus and Bray (25) that sodium iodide has a dissociation constant of 39 x 10 -4 at 18°C. This gives a concentration of iodide ion which is 0.0076 molar for a 0.025 molar solution of sodium iodide at 18 °C. This value of course will be slightly greater at our reaction temperatures. If one considers the constancy of the calculated rate constants for the sodium iodide reactions, it seems quite likely that the rate of reaction of the iodide ion and the sodium iodide ion pair must be approximately the same, perhaps differing by as much as two or three fold.

1 Private communication.

25. C. A. Kraus and W. C. Bray, J. Am. Chem. Soc., 35, 1315 (1913). 32

Consideration of a plot of the logarithm of the rate constant versus the reciprocal of the temperature, "Figure 3", indicates the agreement in the rate constants and the constancy of the activation energy over the temperature range studied. The line for ethyl bromide and iodide ion is very encouraging since the first two points are our data and the third point is taken from the data of Dostrovsky and Hughes (26).

If one draws a curve of the logarithm of the rate constant for the reaction of sodium methoxide with the bromohalomethanes versus those for the sodium iodide with the same compounds, one obtains a curve which closely approximates a straight line. We find that the equation of this line is similar to a Hammett type equation:

log kNaOCH3 = m log kNal and thus there is a constant relationship between the rate of reaction of sodium methoxide and sodium iodide with the methylene halides.

However, as evidenced by the fact that ethylbromide is so far removed from this line, one must use caution in attempting to draw conclusions outside of the families studied.

26. I. Dostrovsky and E. D. Hughes, J. Chem. Soc., 161 (1946). 33

CHAPTER IV

CONCLUSIONS

It has been found in this investigation that the replacing of

a hydrogen in a methyl halide by a halogen atom decreases the reactivity

of the halide in its reactions with iodide ion in acetone and methoxide

ion in methanol.

In any of the halide series studied in the present work the order

in which the second halogen decreases the reactivity of the halide is:

Br> I> Cl> F

In the reactions of the various bromohalomethanes with iodide

ion it is the activation energy which is the dominant factor controlling

the change in the rate of reaction except for bromoiodomethane in which

case it is the entropy which is most important. In the reactions of the iodohalomethanes and of the bromohalo- methanes which were run with methoxide ion, it is the activation energy which is the dominant factor controlling the change in the rate of reaction.

It is impossible to place the reaction of bromoiodomethane with methoxide ion in a series because the true nature of the rate controlling step of the reaction is unknown. 3!

CHAPTER V

RECOMMENDATIONS

It is felt that it would be of considerable value to ascertain whether the steric factors for these compounds are as indicated by the

calculations. This could be done by using triethylamine and quinuclidine since both compounds are quite similar, the only difference being that

the three ethyl groups are tied back in quinuclidine. Thus, steric factors should play a very important role in the reactions of triethyl- amine and a much less important part in the reactions of quinuclidine. It might also be of interest to run iodohalomethanes with radioactive iodine so that it would. be possible to find the effect of bromide upon the reactivity of iodide. This should be possible since the rate of reaction of iodide with the bromide is now known and could also be followed simultaneously with the exchange reaction. 35

APPENDIX A

SAMPLE CALCULATIONS

36

I. Derivation of rate equations. For reactions of type:

CH2XY + NaI CH2XI + NaY

See any good text on Kinetics such as Laidler (27).

k = 2.303 to g b (a-x) t (a-b) a (b-x)

II Derivation of the rate equation for second order reactions of the

type:

CH2XY + OCH3 CH2X0CH3 Fas CH2 (OCH3 ) (a-x) (b-2x)

dx at = k [CH2XY) LOCHd

Let x = amount of product at time t

a = initial concentration CH2XY

b = initial concentration OCHI

dx aT = k (a-x) (3-2x)

kdt - dx (a-x) (b-2x)

Solving by partial fractions.

1 kt + C _ in a-x 2a-b b-2x at t = 0, x = 0

C = 1 in 2a-b a

Therefore

kt = In (a-x) 1 in a 2a-b (b -2x)2a-b

kt = 2.303 b (a-x• ) 7a7:67) log a (b-2x) 37

III Calculation of Alla from absolute rate equation

k = KT - AS e e h RT

K = Boltzmann's constant

h . Planck's constant

In k = In KT Ha + AS h RT R

Using the two sets of data we have available:

AHEL = In ET' - In ki + RT

&Ha = in KT2 - In k2 + LNS RT2

AHrt - Ha = R Eln Ti - in T2 - In ki + In k21 T1 T2

4111a = 2.303 RT1T2 [log k2 - log T2 T2 - T L ki Ti

IV Calculation of LIS a starting with absolute rate equation: Ink = In KT + Asa hh RT

ASa = 2.303 R log k - 2.303 R log KT + h T

V Propagation of errors. For method see Goodwin's Precision

Measurements (28).

Let ZN = Numerical precision measurement of the indirectly determined

quantity. 6 = Numerical precision measurement of the experimentally determined quantity.

38

Error in rate constant due to deviation in temperature. Since the constant temperature baths are regulated to + 0.05, this is the deviation used. Starting with absolute rate equation,

Q Ha Asa RT R k = --KT e e h

LISa R ET L1Ba e:Sk _ K e e (1 + R ) 6T

Let b. T = precision measure of k due to deviation in T.

AS -4121-1, R RT .6. = __6.1 = Ke e (1 .i. 4 11a ) g T 6 T S T - h R T For the reaction:

CH2FBr + NaI CH2FI + NaBr at 50 °C. -6.6 1960o. 1.987 1.987 x 323 ZN = 1.38 x 10 -16 e 19600 T a 1.987 x 323)(+ 0.05) 6.624 x 10-37

DT = 2.083 x 1010 x 1.9835 x 10 -15 (31.54) (+ 0.05)

6T = (4.13 x 10 - 5) (1.577) = 6.51 x 10-5

k = 1350 x 10 -5

Error in rate constant due to inaccuracy in time. k 1 in b (a-x) or k - 1 in b(a-x) t(a-b) a (b-x) t(2a-b) a(b-2x)

_ 1 b(a-x) In Sk 1 , b(a-x) St t2 (a-b) a(b-x) gt t2 (2a-b) a(b-2x)In

39

Let At = p.m. of k due to deviation in t.

A t = r2 . 303 log b(a-x) St a(b-x) l_t2(a-b)

46, = [2.303 b(a-x) log t t2 (2a-b) a(b-2x)

For short time intervals the time was read to the nearest second but, due to inability to determine exact zero and end time, the precision

measure is taken as + 5 secs. CH2FBr + NaI CH2FI + NaBr at 50 ° C.

41 _ 2.303 log .0373 x .3211 t = (+ 5) 1002 (.2971) .3344 x .0240

■ = + 6.7 x 10 -5 (k = 1350 x 10-5) .n t For long time intervals in which t is read to the nearest minute giving an error in t of + 30 secs.

CH2C1I + 2-0CH3 CH2(OCH3 )2 + Cl- + I - at 50°C.

2.303 .2578 x .1873 t = log (+ 30) 371402 (.2470) .2524 x .1276

At -; .003 x 10 -5 1 mole-1 sec -1 (k = .0863 x 10-5)

Error in rate constant due to variations in concentration of nucleophilic reagent. k = 1 In b(a-x) t(a-b) a(b-x)

bk 1 in b(a-x) x sb t(a-b) 2 a(b-x) t(a-b) (b-x)

40

Let 43 = P.M. of k due to deviation in b

[2.303 log b(a-x) 46b cSb 2 t(a-b) a(b-x) t(a-b)(b-x)

Assuming an error + 0.1 ml. of KI03 in the burette reading during the

titration for the reaction:

CH2BrF + Nal ----ow CH2FI + NaBr at 50°C.

The error in b = + 0.000094 12.303 .0373 x .3211 41b log .0133 100(.2971)2 .3344 x .0240 100(.2971)(.0240)]

(0.000094)

41b = [04550 - .01865 ] (0.000094)

.02685 x 0.000094 = + 0.18 x 10 -5 (k = 1350 x 10 -5 ) b

Error in the rate constant due to variation of alkyl halide concen-

tration.

(Sk [2.303 log b(a-x) bx a(b-x) o a • t(a-b)2 at(a-b) [2.303 log b(a-x). - bx ila a(b-x) t(a-b)2 b at(a- ) Sa

Assuming weighing error of + 0.2 mg for CH2BrF, this gives error of

+ 0.0000044 in a.

z 2.303 2 log .0373 x .3211 .0373 x .0133 100(.2971) .3324 x .0240 100 x .3344 x .29713

(+ 0.0000044) Aa 0.04545 x 0.0000044 = 0.02 x 10 -5 (k = 1350 x 10-5) 41

Error in the rate constant due to variations in determination of the quantity of material at time t: 6k = t(a-x)(b-x)

6x 41x = [(a-1)0(10-4.] Assuming an error of + 0.1 ml. KI03 in the titration gives an error of + 6.000094 in x

0.0373 41x = 0.000094 = + 0.54 x 10-5 [100 x .3211 x .0240]

k = 1350 x 10 -5 A 2 „ 2 2 Total error in k = V(6.; 21t2 "a + "I) + 4x

2S = 10-5 V(6.512 + 6.72 + 0.022 + 0.182 + 0.542

10-5 = + 9.36 x k = 1350 x 10 -5 The error is seen to be slightly more than half the observed average deviation of + 18 x 10-5, showing that the indeterminate errors are of considerable magnitude.

Error in LIHEL due to deviation in kl.

( 11 Ha) - RT1T2 ka. ki(T2-T)

For CH2FBr + NaI at 20.3 °C.

Let .6k = P.M. ofAlEla due to deviation in k. 1

2N = x 293.5 x 323.2 .] (+4.55 x 10-5) k1 55.58 x 10 - 5 x 29.7 .52 Kcals Akl = &Ha = 19.6 Kcals

Error in Alla due to deviation in 12.

For CH2FBr + NaI at 50°C.

AHa) _ RT1T2 k2(T2-T1)

LI k 2 = [1.987 x 293.5 x 323.2 I\ (+ 18 x 101 5) 1350 x 10-5 x 29.7

d k2 = .085 Kcals

AHa = 19.6 Kcals

Total error in 4YRa = \d/L1 ki2

v/.522 + .0852 + 0.53 Kcals

L4 Ha 19.6 Kcals

Error in ASa due to deviation in Lid a .

S( &Sp) - 1 Ha) -

For CH2BrF + Na at 20.3 °C. Let P.M. of 21Sa due to deviation in Mai

1 A6H = (+ 520) 1 293.5 -

= + 1.8 e.u. ANHal _

ASa = - 6.6 e.u. 43

From the foregoing calculations we may conclude that the energies

of activation as calculated in this thesis have a precision of about +0.5

Keels and the entropies of activation about + 2 e.u.

It should be noted that the values, for those compounds for which

it was found necessary to obtain the rate constant graphically, are

probably in greater error than the others. This becomes obvious when

one considers the difficulty in obtaining the "best" initial slope for

the "best" curve. In this thesis it has been estimated that the

difficulty in obtaining the initial slope in these cases give an error

of + 8 to 10% in k. It has therefore been assumed that the error is

+ 10% for these values. Wi.

APPENDIX B

TABLES Table 6 Densities of Compounds Used

•

Densit Compounds 20C. 50'C. •0 C. •

Acetone .790 0.785 0.758 0.71+6 Methanol 0.791 0.763 CH2C12 1.27 1.25

CH2C1Br 1.94 1.888 1.84

CH2C1I 2.49 2.4o1+ 2.34 CH2BrF 1.87 1.78

CH2Br2 2.49 2.1+6 2.42

CH2BrI 2.70

CH2I2 3.32 3.24

CH3Br 1.675 1.593

CH3I 2.28 2.20

CH3CH2Br 1.46 1.40 46

Table 7

CH3I + NaOCH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N HC1 = 0.0569 N

Wt. Molarity Molarity Time Ml NaOH k x 104 CH3I CH3I NaOCH3 in per 25 1 mole -1 min-1 Min. ml HC1

.2443 .3388 .2544 30 7.70 89.50

.2286 .3176 .2549 120 15.00 90.51 .2087 .2907 .2554 180 17.70 88.90

.2165 .3014 .2554 300 23.40 88.72

.2303 .3200 .2549 464 26.90 92.93

.2204 .3063 .2549 1034 36.50 94.78

Blank of 5 ml NaOCH3 + 25 ml HC1 = 4.20 nil NaOH

25 ml HC1 = 45.9 ml NaOH

Average k = (90.05 + 2.31) x 10- 4 1 mole -1 min -1

-5 1 mole -1 sec-1 = (15.01 + 0.39) x 10 47

Table 8

+ NaOCH3 in Absolute Methanol at 50 ° C.

NaOH = 0.0310 N HC1 = 0.0589 N

Molarity Time Ml NaOH k x1104 Wt. Molarity CH3I CH I NaOCH3 in per 25 1 mole min-1 3 Min. ml HC1

.2403 .3218 .2444 2.07 12.10 2526

.2070 .2783 .2453 3.13 13.70 2571

.2322 .3110 .2444 4.03 16.65 2649

.2373 .3178 .2444 5.03 18.45 2553

.2375 .3178 .2444 6.05 20.48 2622

.2158 .2896 .2449 7.07 21.15 2687

.2370 .3174 .2444 8.07 23.95 2746

.2292 .3076 .2449 9.03 25.55 2929

.2391 .3203 .2444 15.95 30.70 2538

.2345 .3141 .2444 30 37.90 2639

Blank of 5 ml NaOCH3 + 25 ml HC1 = 6.03 ml NaOH 25 ml HC1 = 47.5 ml NaOH

Average k = (2598 + 52) x 10 -4 1 mole -1 min-1 = (433.0 + 8.7) x 10 -5 1 mole -1 sec-1 48

Table 9 CH2C1I + NaOCH3 in Absolute Methanol at 20.3°C. NaOH = 0.0310 N a HC1 = 0.0569 N b HC1 = 0.0582 N

4.7827 gms. of CH2C1I per 100 ml Solution Molarity of CH2C1I = 0.2710 Molarity of NaOCH3 = 0.2554

Time in M1 NaOH k x 104 Min. per 25 1 mole -1 min-1 ml HC1

0 a 4.70 1420 a 9.78 .548

7217 a 12.40 .549

10293 a 14.30 .511 12525 a 15.90 .502 18990 a 20.20 .505

25472 b 24.41 .505

36245 b 28.81 .492

45653 b 31.40 .490

Average k = (0.518 4- 0.023) x 10 -4 1 mole -1 min-1

. (0.0863 + 0.0038) x 10 - 5 1 mole -1 sec -1 a 25 ml HCl . 45.9 ml NaOH b 25 ml HC1 . 46.9 ml NaOH 49

Table 10 CH2C1I + NaOCH3 in Absolute Methanol at 50°C. NaOH = 0.0221 N HC1 = 0.0642 N

Wt. Molarity Molarity Time Vol. NaOH k x 104 CH2C1I CH2C1I NaOCH in per 25 1 mole -1 min-1 3 Min. ml HC1

.2429 .2622 .2578 129 21.20 25.8 .2426 .2619 .2578 199 26.13 27.4 .2445 .2639 .2578 390 36.80 28.8

.2434 .2628 .2578 559 41.10 25.7

.2338 .2524 .2578 619 42.70 26.5

.2499 .2693 .2573 1154 55.40 25.6

.2546 .2743 .2573 1357 58.60 25.4

Blank of 5 ml of Na0CH3 + 25 ml HC1 = 11.75 ml NaOH

Average k = (26.5 + 1.0) x 10 -4 1 mole -1 min-1 -5 1 mole -1 sec-1 = (4.417 + .167)x 10

25 ml HC1 = 72.6 ml NaOH 50

Table 11 CH2BrI + Na0CH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N a HC1 = 0.0569 N b HC1 = 0.0582 N

5.5832 gms. CH2BrI 100 ml of solution

Molarity of CH2BrI = 0.2528 Molarity of NaOCH3 = 0.2554

Time in ml NaOH Used CH2BrI k x 104 Min. per 25 ml moles/liter 1 mole-1 min-1 HC1

0 4.70 a

1419 5.32 a 0.0019 .210 4363 6.20 a 0.0047 .172

29887 14.90 b 0.0284 .177

45651 17.30 b 0.0353 .152

54933 18.10 b 0.0382 .140 61858 19.25 b 0.0419 .138

72191 20.48 b 0.0457 .135

93983 21.75 b 0.0496 .116 a 25 ml HCl . 45.9 ml NaOH b 25 ml HC1 = 46.9 ml NaOH k = 0.186 x 10 -4 1 mole -1 min-1

= 0.031 x 10 -5 1 mole -1 sec -1

This k was obtained by plotting the amount of CH2BrI used versus time, obtaining the initial slope of this curve and dividing this value by the initial concentrations to obtain the initial rate constant. See "Figure 4". Page missing from thesis 52

Table 12 CH2BrI t NaOCH3 in Absolute Methanol at 50 °C. NaOH = 0.0310 N HC1 = 0.0589 N 8.3815 grams of CH2BrI per 100 ml solution Molarity of CH2BrI = .3796

Molarity of NaOCH3 = .3574

4 Time in ml NaOH k x 10 Min. to back 1 mole -1 min-1 titrate

0 r 8.85 325 s 7.8o 6.65 780 6.82 6.33 1457 15.90 6.14 1810 19.70 6.25 2188 23.10 6.14 2893 27.82 6.13 3250 30.10 6.25 3629 31.90 6.19 4360 34.60 5.88 468o 35.90 5.91 5773 38.65 5.94 7212 41.10 5.84 r Used 35 ml HC1 s Used 30 ml HC1 53

All others used 25 ml HC1 25 ml HC1 = 47.6 ml NaOH Average k = (6.14 + 0.16) x 10 -4 1 mole -1 min-1 -5 1 mole-1 sec-1 = (1.023 + .027) x 10 54

Table 13 CH2I2 + NaOCH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N a HC1 = 0.0569 N b HC1 = 0.0582 N

6.4193 grns of CH2I2 per 100 ml of solution

Molarity of CH2I2 = 0.2396 Molarity of NaOCH 3 = 0.2554

Time in ml NaOH Used CH2I2 k x 104 Min. per 25 ml moles/liter 1 mole -1 min-1 HC1

0 4.70 a

1418 5.02 a 0.0010 .116 4362 5.60 a 0.0028 .107 7215 6.40 a 0.0053 .124

18986 9.80 a 0.0158 .150 29885 12.65 b 0.0214 .134 45649 15.10 b 0.0290 .126

72190 17.50 b 0.0364 .106

93982 19.10 b 0.0414 .096 a 25 ml HCl = 45.9 ml NaOH b 25 ml HCl . 46.9 ml NaOH k = 0.0735 x 10-4 1 mole-1 min-1

. 0.01225 x 10- 5 1 mole-1 sec-1 55

This k was obtained in the manner described at the bottom of

Table 11, with a statistical factor of one-half included.

See "Figure 5". 56

Table 14

CH2I2 + NaOCH3 in Absolute Methanol at 50 °C. NaOH = 0.0310 N HC1 = 0.0589 N

6.0721 grams CH2I2 per 100 ml of solution

Molarity of CH2I2 = 0.2267 Molarity of NaOCH3 = 0.3583

k x 104 Time in ml NaOH Min. to back 1 mole -1 min-1 titrate

0 q 8.70

1085 4.38 6.68

1440 7.6o 6.56

1808 20.40 6.44

252o 15.10 6.34

287o 17.40 6.44

3249 19.2o 6.45

398o 22.30 6.37

4300 23.7o 6.38

5393 27.12 6.37

6832 30.75 6.31 q Used 35 ml HC1 - in all other cases used 25 ml HC1. 25 ml HC1 = 47.5 ml NaOH

Average k = (6.43 + 0.08) x 10- 4 1 mole -1 min-1 . (0.563 4. 0.007) x 10 -5 1 mole-1 sec-1 57

Table 15 CH3Br + NaOCH3 in Absolute Methanol at 20.3 °C. NaOH = 0.0310 N HC1 = 0.0569 N

Wt. Molarity Molarity Time in M1 NaOH k x 104 CH3Br CH3Br NaOCH3 Min. per 25 1 mole -1 min-1 ml HC1

0.1582 0.2020 0.2575 17 5.55 104.5

0.1343 0.2801 0.2565 28 7.28 104.2

0.0911 0.1911 0.2511 59 8.50 106.6

0.0962 0.2014 0.2575 89 10.3o 99.4

0.0831 0.1743 0.2580 118 11.30 106.7

0.0647 0.1360 0.2586 717 20.50 98.6

0.0682 0.1334 0.2586 1248 24.30 96.8

25 m1 HC1 = 45.9 ml NaOH Blank of 5 ml NaOCH3 solution + 25 ml HC1 = 4.10 ml NaOH. Average k = (102.4 + 3.5) x 10 -4 1 mole-1 min-1 = (17.07 + 0.58) x 10-5 1 mole-1 sec-1 58

Table 16 CH3Br + NaOCH3 in Absolute Methanol at 50 °C. NaOH = 0.0310 N HC1 = 0.0589 N

Wt. Molarity Molarity Time in Vol. NaOH k x 104 CH3Br CH3Br NaOCH3 Min. to back 1 mole -1 min-1 Titrate

.1424 .2862 .3550 1.03 1.10 o 2600

.1112 .2243 .3563 2.00 3.70 n 2912

.0751 .1521 .3577 2.12 1.50 o 2803

.0685 .1390 .3584 3.05 2.80 n 2767

.0952 .1924 .3570 3.05 4.70 0 2680

.0781 .1582 .3577 5.07 7.10 o 2911

.2366 .4702 .3509 5.55 15.65 p 2995

.2139 .4267 .3523 7.33 18.95 o 3223

.0498 .1013 .3591 10.3 7.60 o 2979 n Used 30 ml HC1 o Used 35 ml HC1 p Used 25 ml HC1

Blank of 5 ml NaOCH3 + 3 5 ml HC1 = 6.50 ml NaOH

25 ml HC1 = 47.5 ml NaOH

Average k = (2874 ± 145) x 10- 4 1 mole -1 min-1 = (479 + 24) x 10-5 1 mole- 1 sec -1 59

Table 17

CH3CH2Br + NaOCH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N' HC1 = 0.0582 N 2.7804 gms of CH3CH2Br per 100 ml of solution Molarity of CH3CH2Br = 0.2528 Molarity of NaOCH3 = 0.2558

Time in Ml NaOH k x 104 Min. per 25 ml 1 mole-1 min-1 HC1

0 5.40

1086 12.70 7.82

1710 16.20 8.19

2507 19.25 7.94

3162 21.65 8.10

4004 23.80 7.91

5390 26.80 7.89

25 ml HC1 = 46.9 ml NaOH

Average k = (7.98 + .12) x 10 -4 1 mole -1 min-1

.02) x 10 -5 1 mole-1 sec-1 = (1.33 6o

Table 18

CH3CH2Br + NaOCH3 in Absolute Me0H at 50 °C.

NaOH = 0.0310 N HC1 = 0.0582 N

Wt. Molarity Molarity Time in Ml NaOH k x 104 CH3CH2Br NaOCH3 Min. per 25 -lmin 1 mole -1 ml HC1

.1403 .2434 .2488 20 10.00 291.8

.1478 .2555 .2483 49 15.60 273.5

.1514 .2617 .2483 78 20.10 275.4

.1700 .2933 •2479 137 27.90 270.2

.1296 .2249 .2493 198 27.55 281.9

.1448 .2508 .2488 243 31.30 282.o

Blank of 5 ml of NaOCH3 + 25 ml HC1 = 4.80 ml NaOH 25 ml HC1 = 46.9 ml NaOH

Average k = (283.2 + 9.4) x 10 -4 1 mole -1 min-1 -5 1 mole -1 sec-1 = (47.2 + 1.6) x 10 61

Table 19 CH3BrF + NaOCH3 in Absolute Methanol at 20.3 °C. p-Toluene Sulfonic Acid = 0.0550 N

Wt. Molarity Molarity Time in MI Para- k x 104 CH2BrF NaOCH3 Min. Toluene 1 mole -1 min-1 Sulfonic Acid

.2323 .4065 .2616 140 19.00 44.78

.1737 .3058 .2632 352 15.90 46.07

.1009 .1790 .2653 36o 18.75 47.31

.2303 .4030 .2616 480 12.00 44.55

Blank of 5 ml NaOCH3 = 23.85 ml of p-Toluene Sulfonic Acid

Average k = (5.68 + 1.01) x 10 -4 1 mole -1 min-1 . (7.613 + .17) x 10 -5 1 mole -1 sec -1 62

Table 20

CH2BrF + NaOCH3 in Absolute Methanol at 50°C.

p-Toluene Sulfonic Acid = 0.0550 N

Wt. Molarity Molarity Time in Ml Para- k x 104 CH2BrF NaOCH3 Min. Toluene 1 mole -1 min-1 Sulfonic Acid

.2063 .3473 .2473 6 18.40 1317 .2249 .3772 .2463 8 16.51 1334 .1961 .3301 .2473 lo 15.98 1373 .1961 .3301 .2473 15 14.30 1222 .2182 .3666 .2468 24 9.7o 1341 .1516 .2562 .2482 55 7.35 1484

Blank of 5 ml NaOCH3 = 23.65 ml p-Toluene Sulfonic Acid

Average k = (134.5 + 58) x 10- 4 1 mole-1 min-1

= (218.5 + 9.7) x 10 -5 1 mole-1 sec - 1 63

Table 21 CH2C1Br + NaOCH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N a HC1 = 0.0569 N b HC1 = 0.0582 N 3.8007 gms of CH2C1Br per 100 ml Solution

Molarity of CH2C1Br = .2937 Molarity of NaOCH3 = .2523

Time in Ml NaOH k x 104 Min. per 25 ml 1 mole -1 min-1 HC1

0 5.20 a

560o 8.60 a .253

8668 9.95 a .250 10907 10.78 a .238 17370 13.68 a .241

23853 17.70 b .252 34625 20.70 b .235

44033 24.60 b .26o

53317 26.45 b .250 60244 28.04 b .251 a 25 ml HC1 . 45.9 ml NaOH b 25 ml HC1 = 46.9 ml NaOH

Average k = (0.248 + 0.006) x 10- 4 1 mole-1 min-1 . (0.0413 + 0.0010) x 10 -5 1 mole -1 sec -1 64

Table 22

CH2C1Br + NaOCH3 in Absolute Methanol at 36 °C. NaOH = 0.0211 N HC1 = 0.0503 N

Wt. Molarity Molarity Time in Ml NaOH k x 104 CH2C1Br CH2C1Br NaOCH3 Min. per 20 1 mole -1 min-1 ml HC1

.1909 .28568 .1680 3233 20.10 2.25

.1956 .29272 .1680 4606 24.60 2.30

.1744 .26150 .1683 6073 26.58 2.31

.1815 .27162 .1680 8648 31.90 2.30

.1957 .27287 .1667 8658 32.75 2.25

.2079 .31040 .1677 10103 36.20 2.27

.1965 .29406 .1676 10430 35.30 2.20

.1959 .29317 .1676 11550 36.55 2.23

.150o .22530 .1687 11618 33.42 2.36

.1935 .28958 .1680 24798 43.90 2.07

Blank of 5 ml of NaOCH3 + 20 ml HC1 = 6.60 ml NaOH 20 ml HC1 = 47.7 ml NaOH

Average k = (2.27 + 0.04) x 10 -4 1 mole -1 min-1 = (0.378 + 0.0017) x 10-5 1 mole -1 sec-1 65

Table 23 CH2C1Br + NaOCH3 in Absolute Methanol at 50 °C.

NaOH = 0.0221 N HC1 = 0.0532 N

Wt. Molarity Molarity Time in Ni NaOH k x 104 CH2C1Br CH2C1Br NaOCH3 Min. to back 1 mole-1 min-1 Titrate

.2026 c .2977 .2525 120 5.70 e 14.14

.1772 b .2609 .2530 183 7.60 e 14.33 .1854 b .2729 .2530 672 27.70 e 14.21

.1422 a .2101 .1637 1147 28.40 a 14.09

.1.723 c .2541 .2535 1213 31.50 e 14.19

.1592 c .2348 .2535 1217 30.01 e 14.30

.2003 b .2943 .2525 1518 38.75 e 13.83

.1499 a .2215 .1637 1539 33.15 d 14.26

.1577 a .2326 .1634 2667 40.75 d 13.92

.1793 c .2639 .2530 4078 52.90 e 14.00

a Blank of 5 ml NaOCH3 + 20 ml HC1 = 9.87 ml NaOH b Blank of 5 ml NaOCH3 + 30 ml HC1 = 12.18 ml NaOH c Blank of 5 ml NaOCH3 + 30 ml HC1 = 12.10 ml NaOH d Used 20 m1 HC1 = 48.2 ml NaOH e Used 30 m1 HC1 = 72.2 ml NaOH Average k = (14.13 + 0.13) x 10- 4 1 mole-1 min-1 = (2.355 + 0.022) x 10 -5 1 mole -1 sec-1 66

Table 24

CH2Br2 + Na0CH3 in Absolute Methanol at 20.3 °C.

NaOH = 0.0310 N a HC1 = 0.0569 N b HC1 = 0.0582 N

4.8661 gms of CH2Br2 per 100 ml solution

Molarity of CH2Br2 = 0 . 2799 Molarity of NaOCH = 0.2541

Time in M1 NaOH k x 104 Min. per 25 ml 1 mole -1 min-1 HC1

0 4.90 a 5438 5.78 a .0713 10745 6.61 a .0709 28107 10.20 b .0714

53155 14.00 b .0774 70110 16.00 b .0760 91902 18.20 b .0746

129352 21.50 b .0726

a 25 ml HC1 . 45.9 ml NaOH b 25 ml HC1 = 46.9 ml NaOH Average k = (0.03675 + 0.0011) x 10 -4 1 mole -1 min-1 -5 1 mole -1 sec -1 . (0.00613 + 0.0002) x 10 Average k contains a statistical factor of one half. 67

Table 25 CH2Br2 + NaOCH3 in Absolute Methanol at 36 °C.

NaOH = 0.0211 N HC1 = 0.0503 N

Wt. Molarity Molarity Time in Ml NaOH k x 104 CH2Br2 CH2Br2 NaOCH3 Min. per 20 1 mole-1 min-1 ml HC1

.2425 a .27031 .1680 3229 11.40 .731 .2941 a .32656 .1674 4605 14.70 .750 .1881 a .21049 .1687 6071 13.10 .693

.2586 b .28770 .1673 7300 17.22 -737 .2442 a .27221 .1680 10100 19.85 .754

.2159 a .24113 .1683 11545 19.75 .729

.2245 a .25673 .1680 11613 20.10 .735 .2554 a .28469 .1677 16020 26.15 .678

.2594 b .28859 .1677 20273 28.60 .719

.3248 b .35996 .1671 20303 32.10 .723

.2488 b .27733 .168o 27508 32.2o .726

.2002 b .24593 .1683 40743 36.60 .777

a Blank of 5 ml of NaOCH3 + 20 ml HC1 = 6.60 ml NaOH b Blank of 5 ml of NaOCH3 + 20 ml HC1 = 6.70 NaOH

20 ml HC1 = 47.7 ml NaOH

Average k = (0.3705 + 0.007) x 10 -4 1 mole -1 min-1 = (0.06175 + 0.0012) x 10 -5 1 mole -1 sec -1 Average k contains a statistical factor of one half. 68

Table 26 CH2Br2 + NaOCH3 in Absolute Methanol at 50 °C. a NaOH = 0.0211 N HC1 = 0.0503 N b NaOH = 0.0211 N HC1 = 0.0532 N c NaOH = 0.0221 N HC1 = 0.0532 N

4 Wt. Molarity Molarity Time in Ml NaOH k x 10 CH2Br2 CH2Br2 NaOCH3 Min. to back 1 mole -1 min-1 Titrate

.2433 b .2666 .1630 1157 20.60 e 4.34 .2828 c .3087 .2512 1196 16.13 f 4.4o .2278 a .2500 .1674 1465 17.30 d 4.49

.2803 c .3059 .2512 1646 20.45 f 4.4o

.1634 a .1801 .1680 2891 20.10 d 4.32 .1822 c .2004 .2532 3886 27.80 f 4.73 .2366 c .2592 .2522 4336 35.65 f 4.59

.2776 a .3036 .1667 4358 31.40 d 3.95

.2580 c .2821 .2517 5320 41.70 c 4.86

.2540 a .2738 .1670 6048 35.85 d 4.30 .2271 c .2493 .2534 6769 42.80 f 4.73

.2446 a .2680 .1670 14635 46.00 a 4.16

a Blank of 5 ml NaOCH3 + 20 ml HC1 = 6.10 ml NaOH b Blank of 5 ml NaOCH3 + 20 ml HC1 = 9.87 ml NaOH c Blank of 5 ml NaOCH3 + 30 ml HC1 = 12.38 ml NaOH d Used 20 ml HC1 = 47.7 ml NaOH 69

e Used 20 ml HC1 = 50.4 ml NaOH f Used 30 ml HCl = 72.2 ml NaOH Average k = (2.22 + 0.10) x 10 -5 1 mole-1 min-1

= (0.3703 + 0.027) x 10-6 1 mole-1 sec-1 Average k contains statistical factor of one-half. 70

Table 27 CH2C12 + NaOCH3 in Absolute Methanol at 50 °C.

NaOH = 0.0211 N HC1 = 0.0503 N

Wt. Molarity Molarity Time in M1 NaOH k x 104 CH2C12 CH2C12 NaOCH3 Min. per 20 1 mole -1 min-1 ml HC1

.0974 .21925 .1654 7203 10.83 .346

.1472 .32884 .1641 11439 16.05 .359

.1376 .30798 .1644 18798 19.10 .324

.1327 .29758 .1648 24427 22.20 .347

.1340 .29992 .1644 30297 24.50 .327

Blank of 5 ml of NaOCH3 + 20 ml HC1 = 6.70 ml NaOH

20 ml HC1 = 47.7 ml NaOH Average k = (0.1705 + 0.012) x 10- 4 1 mole -1 min-1 .= (0.0285 + 0.0010) x 10-5 1 mole- 1 sec-1 71

Table 28

CH3Br + NaI in Acetone at 0 °C.

KI03 = 0 . 0100 N

Wt. Molarity Molarity Time in Ml k x 104 QH3Br CH3Br NaI Min. KI03 1 mole -1 min-1

0.0937 .2035 .0468 .38 39.54 18017

0.1096 .237o .0466 .68 37.6o 11899

0.0952 .2063 .0467 1.10 32.70 15020

0.1206 .2608 .0466 1.48 19.93 22567

0.0538 .1173 .0470 6.00 23.70 7754

0.0752 .1633 .0468 10.00 12.70 8961

Blank of 5 ml NaI solution = 45.40 ml KI03 solution 72

Table 29 CH3CH2Br + NaI in Acetone at 20.3 °C.

KI03 = 0.0100 N 2.7361 gms of CH3CH2Br per 100 ml of Solution

Molarity CH3CH2Br = .2488

Molarity NaI = .0381

Time in M1 CH3CH2Br k x 104 min-1 Min. KI03 used 1 mole -1 moles/liter

0 38.05

12 26.95 0.0075 746.3

40 20.00 0.0181 675.3

59 16.50 0.0216 600.4

91 12.40 0.0237 528.7

117 10.35 0.0277 478.9

239 6.30 0.0318 330.3 k = 123.1 1 mole -1 sec -1 This k was obtained in the same manner described at the bottom of Table 11.

See "Figure 6". 73

Table 30

CH3CH2Br + NaI in Acetone at 50 °C.

KI03 = 0.0100 N

Wt. Molarity Molarity Time in M1 k x 104 CH3CH2Br NaI Min. KI03 1 mole -1 min-1

.1258 .2170 .0362 2.13 25.10 9372

.1468 .2528 .0362 2.90 18.50 10349

.1664 .2855 .0360 3.78 12.50 10890

.1133 .1959 .0363 4.62 15.50 10676

.1780 .3098 .0360 5.05 8.00 10742

.1607 .2762 .0361 6.00 7.05 10914

Blank of 5 ml of NaI = 38.20 ml KI03

Average k = (10490 + 417) x 10 -4 1 mole -1 min-1

-5 1 mole -1 sec-1 = (1748 + 69) x 10 74

Table 31

CH2BrF + NaI in Acetone at 20.3 ° C.

KI03 = 0 0100 N

Wt. Molarity Molarity Time in Mi k x 104 CH2BrF NaI Min. KI03 1 mole -1 min-1

.2280 .3990 .0390 8 34.02 377.1

.1460 .2575 .0393 55 24.60 344.5

.2157 .3782 .0391 73 17.80 298.8

.1659 .2920 .0393 79 19.70 313.6

Blank of 5 ml NaI = 39.50 ml K10 3

Average k = (334.0 + 27) x 10 -4 1 mole -1 min-1

-5 1 mole -1 sec (55.58 + 4.55) x 10-1 75

Table 32 CH2BrF + NaI in Acetone at 50 °C.

KI03 = 0.0100 N

Wt. Molarity Molarity Time in M1 k x 104 CH2BrF CH2BrF NaI Min. KI03 1 mole -1 min-1

.1803 .3023 .0374 1.05 30.50 8245

.1998 .3344 .0373 1.67 25.00 8058

.1956 .3274 .0373 2.25 21.90 8226

.2392 .3989 .0371 2.78 16.60 8064

.2166 .3618 .0372 3.5o 15.10 7902

Blank of 5 ml NaI Solution = 39.45 ma. KI03

Average k = (8099 + 109) x 10 -4 1 mole -1 min-1 -5 1 mole-1 sec-1 (1350 + 18) x 10 76

Table 33

CH2C1Br + NaI in Acetone at 20.3 °C.

KI03 = 0.0100 N 3.6989 gms CH2C1Br per 100 ml of Solution Molarity CH2C1Br = .2860 Molarity Nal = .0385

Time in Ml k x 104 Min. KI03 1 mole -1 min-1

0 38.5o

64 35.6o 43.01

176 31.20 42.37

253 28.10 44.01

378 23.8o 45.75

571 19.55' 43.29

1360 7.6o 44.8o

1992 4.90 39.27

Average k = (43.27 + 1.48) x 10 -4 1 mole -1 min-1 = (7.12 + .25) x 10 -5 1 mole -1 sec -1 NUMINNEWN ofteatilariloimo,

77

Table 34 CH2C1Br + NaI in Acetone at 36 °C.

KI03 = 0.0100N

Wt. Molarity Molarity Time in Ml k x 104 CH2C1Br CH2C1Br NaI Min. KI03 1 mole -lmin-1

0.1851 .2763 .02992 30.67 24.52 282 0.2052 .3056 .02987 90.0 14.30 287

0.2064 .3073 .02987 90.0 14.50 283

0.1611 .2408 .02998 120.0 14.21 280

0.1466 .2196 .03004 150.0 12.60 286

0.1675 .2505 .02998 180.0 9.13 284

0.1812 .2703 .02992 210.33 6.71 284

Blank of 5 ml of NaI = 31.00 ml KI03 Average k = (284 + 3) x 10- 4 1 mole-1 min-1 = (47.33 + .50) x 10-5 1 mole-1 sec-1 78

Table 35

CH2C1Br + NaI in Acetone at 50 °C.

KI03 = 0.0100N

Wt. Molarity Molarity Time in Ni k x 104 CH2C1Br CH2C1Br NaI Min. KI03 1 mole -1 min-1

0.0824 .1220 .02567 5.0 25.05 1308

0.1257 .2931 .02703 10.5 19.85 1211 0.1615 .2363 .02542 15.0 17.00 1313 0.1480 .3172 .02698 21.0 12.42 1293 0.1551 .1440 .02734 25.92 17.92 1301 0.1308 .1925 .02552 31.83 12.30 1324

0.0734 .1086 .02567 34.92 16.80 1302

0.1353 .1992 .02552 46.o 8.30 1349

0.1224 .1802 .02552 50.0 8.60 1335

0.1558 .2288 .02547 54.58 5.55 1337

Blank of 5 ml NaI = 26.80 ml KI03 Average k = (1324 + 14) x 10- 4 1 mole -1 min-1 = (220.7 + 2.4) x 10-5 1 mole -1 sec-1 79

Table 36

CH2Br2 + NaI in Acetone at 20.3 ° C.

KI03 = 0.0100 N

4.9107 gms CH2Br2 per 100 ml of Solution

Molarity of CH2Br2 = .2825

Molarity of Ned = .0315

Time in Ml k x 104 Min. KI0 1 mole -1 min-1 3

0 31. 50

146 28.38 25.25 382 24.50 23.58 1092 15.60 24.36

1840 9.80 25.58

2877 4.50 25.58

Average k = (12.435 + 0.360) x 10 -4 1 mole -1 min-1

-5 1 mole -1 sec -1 = (2.073 + 0.060) x 10

Average k contains a statistical factor of one-half. 80

Table 37 CH2Br2 + NaI in Acetone at 36 °C.

KI03 = 0.0100 N

Wt. Molarity Molarity Time in M1 k x 104 CH2Br2 CH2Br2 NaI Min. KI03 1 mole -1 min-1

.2734 a .3031 .02809 31 24.30 196 .2252 a .2505 .02820 90.17 19.20 188

.2705 a .2998 .02809 119.25 15.37 186

.2430 a .2699 .02815 319.83 5.88 197 .2002 b .2232 .02773 362 6.30 199

a Blank of 5 ma. NaI = 29.15 ml KI03 b Blank of 5 ml NaI = 28.62 ml KI03

Average k = (96.0 + 2.5) x 10 -4 1 mole -1 min-1

= (16.0 + .4) x 10-5 1 mole -1 sec -1 Average k contains a statistical factor of one-half. 81

Table 38 CH2Br2 + NaI in Acetone at 50 °C.

KI03 = 0.0100 N

Wt. Molarity Molarity Time in M1 k x 104 CH2Br2 CH2Br2 NaI Min. KI03 1 mole -1 min-1

.2477 b .2699 .0274 5.17 25.90 805 .2411 b .2627 .0274 11.00 22.90 831 .2112 a .2306 .0276 14.92 21.82 851

.2358 b .2568 .0274 21.50 18.68 817

.2636 a .2866 .0275 32.00 13.80 834 .1534 a .1683 .0277 34.58 18.30 826

.1235 a .1358 .0278 39.83 18.80 837

a Blank of 5 ml NaI = 29.05 ml KI03 b Blank of 5 ml NaI = 28.98 ml KI03 Average k = (414.5 + 4.5) x 10 -4 1 mole -1 min-1 = (69.08 + 0.92) x 10 -5 1 mole-1 sec-1 Average k contains a statistical factor of one-half. 82

Table 39 CH2BrI NaI in Acetone at 20.3 °C.

KI03 = 0.0100 N

5.7653 gms CH2BrI per 100 ml of Solution Molarity CH2BrI = 0.2610

Molarity NaI = 0.0395

4 Time Ml CH2BrI k x 10 KI03 used 1 mole -1 min-I Moles/liter

0 39.50

63 38.00 0.0015 24.0

172 34.42 0.0051 46.4

249 32.18 0.0073 27.8

376 30.10 0.0094 23.1

57o 25.6o 0.0139 29.9

1357 15.70 0.0238 30.3

1986 11.10 0.0284 26.9

2879 9.10 0.0304 21.2

4568 6.90 0.0326 16.1

6324 6.50 0.0330 12.0

-4 -1 -1 Average k = 33.95 x 10 1 mole min

= 5.66 x 10 -5 1 mole -1 sec -1 This k was obtained as described on the bottom of Table 11. See "Figure 7". 83

Table 40

CH2BrI + NaI in Acetone at 50 °C.

KI03 = 0.0100N

Wt. Molarity Molarity Time in M1 CH2BrI k x 104 CH2BrI CH2BrI NaI Min. KI03 used 1 mole -1 min-I mil

0.3291 .2811 .0270 10 22.45 .0058

0.2784 .2387 .0282 29 19.85 .0094 572

0.2889 .2473 .0281 59 15.00 .0139 485

0.3049 .2610 .0281 88 11.92 .0169 414

0.2785 .2388 .0282 112 9.50 .0192 449

0.2937 .2545 .0281 147 7.30 .0212 396

0.2540 .2182 .0282 120 9.40 .0193 465

0.2975 .2546 .0281 161 6.20 .0223 405

0.3322 .2838 .0281 177 5.80 .0226 343

Blank of 5 ml NaI = 29.75 ml KI03 Average k = 698 x 10 -4 1 mole -1 min-1 = 116 x 10-5 1 mole -1 sec-1 The value of k was obtained by multiplying the value shown for the

CH2BrI used by the ratio of original CH2BrI of the second line of data and plotting this versus time. The initial slope was then obtained and divided by initial concentrations in the second line of data to obtain the initial rate constant. See "Figure 8". 84

Table 41 CH2C12 + NaI in Acetone at 50 °C.

KI03 = 0.0100 N

Wt. Molarity Molarity Time in Ml k x 104 CH2C12 CH2C12 NaI Min. KI03 1 mole -1 min-1

.0606 a .1363 .02734 4323 24.70 2.50 .0428 a .0967 .02745 5814 24.80 2.55 .0549 a .1238 .02739 7200 23.05 2.47

.0266 a .0602 .02750 8640 25.40 2.51 .0857 a .1922 .02724 10080 18.04 2.48 .1085 b .2423 .02998 17286 11.35 2.56 .0568 b .1282 .03027 18732 17.70 2.56

a Blank of 5 ml NaI = 28.60 ml KI03 b Blank of 5 ml NaI = 31.60 ml KI03

Average k = (1.265 + 0.020) x 10 -4 1 mole -1 min-1

-5 1 mole -1 sec-1 = (0.211 + 0.003) x 10 Average k contains a statistical factor of one-half. 85

Table 42

CH2C12 + Nal in Acetone at 60 °C.

KI0 = 0100N 3

Wt. Molarity Molarity Time in Mi k x 104 CH2C12 CH2C12 Nal Min. KI03 1 mole -1 min-1

0.1966 .4262 .02528 1122.5 17.70 9.2

0.2172 .4715 .02523 1440 15.20 8.9

0.1704 .3722 .02537 1801.5 15.10 9.1

0.2426 .5248 .02514 2568 8.5o 8.8

0.2484 .5366 .02509 2882.5 7.35 8.7

0.1830 .3974 .02532 3238.5 9.00 8.9

Blank of 5 ml Nal = 27.4 ml KI03 Average k = (4.45 + 0.10) x 10 -4 1 mole-1 min-1

. (0.842 + 0.017) x 10 -5 1 mole -1 sec -1 Average k contains a statistical factor of one-half. 86

Table 43

CH2C1I + NaI in Acetone at 50 ° C.

KI0 =- 0 0100 N 3

Wt. Molarity Molarity Time in ML k x 104 CH2C1I CH2C1I Nal Min. KI03 1 mole-1 min-1

.2110 a .2269 .02985 18430 21.15 .974

.2704 b .2892 .02807 14499 20.30 .927

.2037 b .2191 .02823 25892 17.80 .932

'.2523 a .2703 .02974 27415 16.75 .875

.2274 b .2441 .02817 36013 13.00 .978

.2633 b .2821 .02812 54990 7.90 .898

.2781 b .2974 .02807 54995 7.42 .885

a Blank of 5 ml of NaI = 31.45 ml KI03 b Blank of 5 ml of NaI = 29.75 ml K103

Average k = (0.931 + 0.031) x 10 -4 1 mole -1 min-1

-5 1 mole -1 sec -1 = (0.155 + 0.005) x 10 87

Table 44

CH2C12 + NaOCH3 in Absolute Methanol at 50 °C.

NaOH = 0.0218 N HC1 = .0503 N

Molarity Time in Ml NaOH k x 104 Wt. Molarity 1 -1 CH2C12 CH2C12 NaOCH3 Min. per 20 1 mole min ml HC1

.0974 .2193 .1654 7203 10.83 .346

.1472 .3288 .1641 11439 16.05 .359

.1137 .2555 .1651 14331 16.05 .368

.1376 .3080 .1644 18798 19.10 .324

.1327 .2976 .1648 24427 22.20 .347

.1340 .2999 .1644 30297 24.50 .327

Blank 5 ml NaOCH3 + 20 ml HC1 = 6.70 ml NaOH

20 ml HC1 = 45.9 ml NaOH Average k = (0.1705 + 0.0060) x 10 -4 1 mole -1 min-1 -5 1 mole -1 sec -1 = (0.0285 + 0.0010) x 10

Average k contains a statistical factor of one-half. 88

APPENDIX C

GRAPHS CH2Ig

CHBr I

C B H2C I CH31 .,

CH2C10Br

CH2BrF

0 CHr 0 CHC1-(1er -0-- CH CI -

1,2 Figure 10 Densities of Holomethanes

.80

.6 40 50 20 30 10 Temperature in °C 90

F igure 2 Log of Rate Constant Versus Reciprocal of the Absolute Temperature o This work • Dostrovsky and Hughes

Br NaOCH 70 CH2 2 3

H2CIBr 60 NaOCH 3

0 CH2Bre Na I

40 CH2BrC Nal

Br Nal CH3CH

20 —

1.0 350 340 3.30 3.20 3.10 3.00 1 / T x 10 3 70

65 0 CH2CIBr

60 re) 0 / 0

crb 0 —J 5.0 )CH3CH2 Br

Figure 3 45 Relationship Between Rate of Reaction of Alkyl Bromides with Nal and NaOCH 40 3

3.5 2.9 3.0 3.5 4.0 4.5 50 - Log k for Nal 4.5

3.5 0.1 0 x 5.0 -0 25

Initial slope =0.0000012 oP 1.5

I.0 Figure 4 CH BrIs NaOCH in Methanol 2 3 20 °C

10 20 30 40 50 60 70 80 90 100 110 120 130 140 Time in Minutes x 10 Figure 5 3.6 CH21 2 * NaOCH3 in Methanol 20,3 ° C 3.2

cO —x 24

3 2D D (-.11.6 JJ 0 1.2

10 20 30 40 50 60 70 Titne in Minutes x I0-3 0 0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 Time in Minutes Figure 7

1.8 CH2Br1+ Nal in Acetone 20.3°C 1.6

T3 .8

.6

4

.2

o o 100 200 300 400 500 600 700 Time in Minutes 0 0 10 20 30 40 50 60 70 80 90 100 110 120 130 140 150 Time in Minutes 97

Bibliography

1. E. D. Hughes, Transactions of the Faraday Society, 37, 603 (1941).

2. J. Hine and D. E. Lee, Journal of the American Chemical Society, 73, 22 (1951).

3. J. Hine and D. E. Lee, ibid, 74, 3182 (1952).

4. P. Petrenko-Ktitschenko, D. Talmud, B. Talmud, W. Butmy-de-Katzman and A. Gandelman, Zeitschrift fur Physikalische Chemie, 116, 313 (1925).

5. P. Petrenko-Kritschenko, Journal fur Praktische Chemie, 111, 23 (1925).

6. P. Petrenko-Kritschenko and V. Opotsky, Berichte der Deutschen Chemischen Gesellschaft, 59D, 2131 (1926).

7. P. Petrenko-Kritschenko, A. Ravikovich, V. Opotsky, E. Putyatui and M. D'yakova, Journal of the Russian Physical Chemical Society, 6o 149 (1928).

8. P. Petrenko-Kritschenko, A. Ravikovich, V. Opotsky, E. Putyatui and M. D'yakova, Berichte der Deutschen Chemischen Gesellschaft,

61B , 845 (1928).

9. P. Petrenko-Kritschenko, V. Opotsky, M. D'yakova and A. Losovoi, ibid, 62B, 582 (1929).

10. P. Petrenko-Kritschenko, V. Opotsky, E. Putyatui and M. D'yakova, Journal of the Russian Physical-Chemical Society, 61, 1778 (1929).

11. P. Petrenko-Kritschenko, Ukrainskii Khemichnii Zhurnal, 1, 304 (1929).

12. J. B. Conant and W. R. Kirner, Journal of the American Chemical Society, 46, 232 (1924).

13. J. B. Conant and R. E. Hussey, ibid., 47, 476 (1925).

14. J. B. Conant, W. R. Kirner and R. E. Hussey, ibid., 47, 488 (1925).

15. L. W. Andrews, ibid., 25, 756 (1903). 98

1 . K. L. Senior, R. R. Hetrick and J. G. Miller, ibid., 66, 1987 (1944).

17. J. Stastny, Chemische Zentralblatt, (1942), II, 1039.

18. Beilsteins Handbuch der Organischen Chemie, Band I, Verlag von Julius Springer, Berlin, Germany, 1918, pp. 273 and 635. 19. J. Timmerman and Mme. Hennault-Roland, Journal de Chimie Physique, 29, 529 (1932).

20. International Critical Tables, Vol. III, 1st ed., McGraw-Hill Book Co., Inc., New York, N. Y. (1926), p. 27.

21. L. F. Fieser, Experiments in Organic Chemistry, Part II, 2nd ed. D. C. Health and Co., New York, N. Y. (1941), p. 359.

22. F. Swarts, Bulletin de 1'Academie Royale de Beige, 113 (1910). 23. O. Ruff, O. Bretschneider, W. Luchsinger and G. Miltschitsky, Berichte der Deutsche Chemische Gessellschaft, 69, 299, (1936).

24. H. Gilman, Organic Chemistry, Vol. I, 2nd ed., John Wiley and Sons, Inc., New York, N. Y. (1948), p. 362.

25. C. A. Kraus and W. C. Bray, Journal of the American Chemical Society, 35, 1315 (1913).

26. I. Dostrovsky and E. D. Hughes, Journal of the Chemical Society, 161, (1946).

27. K. J. Laidler, Chemical Kinetics, 1st ed., McGraw-Hill Book Co., Inc., New York, N. Y. (1950), p. 9.

28. H. M. Goodwin, Elements of the Precision of Measurements and Graphical Methods, 2nd ed., McGraw-Hill Book Co., Inc., New York, N. Y., (1920), p. 26 ff. 99

PART II A STUDY OF THE RATE OF PROTON TRANSFER REACTIONS 100

CHAPTER I

INTRODUCTION

It has long been assumed that the rate of proton transfer from an oxygen, nitrogen, sulfur or halogen acid to a base containing the same elements, proceeds at an immeasurably fast rate and requires zero energy of activation. However, it is useful to postulate such proton transfers as slow, rate controlling steps in some reactions which are

considered to be general acid and/or base catalyzed (1, 2). It seems desirable to make a study of this type reaction in order to determine if any such proton transfer does indeed occur at a rate which is slow enough to measure. The literature contains a few references to such cases. A controversial paper by Orr (3) states that, by following the exchange of deuterium between heavy water and ordinary ethanol, he was able to follow the rate of proton transfer. He found for this reaction: k2 C2HSOH + HOD -457== c2HSO + H2O

k = 2.2 + 0.2 x 10-6 1 mole-1 sec -1 2 6 1 mole-1 sec-1 k = 2.4 + 0.2 x 10-

1. R. P. Bell, Acid-Base Catalysis, Oxford University Press, London, (1941), Chap. VII.

2. L. P. Hammett, Physical Organic Chemistry, McGraw-Hill Book Co., Inc., New York, N. Y., (1940), p. 241.

3. W. J. C. Orr, Trans. Faraday Soc., 32, 1033 (1936). 101

To follow the reaction he removed the water from the reaction mixture as a hydrate of calcium sulfate, removed the water from this hydrate under reduced pressure and determined its density with a pycnometer and also its refractive index with an interferometer and was thus able to obtain the amount of deuterium exchange even though a small amount of ethanol contaminated the sample. Junger and Wirtz (4) disagree with this work since they have found that the exchange between alcohol vapors and solid calcium deutoxide is extremely rapid, being completed in less than an hour. Jungers and Bonhoeffer (5) have also found the exchange between various alcohols and heavy water to be too fast to measure using an extraction technique whereby the alcohol is removed with . On the other hand Geib (6) has reported that the rate of exchange of deuterium in the reaction:

H25 + CH3OD CH3OH + D2S is measurable at low temperatures; i.e., -117 degrees. Brodskii and Sulima (7) found that the exchange between NH4+ and D20 is very slow but that NH3 and D20 exchange almost instantaneously. They attributed this to the fact that nitrogen in the ion is like carbon and silicon in that it lacks free electron pairs for coordination,

4. J. C. Jungers and K. Wirtz, Bull. Soc. Chim. Belg., 45, 679 (1936).

5. J. C. Jungers and K. F. Bonhoeffer, Z. Physik. Chem. A177, 460 (1936).

6. K. F. Geib, Z. Elektrochem., 45, 648 (1938).

7. A. J. Brodskii and L. V. Sulima, Doklady Acad. Nauk. S.S.S.R., 74, 513 (1950). 102

whereas the nitrogen in has an available free electron pair. This work has been successfully checked by Professor C. G. Swain at the Massachusetts Institute of Technology (private communication).

Briscoe and co-workers (8, 9, 10) and Garrick (11) found that the cobaltammines exchange the amine with heavy water slowly. Ogston (12) found that the conductivity of certain amines in methanol and ethanol solution changed gradually over a period of ten minutes and forty minutes, respectively. This he attributed to a reaction as follows: B C2H5OH + C2H5OH C2HSO- + C2HSOH2+ C2HSOH + BH + where B indicates a base; i.e., the amine. This explanation fits the data fine except that approximately two-thirds of the reaction is completed in the first two minutes, after which it follows a first order course and this phenomena is not explained by this mechanism. Schaefgen, Newman and Verhoek (13) found that amines in methanol solution

8. F. W. James, F. S. Anderson and H. V. A. Briscoe, Nature , 139, 109 (1937). 9. H. V. A. Briscoe et al, J. Chem. Soc., 1492 (1937).

10. J. S. Anderson, N. L. Spoor and H. V. A. Briscoe, Nature, 139, 508 (1937). 11. F. J. Garrick, Nature, 139, 507 (1937).

12. A. G. Ogston, J. Chem. Soc., 1023 (1936). 13. J. R. Schaefgen, M. S. Newman and F. H. Verhoek, J. Am. Chem. Soc., 66, 1847 (1944). 103

caused bromcresol purple to fade slowly, which can be explained either

on the basis of a slow proton transfer or a slow indicator reaction. Professor R. G. Pearson of Northwestern University has informed us privately that a reinvestigation of these experiments shows that the presence of small amounts of carbon dioxide was apparently responsible

for the results. Nyman, Fung and Dodgen (14) found that the exchange reaction

between ammonium ions and ammonia , using radioactive nitrogen, is not slow enough to measure using their technique. This technique required considerable time to distil off the ammonia and they suggest that a technique using radioactive hydrogen might enable one to measure the rate of the reaction. They point out that if the reaction is

either of the following possibilities:

NH4 + NH3 NH3 + NH4 (1) 2NH3 NH4 + NH2- (2) * NH2 - + NH4 NH3 + NH3 then it is apparently a proton transfer which is slow.

Nelson and Butler (15) report that in the alkaline decomposition of diacetone alcohol, deuterium in either the solvent or the alcohol slows the rate of decomposition. This, they claim, is an indication

that the rate determining step is a proton transfer from the alcohol

14. C. J. Nyman, S. C. Fung and H. W. Dodgen, J. Am. Chem. Soc., 72, 1033 (1950).

15. W. E. Nelson and J. A. V. Butler, J. Chem. Soc., (1938), 947. to an OH - ion or an OD- ion. LaMer and co-workers (16, 17) while studying the solvolytic decomposition of nitramide in light and heavy water mixtures proposed that the rate controlling step is a proton transfer from the nitrogen atom in nitramide to the solvent. Ogawa

(18) has found that at sixteen degrees Centigrade the exchange between urea and heavy water is only sixty per cent of the calculated

theoretical equilibrium value after five minutes. This calculated value was not shown experimentally to be the actual value. Moreover, these results were obtained by precipitating the urea with mercuric nitrate, and it is known that in many instances, precipitation causes highly erroneous results in isotope work since the exchange is quite

often driven to completion by the precipitation. It is possible that a reaction of this type, in which deuterium is used as a tracer, may go through either of two mechanisms. The first is an ionic mechanism

ROH + 020 --a. RO- + D2011+

RO- + D20H+ ROD + DOH, etc., which seems improbable if the work of Brodskii and Sulima is confirmed

16. V. K. LaMer and J. Greenspan, Trans. Faraday Soc., 33, 1266 (1937). 17. V. K. LaMer and S. Hochberg, J. Am. Chem. Soc., 61, 2552 (1939). 18. Eijiro Ogawa, Bull Chem. Soc. Japan, 11, 367 (1936). 105

by further experiments. The second mechanism involves cyclic inter- mediates of the types: / D 7H ci—ok R - 0 N4DzO or etc. \D R - 0'''1D

For either of these mechanisms it seems reasonable that those compounds, for which the exchange has the best possible chance of being slow, are those which are weak acids and weak bases. It also would seem to be highly advantageous if the two compounds involved were of nearly the same acidity and basicity. io6

CHAPTER II

INSTRUMENTATION AND EQUIPMENT

Of the various methods which are available for the determination of heavy water, it seems that the one which is most adaptable to kinetics work is the falling drop method. This method depends upon the rate of fall of droplets, of exactly the same size but different densities, through an immiscible liquid whose density is very close to that of the droplets. For this, ortho fluoro toluene, density 1.004 at 13.2 °C. has been chosen as the immiscible liquid.

In order to obtain reasonable accuracy by this method, highly specialized equipment had to be built. A very sensitive constant temperature bath and a micropipette were built. The constant temperature bath was made from a two foot piece of glass tubing of approximately four inches outside diameter. The bottom was sealed with a large rubber stopper, cemented on the inside with rubber cement and on the outside with Jewelers wax. A long glass stirrer having blades every two inches was used for agitation. Two heat sources and a cooling coil were placed in the bath. The heat sources were long narrow U-tubes of 3 mm glass tubing containing a single strand of nichrome wire. The heat was regulated by the use of variacs. One of these elements was used as a constant source of heat and the other was used as an intermittent source of heat which was 107

controlled through a relay by the use of a mercury-toluene thermo- regulator. The thermoregulator, "Figure 1", consisted of a two centimeter by fifteen centimeter bulb full of toluene (A), which was connected, through a short length of five millimeter glass tubing (B), to three, two foot lengths of glass tubing. One of these three arms (C) was five millimeter tubing connected through a tungsten to glass seal to be used as one of the contacts. The second arm (D) was one millimeter capillary tubing. The third arm (E) was five millimeter glass tubing having a one millimeter straight bore stopcock (F) sealed onto the top. All three sidearms were filled with mercury and a length of nichrome wire was immersed in the mercury of the first arm and connected to one side of the relay. The second arm had a threaded bakelite sleeve cemented to the top through which a steel pointer connected to a screw (H) was passed as the other contact. This screw was connected to the other side of the relay. The stopcock on the third arm was used to make rough adjustments in the thermoregulator setting by adding or removing mercury from the system. The screw arrangement was used to make the fine adjustments in the setting.

The cooling coil was a long U - shaped tube of aluminum tubing through which cooling water could be circulated. The bath itself was insulated with two layers of heavy asbestos paper, in which windows had been cut for observation purposes. It has been found possible to regulate the intermittent heat in this system so that the temperature of the bath varies only 0.002 °C. between turning on and shutting off of the heat. The temperature i o8

G

C D

A

B Figure I Thermoregulator 109

difference was measured with a Philadelphia Precision Differential

Thermometer. The design chosen for the micropipette is that of Roseburg and Van Heyningen (19). The pipette, "Figure 2", consists of a threaded drive shaft (A) which passes through a heavy top which is bolted to the casing. The drive rests upon a piston (C) which is supported by a sponge rubber cusion (D). The piston fits snugly into the block (E) which in turn fits snugly into a short section of glass tubing on the end of one millimeter capillary tubing (I). The glass tubing and block are held together with a short length of pressure tubing (F). The capillary tubing passes through rubber gaskets (G) and a knurled, threaded cap (H) and is bent in the shape shown. It is connected through stopcock (J) to a mercury reservoir

(K). A hole in the top (B), for filling the chamber with mercury, is fitted with a loose fitting plug (L). A peg (M) is threaded to the driving shaft (A). Two moveable stops (N) are clamped to the top and limit the arc through which the shaft may be turned.

The pipette is mounted on an assembly, "Figure 3", which makes it possible to raise and lower the pipette evenly. Two bars (A) of solid brass are drilled to accommodate two rods (B) and a heavy screw

(C). The two bars (A) are fastened to the rods (B) and the ring stand

(C) with set screws. A heavy, solid brass bar (E) is also drilled to accommodate the rods (B) and tapped out to fit the screw (D). The

19. F. Roseburg and W. E. Van Heyningen, Ind. Eng. Chem. Anal. Ed., 14, 363 (1942). 1 10

M N1

I

Figure 2 Micropipette 111

yoke (F) holds the pipette and is fastened to the traveling bar (E).

Screw (D) is threaded along the entire length which is between the two bars (A) and by turning the crank (G) the traveling bar (E) is caused to move up or down. This arrangement, when thoroughly greased, gives quite easy control over the pipette.

To operate, the shaft is backed off to the stop and sufficient mercury is admitted from the well through the stopcock to just fill the capillary to the delivery tip. The tip is lowered into the sample and some of the sample is drawn up into the capillary by allowing a small portion of the mercury to run back into the well. The tip is wiped dry, being careful not to remove any liquid from inside the capillary, and is immersed in the immiscible liquid. The knob is turned until it hits the other stop. The droplet which is formed is too small to detach itself from the tip of the delivery tube but if the tip is slowly raised out of the liquid, the droplet is detached as the tip breaks the surface.

The pipette was calibrated by making up mixtures of ordinary water and heavy water by weighing accurately on a Christian-Becker

Chainomatic Balance equipped with a magnetic damper. Droplets of these mixtures were allowed to fall a given distance through the immiscible liquid and the time of fall determined with an electric timer. The data in Table 1 was obtained and used to plot the curve in "Figure 4".

For determination of the deuterium content of ethylamine and tertiary butyl alcohol, the Beckman IR2 Spectrophotometer was used.

The ethyl amine was analyzed in a gas cell 10 cm. long with rock G

0 A

E F I j I

0 A

Figure 3 Elevating Assembly 113

.14 ow J3 o

°al .12 a) a) II 0 =a .10 LL. ,0- 09

g i= 08 II

05

04

03 Figure 4 .02 Calibration Curve for Micropipette .0I 30°C

0 0 10 20 30 40 50 60 70 80 90 114

Table 1. Calibration data for Micropipette

Wt. of Wt. Mole % Average 1 - 1 99. HOH D time of ti t2 DOD fall

.4427 .0974 80.4 8.3 .0184

.3365 .1977 60.5 10.0 .0389

.2749 .2465 50.1 11.3 .0504

.2213 .2972 40.1 13.4 .0634

.1897 .2956 36.6 13.3 .0656

.1682 .3396 30.8 15.55 .0744

.1368 .3704 25.o 17.4 .0811

.1143 .3973 20.6 18.8 .0875

.0847 .4175 15.5 22.9 .0950

.0587 .4482 10.7 30.65 .1063

.0408 .4503 7.5 32.4 .1096

.0350 .4679 6.3 40.0 .1158

.0219 .4948 3.9 47.o .1295 115

windows and the tertiary butyl alcohol in a liquid cell 0.1 mm. thick. with rock salt windows. The spectra of these compounds using a 0.1 mm. cell "Figures 5 and 6", show that N-H has an absorption maximum at 2.9 microns and 0-D has one at 3.86 microns.

Figure 5 Infrared Absorpt ion Spectrum Partially Deuterated t- Butyl Alcohol - - -- t - Butyl Alcohol Beckmann 1R-2 using 0.1 mm. Cell c 70 - 7 „op •••......

-(7) •Mbr ww ■ .... 60 N 60- u) c 50- % 1-- 1 % t I (1) 40- I U I%■ % i % i a) 30- i a_ I , I % i 20 11 i lk I i i 1 i I i / I 10 i / I t i % , 0 I 'r- - ' 1 I I I I mill 2.6 28 30 32 34 36 3.8 4.0 4.2 4.4 4.6 4.8 5.0 5.2 Wave Length in Microns Figure 6 90- Infrared Absorption Spectrum Ethylamine 80 - Partially Deuterated Ethylamine---- Beckmann IR-2 using 10 cm gas cell at 642 mm c 70 - / / ._c. co 60 E cto 2 50- i--

1)40 - o I I /1 1330 - 1 1 il a. / 1 I / I I 20 - \I /

/ 10 I / .1 II i 1 1 t 0 1 1 1 ii ii 2.6 2.8 3.0 32 34 3.6 3.8 4.0 4,2 4.4 4.6 4.8 5.0 Wave Length in Microns 118

CHAPTER III

PROCEDURE

Experimental Since no attempt was made, in any of the work which contradicts

Orr's (3) claim, to duplicate the conditions used by him and due to the great theoretical importance of his claim, it was decided to duplicate as nearly as practical these conditions. The method employed by Orr was to make up solutions of ethanol and water containing about twenty per cent deuterium in glass stoppered bottles. These were placed in a thermostat at 25 ° + 0.01°C. for various lengths of time. At the desired time, three to four grams of calcium sulfate were added and shaken for thirty minutes. The hydrated calcium sulfate was collected on a filter and dry air was drawn through the filter for forty-five minutes to remove adherent ethanol. The calcium sulfate was then dehydrated for one and a half hours at 150 °C. The water obtained was analyzed by obtaining the density with a pycnometer and the index of refraction with an interferometer. From the data obtained he arrived at a value for the amount of residual ethanol

(<0.4 per cent) and the deuterium content. In order to obtain data in which the deuterium content at random distribution is sufficiently different from that of the starting material to give interpretable results, it was decided to use a higher concen- tration of deuterium and use equimolar quantities of ethanol and water. 119

Hence, a solution of approximately twenty per cent deuterium oxide was prepared and its composition determined at each use. One milliliter of this solution at room temperature was pipetted into a twenty-five milliliter, glass stoppered Erlenmeyer flask and cooled to 0°C. Three and two-tenths milliliters of absolute alcohol at 0 °C. is then pipetted

into this flask and mixed, and allowed to stand at 0 °C. for various lengths of time. The reaction mixture was then poured upon ten grams

of anhydrous calcium sulfate, and shaken for ten minutes. Ten grams

of calcium sulfate were used in an attempt to halt the hydration at the half hydrate stage because the vapor pressure of this hydrate is

especially low. The mixture was then placed in a drying pistol, held

at 35 °C. and 4 mm. pressure for two hours to remove the adhering

ethanol. The temperature was raised to 190 °C., the pressure maintained

at 4 mm. for two hours and the water collected in a dry ice trap.

The water sample was analyzed for ethanol by a modification of the method of Williams and Reese (20). Two-tenths milliliter of

sample was diluted to 2 ml. in an ordinary test tube. One milliliter of a solution of potassium dichromate in concentrated sulfuric acid was added and shaken vigorously. The concentration of the potassium dichromate need not be known since the same solution must always be used and the colorimeter calibrated against known concentrations of ethanol. The test tube was placed in a boiling water bath for exactly two minutes, in an ice bath for one minute and then diluted. The

20. M. B. Williams and H. D. Reese, Anal. Chem., 22 1556 (1950). 120

contents of the test tube were transferred to a 250 ml. volumetric flask, the test tube rinsed into the flask three times, 2.0 ml. of a saturated solution of symdiphenyl carbazide in ethanol added and water added to the mark. 2 ml. of this solution was transferred to a

50 ml. volumetric flask and diluted to the mark. The intensity of the color of the solution was compared to the color of a known standard, which has been freshly prepared in the same manner and has the same ethanol concentration as that used for the calibration curve (Figure 7).

The calibration curve is extremely smooth even with an instru- ment as inaccurate as the DuBosq colorimeter, thus attesting to the accuracy of the measurements. It is seen from the curve that the method gives an accuracy of about + 0.005 per cent in the determination of ethanol. The use of the values obtained from this curve in the calculations leads to an error in density of + 0.00004 density units.

This in turn gives an error in deuterium content of + 0.037 per cent which is less than the error claimed by Roseburg and Van Heyningen for the falling drop method.

The apparent heavy water content was determined by the falling drop method using the equipment previously described. From this data it is possible to calculate the actual deuterium content in the water sample.

Exchange reactions were run between deuterated ethyl amine and heptylamine and tertiary-amyl alcohol, as well as between deuterated tertiary-butyl alcohol and di-tertiary butyl-iso-propyl carbinol. One milliliter of the deuterated compound was sealed in thin walled, glass Figure 7 Calibration Curve for Determination of Ethanol Left Gen set at 25 Blank Contains .608gms Ethanol per 100 ml Solution = 0 a) 0 t.-ow _ ix cil c 76 0 w cr 25 — a co o 03 = 0 20—

I I I I I 15 I I I I I I I I I .300 .400 .500 .600 .700 .800 Grams Ethanol per 100m1 Solution 122

ampoules and weighed. These ampoules were placed in glass stoppered flasks containing the second compound in amounts having exchangeable hydrogen equivalent to that of the compound in the ampoule. The ampoules were then broken and after an appropriate length of time, the volatile component was pumped off and the transmission determined at a frequency where the non-deuterated ethyl amine and the deuterated tertiary-butanol absorbed predominantly.

Only two different time intervals were used in these reactions.

In one case the reactants were mixed and separated in as short a time as possible, approximately two minutes, and in the other the reaction was allowed to run for considerable lengths of time, about eight hours, to allow equilibrium to set in.

Preparation and Purification of Reagents

Heptylamine - Heptaldoxime was prepared by the Organic Synthesis method (21). The crude product was distilled under 6 mm. pressure using a glascol heater maintained at 140-145 °C. The heptaldoxime was reduced with sodium in ethanol after the Organic Synthesis method (22).

The crude heptylamine was fractionated through a forty-five centimeter

Vigreux column and the fraction boiling at 152 °C. was collected.

21. Org. Syn., Collective Vol. II, John Wiley and Sons, New York, N. Y. (1941), p. 313. 22. Ibid., p. 318. 123

Ethyl Amine - Matheson's anhydrous ethyl amine was distilled through a glass helices packed column having an inside diameter of 8 mm., a length of two feet and equipped with a water jacket so that the temperature of the column could be maintained near the boiling point of the desired compound. The ethyl amine, boiling point 16.5-17.0 °C., was collected and stored in the freezing compartment of the refrigerator in a glass stoppered Erlenmeyer flask with the stopper wired down.

Deuterated Ethylamine - Matheson's anhydrous ethylamine was treated with deuterium oxide, distilled and stored as described above.

Deuterium Oxide - Samples of heavy water, obtained from Stuart Oxygen

Company of San Francisco and containing 99.8 per cent deuterium oxide, were used without further purification.

Tertiary Amyl Alcohol - Sharples' tertiary-amyl alcohol was refluxed over sodium and fractionated through a vacuum jacketed column of approximately 2 cm. inside diameter, one meter long, packed with three-eighths inch glass helices.

Deuterated Tertiary Butyl Alcohol - Commercial tertiary-butyl-alcohol was treated with deuterium oxide. The azeotrope was distilled and dried over calcium sulfate. This was next dried over sodium and distilled, boiling point 82.5 °C.

Di-tertiary-butyl-iso-propyl-carbinol - In a three liter, three necked flask equipped with a five hundred milliliter dropping funnel, a

Herschberg stirrer with a mercury seal, and a reflux condenser, was placed 6.8 grams (2.5 moles) of magnesium turnings. Two hundred milliliters of anhydrous ether and approximately 5 ml. of iso-propyl chloride were added along with a crystal of iodine. Stirring was begun 124

and addition of a solution of 196 grams (2.5 moles) of iso-propyl chloride in 800 ml. of anhydrous ether was started. There was no immediate reaction so 1 ml. of pure dry methyl iodide was added to initiate the reaction. An extremely vigorous reaction occurred almost immediately. The addition of the iso-propyl chloride solution was continued at such a rate as to maintain vigorous reflux over a period of approximately six hours.

Gaseous carbon dioxide obtained from the sublimation of dry ice was passed into the reaction mixture until no more gas was absorbed.

It was found necessary to keep the mixture at room temperature since solidification occurred if the temperature was lowered very much.

A solution of 200 ml. of concentrated sulfuric acid in 500 ml. of absolute methanol was added to this mixture at a rate sufficient to maintain vigorous reflux. The mixture separated into two layers and after fractionation of both layers it was found that only the methanol layer gave methyl iso-butyrate, obtained as an azeotrope with methanol. The azeotrope was broken by adding a large amount of salt water and separating the layers. The lighter layer was fractionated through the Todd Still and found to be almost entirely methyl iso- butyrate, boiling point 92.5 °C., which was obtained in 29 per cent yield (0.72 moles).

In a one liter, three necked flask equipped with a Herschberg stirrer with a mercury seal, a 500 ml. dropping funnel and a reflux condenser was placed 200 grams of tertiary-butyl chloride, 300 ml. of anhydrous ether and 80 grams of coarse sodium sand (approximately two milliliter average diameter). The methyl iso-butyrate was added, 125

two milliliter average diameter). The methyl iso-butyrate was added, with stirring, over a period of approximately forty minutes. After the first eight to ten milliliters had been added a vigorous reaction began. Stirring was continued for six hours after the final addition of methyl iso-butyrate. At the end of this time, considerable amounts of unreacted sodium were found in the mixture, so 50 ml. of ethanol was added to destroy it. The mixture was hydrolyzed with a solution of 160 grams (3 moles) of ammonium chloride in 480 milliliters of water. The ether layer was washed with a solution, dried over calcium sulfate and fractionated. A 10 ml. fraction having a boiling point of 120 °C. at 30 mm. pressure was obtained and found to 2 5 have n ip = 1.4592. Bartlett and Schneider (23) report nE5 = 1.4623 for the fraction from 120-122 °C. at 30 mm. pressure.

23. P. D. Bartlett and A. Schneider, J. Am. Chem. Soc., 67, 141 (1945). 126

CHAPTER IV

DISCUSSION OF RESULTS

It was of considerable interest in this investigation to

recalculate the data given by Orr (3) from the data in Table 2 which he had obtained.

Table 2. Data Obtained by Orr

Run Init. Weight Weight Vol. Time Equil Alpha Mole EtOH HOD Soln. in mole % HOD gms. gms. cc Hrs. % HOD

1 21.74 3.3667 5.5172 9.44 2 19.4 .44

2 20.9 2.0664 3.9926 6.38 3.8 18.0 .72

3 20.9 3.772 8.825o 13.16 6.8 18.0 .93

4 20.9 3.7570 6.445o 10.81 26 16.75 1.10

5 20.9 3.7170 6.0101 10.34 33 16.4 1.12 6 20.9 2.2521 3.7694 6.39 108 16.6 1.11

x(c + x) Alpha = (a-x)(b-x)

This equation is derived from the equation: ki C2H5OH + HOD C2H5OD + HOH k2

(a-x) (b-x) 127

Using these equations and the definitions:

K _ (a + b)ae + c c( e - 1 cepab Col e - 1

B =V K2 - 4>• the following rate equation was derived:

2.303 2 - x (K-B) 1 k2t = log B( e - 1 ) 2 - x (K+B)

Using these data and equations, one obtains the values shown in Table 3.

Table 3. Calculated Rate Constants

Run B 2.303 log 2)-x(K-B) k2 tB O0(e-1) 2-x(K+B) 1 mole-1 se c -1

1 377.0 545.1 374.1 7.8 x 10 -6 .30856 2.41 x 10 -6

2 391.5 510.0 388.8 3.94 x 10 -6 .58737 2.31 x 10 -6

3 405.5 484.0 403.1 2.1 x 10 -6 .93991 1.97 x 10 -6

4 380.5 519.7 377.8 • 59 x 10 - 6 2.12856 1.26 x 10 -6

5 375.3 525.6 372.5 .47 x 10-6

6 378.8 523.3 376.0 .15 x 10 -6 1.82706 .274 x 10 -6

From this Orr claims the rate constant to be 2.2 + 0.2 x 10 -6 1 mole -1 sec -1 . It can be seen that there is no constant value for the

1 See p.134 for derivation. 128

rate constant and it is, in fact, steadily decreasing.

Further, if one calculates the amount, of deuterium which would be expected in the resulting water if the distribution of the deuterium were random among all the species the percentages in Table 4 are obtained.

Table 4. Calculated Equilibrium Percentages

Run millimoles d13 D Mole Mole EtOH HOD HOH Total in % HOD '/40 HOD H & D soln. Calc. Orr

1 73.19 65.83 237.03 678.91 9.70 19.40 19.4

2 44.92 45.87 173.39 483.44 9.49 18.98 18.0

3 82.03 101.39 383.25 1051.31 9.64 19.28 18.0

4 81.67 74.05 279.89 789.55 9.38 18.76 16.75

5 80.80 69.05 261.01 740.92 9.32 18.64 16.4

6 48.96 43.31 163.70 462.98 9.35 18.70 16.6

It now seems fairly obvious that all of the samples had reached approximately the same point and therefore must be at equilibrium. The starting and equilibrium amounts of deuterium are so nearly alike that the accuracy of the experiments is not sufficient, for the results to be considered valid especially when one realizes that in no case would an error in weighing of more than 0.4 milligrams be necessary to account for the deviations from equilibrium. Since the densities are taken with 129

an ordinary pycnometer, the errors of filling the pycnometer and the errors of weighing are more than likely greater than this allowable error. It has been possible to substantiate these beliefs using the procedure described on p. 119. The data obtained and calculated?- are tabulated in Table 5.

It can be seen that in these three experiments, all run for different lengths of time, that the reactions have all proceeded to approximately the same point relative to the calculated equilibrium value. Thus, the reaction between ethanol and heavy water has been found to be too fast to measure under these conditions.

It was felt that compounds of low acidity and basicity and having approximately the same acidity and basicity constants would have the best chance of undergoing a slow proton transfer. It is to be expected that most aliphatic amines have comparable acidity constants and that these constants are quite low so that it was decided to use a low boiling amine and one which was quite high boiling. The two chosen were ethylamine and heptylamine. At the same time it was decided to run ethylamine with a tertiary alcohol since these alcohols are also of low acidity. The alcohol chosen was tertiary-amyl alcohol.

Since there might be some steric effects which would cause a decrease in the rate of proton transfer it was decided to study the exchange between deuterated tertiary-butyl alcohol and di-tertiary-butyl- iso-propyl carbinol.

1 See sample calculations p. 139. 130

Table 5

Results of Exchange Between Heavy Water and Ethanol

1 2 3

Volume of ethanol 3.2 3.2 3.2

Weight of ethanol 2.5800 2.488 2.5800

Millimoles of ethanol 56.00 54.01 56.00

Volume of water 1.0 1.0 1.0

Reaction time 0.5 hrs. 40.75 hrs. 63 secs.

Millimoles water 55.26 55.26 55.46

Drop time of water (start) 21.35 21.4 21.4

Atom % D in water (start) 17.0 17.0 17.0

Milliequivalents D 18.79 18.79 18.86

Total milliequivalents H and D 166.38 161E.39 166.92

Ethanol in product gm5/100 ml 0.75 0.76 0.87

Drop time of product 35.0 32.8 37.3

Corresponding atom % D 8.1 8.8 7.5 Density of product 1.00443 1.00518 1.00378

Atom % D in product 9.38 10.12 9.00 Atom % D from random distribution 11.28 11.41 11.30

Temperatures 0° 30° 0° 131

The data obtained from these reactions are tabulated in Table 6.

Table 6. Results of Deuterium Exchange Reactions

Reaction Wt. Wt. Wave % Transmission A + B A B length Short Extended microns Reaction Reaction

Ethylamine .8316 2.1253 2.90 51+

n-heptylamine .7675 1.9628 2.90 50

Ethylamine .7738 3.0397 2.90 95

t-amylalcohol .7883 3.0829 2.90 95 t-butylalcohol .7961 2.0032 3.86 27. 0

di-t-butyl- .7853 1.9786 :3.86 28.5 i-propyl carbinol

It can be seen that in every case the transmission after the very short reaction time is essentially the same as after the extended reaction time. Thus it can be concluded that the exchange was complete in the very short reaction time. 132

CHAPTER V

CONCLUSIONS

From this study it can be concluded that the prototopic exchange between ethanol and water occurs at a rate too fast to measure by the

method used.

The exchange of deuterium between ethylamine and normal heptyl-

amine occurs at a rate too fast to measure by the method used.

The exchange of deuterium between ethylamine and tertiary amyl

alcohol occurs at a rate too fast to measure.

The exchange of deuterium between tertiary butyl alcohol and

di-tertiary-butyl-iso-propyl carbinol occurs at a rate too fast to

measure. 133

APPENDIX B

SAMPLE CALCULATIONS 134

I Derivation of rate equation used by Orr.

kl ROH + HOD ROD + HOH k2 (a-x) (b-x) x (c + x)

Rate of formation of ROD dx = [ROH] [HOD] - k2 [ROD] [1:10113 dt dx = k (a-x) (b-x) - k2 (c + x) x dt 1

At equilibrium dx = 0 and therefore dt kl _ (c + x) x k2 (a-x) (b-x)

0( e = exchange coefficient at equilibrium = therefore kl = k2

Substitute this for la in rate equation dx = c< k (a-x)(b-x) - k2 (c + x) x dt e 2 dx k2 [cc (a-x)(b-x) - (c + x) x] dt k2dt = dx ckle (a-x)(b-x) - x (c + x)

dx k2dt = ale ale x2 - (a + b) x + c

135

This is an equation of the type dx k2dt = 2 Nbc + Nx + Q

Let M = CKe -1

N = -(°(ea + (Kelp + c)

Q = c< eab R = integration constant k2t 2.303 log 2Mx + N + R N2 = 1NQ 2Mk + N + V/ N2 - at t = 0, x = 0

R =- 2.303 log ,N2 - N2 2414Q. N + 1/N2 - 4MQ Substituting k2t = 2.303 log 2MNx + 2Mx V N2 - i[MG. + li-MQ. i/N2 _ 4.m0. 2MNx - 2Mx 1/N2 - 4-MG. + 4-MQ. k2t -- 2.303 log 2Mx (N + i/N2 - 4MQ.,) + 4ma N2 - 14mQ. 2Mx (N - 111277:14a) + 4MQ.

Let K = [(a + b) c.4 e =- N 1 M

1 icieeab] = G. °‹e 1 M

B= /K2 - 136

k2t = 2 .303 log xkiM2K2 - 414. - ME] + 2MX m2K2 41,4X q - -x [4 M2K2 - 4M + MK, + 2M A 2.303 x [B-1C1 + 2X log k2t _ MB -x [13+K J + 2A

2.303 2>+ - x (.1c-A k2t B(cice -1) log 2 ). - x [IC+B]

II Calculation of mole fraction of heavy water in falling drop experiment.

If we assume that the partial molar volume of ethanol in heavy water is the same as in ordinary water, then

N1 = mole fraction H2O

N2 = mole fraction HOD N3 = mole fraction Ethanol

Ml = molecular weight H2O

M2 = molecular weight HOD

M3 = molecular weight Ethanol V1 = partial molar volume H2O

V2 = partial molar volume HOD

V3 = partial molar volume Ethanol d = N1M1 + N2M2 + N3M3 Ni Vi + N2V2 + N3V3

137

niMi + n2M2 + n3.43 wi + w2 + w3 d = w2 niVi + n2V2 + n3V3 wl + Vi V2 + 3 M 4- i m2 m3 V3

Thus since the falling drop method gives the concentration of the sample if it had not contained any ethanol. We can therefore calculate the density of the actual sample.

d = BA N2M2 N1V1 + N2V2

For run number I

d = .17 x 20.028 + .83 x 18.016 .17 x 18.153 + .83 x 18.094

d = 1.00443

Using this value: w w2 1.00443 = 2 '73 ri 4 m2142v. 2 4. 143128 v3 1

Basing the calculations on one ml of solution,

N./ + w + w = 1.00443 1 2 3 w3 = 0.00760

w = 1.00443 - 0.00760 - 1 w2 = .99686 - Thus

18.153 18.094 _ (•99683 - wi) + 4. (0.00760) = 1 20.028 wl 18.016 456.0579

138

w2 = .10317

1172 .10317 M2 N2 20.028 w w .8879 + .1042 + .0079 + M1 M2 18.016 20.028 46.07

N2 = .0938 139

Calculation of equilibrium per cent HOD if random distribution is assumed

If 1.0 mole of water containing 20 mole per cent HOD is added to 1.0 mole of ethanol, then the water will have 2.0 moles of exchangeable hydrogen (of both isotopic species) and the ethanol will have 1.0 mole of exchangeable hydrogen. There are 0.2 moles of exchangeable deuterium from the water. Since at equilibrium the deuterium will be distributed uniformly among the exchangeable hydrogens, the mole per cent deuterium in any species will be equal to the mole per cent deuterium in the solution. Therefore

Mole % D .2 x 100 - 13.3 1 + 2 140

Bibliography

1. R. P. Bell, Acid-Base Catalysis, Oxford University Press, London, (1941), chap. VII. 2. L. P. Hammett, Physical Organic Chemistry, McGraw-Hill Book Co., Inc. New York, N. Y., (1940), p. 241.

3. W. J. C. Orr, Transactions of the Faraday Society, 2, 1033 (1936). 4. J. C. Jungers and K. Wirtz, Bulletin de la Societe Chimique de Belgique, 45, 679 (1 936).

5. J. C. Jungers and K. F. Bonhoeffer, Zeitschrift fur Physikalische CheLie A177, 460 (1 936).

6. K. F. Gelb, Zeitschrift fur Electrochemie l 125, 648 (1938).

7. A. J. Brodskii and L. V. Sulima, Doklady Academii Nauk SSSR, 74, 513 (1950).

8. F. W. James, J. S. Anderson and H. V. A. Briscoe, Nature, 139, 109 (1937).

9. J. S. Anderson, R. H. Purcell, T. G. Pearson, A. King, F. W. James, H. J. Emeleus, and H. V. A. Briscoe, Journal of the Chemical Society, (1937), 1492.

10. J. S. Anderson, N. L. Spoor and H. V. A. Briscoe, Nature, 139, 508 (1937). 11. F. J. Garrick, Nature, 139, 507 (1937). 12. A. G. Ogston, Journal of the Chemical Society, (1936), 1023.

13. J. R. Schaefgen, M. S. Newman and F. H. Verhoek, Journal of the American Chemical Society, 66, 1847 (1944).

14. C. J. Nyman, S. C. Fung and H. W. Dodgen, Journal of the American Chemical Society, 72, 1033 (1950). 15. W. E. Nelson and J. A. V. Butler, Journal of the Chemical Society, (1938), 957. 16. V. K. LaMer and J. Greenspan, Transactions of the Faraday Society, 33, 1266 (1937). 17. V. K. LaMer and S. Hochberg, Journal of the American Chemical Society, 61, 2552 (1939).

18. EiJiro Ogawa, Bulletin of the Chemical Society of Japan, 11, 367 (1936). 19. F. Roseburg and W. E. Van Heyningen, Industrial and Engineering Chemistry, Analytical Edition, 14, 36T(1942). 20. M. B. Williams and H. D. Reese, Analytical Chemistry, 22, 1556, (1950).

21. Organic Synthesis, Collective Volume II, John Wiley and Sons, New York, N. Y. (1941), p. 313. 22. Organic Synthesis, Collective Volume II, John Wiley and Sons, New York, N. Y. (1941), p. 318. 23. P. D. Bartlett and A. Schneider, Journal of the American Chemical Society, 67, 141 (1945). 142

VITA

Cyrus Henry Thomas was born on February 14, 1925 at Johnson City,

New York. He is the son of Mildred Evelyn (nee Nicholas) and Stewart

Thomas. After attending Frank M. Smith School and C. Fred Johnson Junior High School, he graduated from Johnson City High School in June of 1942. In July of that year he took a job as a "sole breaker" with Endicott Johnson Shoe Corporation. In September he left this job to take a position as a Chemists' Helper at Ansco Corporation in

Binghamton, New York. In June, 1943 he left this job to enlist in the United States Navy. After taking "Boot Training" at Sampson Naval Training Station he was sent to aviation radioman and radar school at Memphis, Tennessee. In January of 1944 he was sent to the University of the South at Sewanee, Tennessee to begin officers training. After spending his freshman year there, he was transferred to the Naval

Reserve Officer Training Corp. at the Georgia Institute of Technology, from which he graduated in June of 1946 with a Bachelor of Science degree in Naval Science and an Ensign's commission. In August of that year he married Colleen Nicholson. After serving one and one-half years with the fleet, he resigned his commission and was placed on inactive duty. In January of 1948 he returned to his home in Johnson

City and his old job at Ansco to wait for the new school year to open. In September of 1948 he re-entered Georgia Tech. and obtained the 143

Bachelor of Science degree in Chemistry in June of 1950. He immediately entered graduate school at the same institution and in September became a Research Corporation of New York Fellow. The following year he became an Atomic Energy Commission Research Assistant, which position he retained until his graduation with the degree of Doctor of Philosophy in the School of Chemistry in January of 1953. He is now employed as a research chemist in the Chemical Division of the Standard Oil Development

Company at Linden, New Jersey.