In Presenting the Dissertation As a Partial Fulfillment Of

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In presenting the dissertation as a partial fulfillment of • the requirements for an advanced degree from the Georgia Institute of Technology, I agree that the Library of the Institute shall make it available for inspection and circulation in accordance with its regulations governing materials of this type. I agree that permission to copy from, or to publish from, this dissertation may be granted by the professor under whose direction it was written, or, in his absence, by the Dean of the Graduate Division when such copying or publication is solely for scholarly purposes and does not involve potential financial gain. It is understood that any copying from, or publication of, this dissertation which involves potential financial gain will not be allowed without written permission.

3/17/65 b

LTh PART I THE RELATIVE REACTIVITIES OF COMMON NUCLEOPHILIC REAGENTS TOWARD DIFLUOROMETHYLENE--A SEARCH FOR BIFUNCTIONAL CATALYSIS

PART II THE REACTIONS OF THE METHYLENE HALIDES WITH ALKOXIDES IN ALCOHOLIC SOLVENTS--A - SEARCH,FOR AN a-ELIMINATION MECHANISM AND METHYLENE INTERMEDIATES

PART III THE REACTION OF METHYLMAGNESIUM BROMIDE WITH BENZOPHENONE-- THE MECHANISM OF THE GRIGNARD REACTION

A THESIS

Prethented .to

The Faculty of the Graduate Division

by 4/

Roy lAuke

In Partial Fulfillment

of the Requirements for the Degree

Doctor of Philosophy in the

School of Chemistry

Georgia Institute of Technology

April, 1967 Original Page Numbering Retained.

PART I THE RELATIVE REACTIVITIES OF COMMON NUCLEOPHILIC REAGENTS TOWARD DIFLUOROMETHYLENE--A SEARCH FOR BIFUNCTIONAL CATALYSIS

PART II THE REACTIONS OF THE METHYLENE HALIDES WITH ALKOXIDES IN ALCOHOLIC SOLVENTS--A SEARCH FOR AN ccELIMINATION MECHANISM AND METHYLENE INTERMEDIATES

PART III THE REACTION OF METHYLMAGNESIUM BROMIDE WITH .BENZOPHENONE-- THE MECHANISM OF THE GRIGNARD REACTION

Approved:

■■•

Date app ved by Chairman I 1 6 ACKNOWLEDGMENTS

The author wishes to thank Drs. Jack Hine and E. C. Ashby for their guidance and encouragement throughout the course of this work.

The helpful comments offered by Dr. H. M. Neumann while serving on the reading committee are also appreciated.

Financial assistance from the Army Ordinance Department and the

Alfred P. Sloan Foundation is gratefully acknowledged.

The author also wishes to express his gratitude to his wife for her understanding and help at every step of the way. TABLE OF CONTENTS

Page

ACKNOWLEDGMENTS ii

LIST OF TABLES. ix

LIST OF ILLUSTRATIONS xviii

SUMMARY xxii

PART I

Chapter

I. INTRODUCTION 1

Background The Basic Hydrolysis of the Haloforms The Basic Hydrolysis of Difluarohalomethanes Bifunctional Capturing Agents

Objectives Approach

II. EXPERIMENTAL 8

Instrumentation pH Measurements and. Potentiometric Titrations Constant-Temperature Bath Nuclear Magnetic Resonance Spectra Infrared Spectra

Chemicals

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Various Nucleophilic Reagents Apparatus and Procedure

Analyses Hydroxide, Phenoxides, and Sulfides Chloride Fluoride Catechol Mono(Difluoromethyl) Ether ,iv

Table of Contents (continued) Page

Preparations Preparation of Catechol Mono(Difluoromethyl) Ether Reaction of Chlorodifluoromethane with Phenoxide in Deuterium Oxide Identification of Reaction Products

III'. DISCUSSION AND RESULTS 18

The Basic Hydrolysis of Chlorodifluoromethane Determination of the Product Distribution from the Stoichiometry Determination of the Relative Rate Constants The Basic Hydrolysis of Chlorodifluoromethane Comparison of the Basic Hydrolyses of Chlorodifluoro- methane and Chloroform Rate Constants and the Appropriate Concentration Terms Carbon Monoxide as a Reaction Product

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Monofunctional Reagents The Reaction of Difluoromethylene with Fluoride The Reaction of Difluoromethylene with Azide The Reaction of Difluoromethylene with Cyanide

The Relative Reactivities of Azide, Cyanide, Hydroxide, and Fluoride Summary The Reaction of Difluoromethylene with Phenoxide The Reaction of Difluoromethylene with 2,4- Dichlorophenoxide The Reaction of Difluoromethylene with p-Methyl- thiophenoxide

The Relative Reactivities of Some Nucleophilic Reagents Toward Difluoromethylene

The Relative Reactivities,of Water and Hydroxide toward Difluoromethylene The Reaction of Difluoromethylene with Phenoxide

The Reaction of Difluoromethylene with Phenoxide in Deuterium Oxide..

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Bifunctional Reagents The Reaction of Difluoromethylene with Catechol Monoanion The Reaction of Difluoromethylene with Pyrogallol

ii Table of Contents (continued) Page

The Reaction of Difluoromethylene with 2-Mercaptoethanol

The Basic Hydrolysis of Chlorodifluoromethane in 60 Percent Ethanol The Reaction of Difluoromethylene with Hydroxide in 60 Percent Ethanol The Reaction of Difluoromethylene with Phenoxide The Reaction of Difluoromethylene with m-Chloro- phenoxide The Reaction of Difluoromethylene with p-Chloro- phenoxide

IV. CONCLUSIONS 72

APPENDIX 74

LITERATURE CITED 99

PART II

I. INTRODUCTION 101

Background The Reactions of Mono-, Di, and Trihaiomethanes with Strong Bases

Purpose Approach

II. EXPERIMENTAL 108

Instrumentation Ultraviolet Spectra Nuclear Magnetic Resonance. Spectra Infrared Spectra Potentiometric Titrations Constant-Temperature Bath

Chemicals

Reaction of the Methylene Halides with Alkoxides Determination of the Rates of Formal Formation Preparation of-Di(tert-Butyl)- and Diisopropyl Formal

Determination of the Stoichiometry of. the Reactions of Potassium tert-Butoxide and Isopropoxide with the Methylene Halides

I 17 vi

Table of Contents (continued) Page

Deuterium Exchange of the Methylene Halides The Nuclear Magnetic Resonance Method The Near-Infrared Method- The Infrared Method

Determination of the Deuterium Content of Exchanged Methylal Preparation of Deuterium-Exchanged. Methylal : Isolation of Deuterium-Exchanged Methylal Determination of the CH 2 /CH 3 Ratio of Deuterium- Exchanged Methylal

Preparation of Methyl Alcohol-d Procedure

Preparation of Isopropyl and tert-Butyl Alcohol-d Procedure

,III. DISCUSSION AND RESULTS ., . • 133

Readtions of the Methylene Halides with Alkoxides The Rates of Fotmal Formation

Reaction of Chloroform with Alkoxides The Rate of Reaction

Reactions of the Methylene Halides with Alkoxides The Solvent Kinetic Isotope Effect The Rates, of Deuterium Exchange Derivation of the Rate Equations for Deuterium Exchange Results of the Study of Deuterium Exchange Study of the .Deuterium Content of the Formal The Mass-Law Effect Product and Stoichiometry Studies

IV• CONCLUSIONS 158

APPENDIX 160

LITERATURE CITED 231

vii

Table of Contents (continued)

Page

PART III

INTRODUCTION 234

Background The Reactive Grignard Species Association of the Grignard Reagent The Grignard-Ketone Complex Kinetics of the Grignard Reaction Reaction Mechanism

Purpose Approach

II. EXPERIMENTAL 243

Apparatus Instrumentation Constant-Temperature Bath Timer Inert-Atmosphere Box Reaction Flasks

Chemicals

Methods for Following the Grignard Reaction

Kinetic Studies Preparation of the Reaction Flasks and Addition of Benzophenone Preparation of Methylmagnesium Bromide Solutions Initiating, Timing, and Quenching of the Reactions Analyses

Extinction Coefficients

Preparations Methylmagnesium Bromide

Inert-Atmosphere Box and Purification System Purification System Preparation of Manganese(II) Oxide viii

Table of Contents (continued) Page

III. DISCUSSION AND RESULTS 258

Kinetic Studies The Grignard Mechanism The Meisenheimer Mechanism The Dimer Mechanism The Swain Mechanism

IV. CONCLUSIONS 286

APPENDIX 288

LITERATURE CITED 321

VITA 324

LIST OF TABLES

Table Page

PART I

1. The . Basic Hydrolysis of Chlorodifluoromethane at 36° . . . 22

2, The Basic Hydrolysis of Chlorodifluoromethane at 36° . 32

3. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Fluoride at 36° . . 33

4. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Azide at 36° 34

5. Comparison of Directly and Indirectly Determined Fractions of Difluoromethyl Azide 36

6. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Cyanide at 36° 39

7. The Relative Reactivity of Azide, Cyanide, Hydroxide, and Fluoride toward Difluoromethylene 41

8. The Basic Hydrolysis of Chlorodifluoromethane in the. Presence of Phenoxide at 36° 42-

9. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2,4-Dichlorophenoxide at 36° 44

10. The, Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Methylthiophenoxide at 36° 45

11. The Relative Reactivities of Some Nucleophilic Reagents Toward Difluorotethylene 46

12. The Basic Hydrolysis of Chlorodifluoromethane in the rresence of Catechol at 36° 56

13. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36° 56

14. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36° 57

15. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36° 57 List of Tables (continued)

Table Page

16. Direct Determination of Catechol Mono(Difluoromethyl)Ether . 59

17. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2-Mercaptoethanol at.36° 63

18. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide in 60 Percent Ethanol 69

19. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of m,-.Chlorophenoxide in 60 Percent Ethanol at 25° . 69,

20. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Chlorophenoxide in 60 Percent Ethanol at 25° . 71

21. The Basic Hydrolysis of Chlorodifluoromethane at 36° . . . . 75

22. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Fluoride at 36° 76

23. The Basic Hydrolysis of Chlorodifluoromethane in the Presence, of Azide at 36° 77

24. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Cyanide at 36° 78

25. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide at 36° 79

26. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2,4-Dichlorophenoxide at 36° 80

27. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Methylthiophenoxide at 36° 81

28. The Basic Hydrolysis,of Chlorodifluoromethane in the Presence of Phenoxide at elso 82

29. The Basic Hydrolysis of ChlorOdifluoromethane in Deuterium Oxide in the Presence of Phenoxide at 36° 83

30. The Basic Hydrolysis of Chlorodifluoromethane in the Presence. of Catech of Dianion at 36° 84

31. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol Monoanion at 36° 85

List of Tables (continued)

Table Page

32, The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol Monoanion at 36° 86

33. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Pyrogallol Mono- and Dianions at 36° 87

34. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2-Mercaptoethanol at 36° . . . . 88

35. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 60 Percent Ethanol at 25°

36. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide in 60 Percent Ethanol at 25° 90

37.' The Basic Hydrolysis of Chlorodifluoromethane in the Presence of m-Chlorophenoxide in 60 Percent Ethanol at,25° . 91

38. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Chlorophenoxide in 60 Percent Ethanol at 25° 92

PART II

1, Relative Rates of Reaction of the Methylene Halides with Iodide in Acetone and Methoxide in Methanol at 50° 103

2. Rate Constants for the Reaction of the Methylene Halide with Alkoxides in Alcoholic Solvents at 36 0; ,, . . 134

3. Relative Reactivities of Several Methylene Halides with Methoxide, Isopropoxide, and tert-Butoxide 137

4. Reaction of Chloroform, with Methoxide, Isopropoxide, and teelt-Butoxide in the Corresponding Alcohols at 36° 138

5. The Solvent Kinetic Isotope Effect in the Reactions of the Methylene Halides with Alkoxides 141

6. Rate Constants for Deuterium Exchange of the Methylene Halides 143

7. Deuterium Content of the Methylal Produced in the Reaction of Methylene Bromide with Methoxide in Methyl Alcohol-d . . 150 xii

List of Tables (continued)

Table Page

8. Reaction of Methylene Iodide with Methoxide in the Presence of Iodide and Perchlorate 152

9. Stoichiometry of the ReaCtions of Isopropoxide-and tert-Butoxide with the :Methylene Halides 157

10. Extinction Coefficients for Methyl Alcohol in Methyl Alcohol-d 167

11. Extinction Coefficients for Isopropyl Alcohol in Isopropyl Alcohol-d 168

12. Extinction Coefficients for tert-Butyl Alcohol in tert- Butyl Alcohol-d . . . ..... ..... ...... 169

13. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol -d at 36° 170

14. Reaction of Methylene,Iodide with Sodium Methoxide in Methyl Alcohol-d at 36° 171

15. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36° 172

16. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36° 173

17. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36° 174

18. Reaction of Methylene Iodide with Potassium Isopropoxide in Isopropyl Alcohol at 36° 175

19. Reaction of Methylene Iodide with Potassium Isopropoxide in Isopropyl Alcohol at 36° 176

20. Reaction of Methylene Iodide with Potassium Isopropoxide in Isopropyl Alcohol-d at 36° 177

21. Reaction of Methylene Iodide with Potassium tert-Butoxide, in tert-Butyl Alcohol at 36° 178

22. Reaction oflilethylene Iodide with Potassium tert,Butoxide in tert-Butyl Alcohol-d at 36° 179

List of Tables (continued)

Table. Page

23. React;ion of Methylene Bromide with Sodium Methoxide in Methyl Alcohol-d at 36° 180

24. Reaction of Methylene Bromide with Sodium Methoxide in Methyl Alcohol-d at 36° 181

25. Reaction of Methylene Bromide with Potassium Isopropoxide in Isopropyl Alcohol at 36 ° 182

26. Reaction of Methylene Bromide with Potassium Isopropoxide in Isopropyl Alcohol-d at 36° 183

27. Reaction of Methylene Bromide with Potassium tert-Butoxide in tert-Butyl Alcohol at 36° 184

28. Reaction of Methylene Bromide with Potassium tert-Butoxide in tert-Butyl Alcohol at 36° . . . 185

29. Reaction of Methylene Bromide with Potassium tent-Butoxide in tert-Butyl Alcohol-d at 36° 186

30. Reaction of Methylene Chlorobromide with Sodium Methoxide in Methyl Alcohol-d at 36° 187

31. Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol at 36° 188

32. Reaction of Methylene Chlorobromide with Potassium Isopropoxide in Isopropyl Alcohol-d at 36° 189

33. Reaction of Methylene Chlorobromide with Potassium tert-Butoxide in tert-Butyl Alcohol at 36° 190

34. Reaction of Methylene Chlorobromide with Potassium tert-Butoxide in tert-Butyl Alcohol-d at 36° 191

35. Reaction of Methylene Chloride with Sodium Methoxide in Methyl Alcohol-d at 36° 192

36. Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol at 36° 193

37. Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol-d at 36° 194

xiv

List of Tables (continued)

Table Page

38. Reaction of Methylene Chloride with Potassium tert- Butoxide in tert-Butyl Alcohol at 36° 195

39. Reaction of Methylene Chloride with Potassium tert- Butoxide in tert-Butyl Alcohol at 36° 196

40. Reaction of Chloroform with Sodium Methoxide in Methyl Alcohol at 36° 197

41. Reaction of Chloroform with Potassium Isopropoxide in Isopropyl Alcohol at 36° 198

42. Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d at 36° 199

43. Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d at 36° 200

44. Deuterium Exchange of Methylene Iodide in Methyl Alcohol-d at 36° . . . 201

45. Deuterium Exchange of Methylene Chloride in Isopropyl Alcohol-d at 36° . . . 202

46. Deuterium Exchange of Methylene Chlorobromide in Isopropyl Alcohol-d at 36° 203

47. Deuterium Exchange of Methylene Bromide in Isopropyl Alcohol-d at 36° 204

48. Deuterium Exchange of Methylene Bromide in Isopropyl Alcohol-d at 36° 205

49. Deuterium Exchange of Methylene Iodide in Isopropyl Alcohol-d at 36° 206

50. Deuterium Exchange of Methylene Chloride in tert-Butyl Alcohol-d at 36° 207

51. Deuterium Exchange of Methylene Chlorobromide in tert- Butyl Alcohol-d at 36° 208

52. Deuterium Exchangeof Methylene Chlorobromide in tert- Butyl Alcohol-d at 36° 209 xv

List of Tables (continued)

Page

53. Deuterium Exchange of Methylene Bromide in tert-Butyl Alcohol-d at 36° 210

54. Deuterium Exchange of Methylene Bromide in tert-Butyl Alcohol-d at 36° 211

55. Deuterium Exchange of Methylene Iodide in tert-Butyl Alcohol-d at 36° 212

56. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Iodide at 36°. . . 213

57. Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Iodide at 36° . . 214

58. Reaction of Methylene Iodide with Sodium Methoxide in Methanol-d in the Presence, of 0.5 M Sodium Perchlorate at 36° 215

59. Reaction of Methylene Iodide with Sodium Iodide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Perchlorate at 36° 216

60. Stoichiometry of. the Reaction of Potassium tert-Butoxide with Methylene Bromide 217

61. Stoichiometry of the Reaction of Potassium Isopropoxide with Methylene Chlorobromide 217

PART III

1. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at High Grignard-to-Ketone Ratios at 25° . 259

2. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at High Crignard-to-Ketone Ratios at 25° . 260

3. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at Low Grignard-to•Ketone Ratios at 25° . 283

4. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 295

5. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 296

xvi

List of Tables (continued)

Table Page

6. Reaction of Methylmagnesium Bromide with Benzophenone-in Ether at 25° 297

7. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° ...... . ..... ...... 298

8. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 299

9. ReaotiOn, of Methylmagnesium Bromide with Benzophenone in Ether at 25° 300

10. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 301

11. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 302

12. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 303

13. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 304

14. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 305

15. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 306

16. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 307

17. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 308

18. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 309

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 310

20. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 311 xvii

List of Tables (continued)

Table Page

21. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 312

22. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 313

23. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 314

24. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° ...... ...... 315

25. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 316

26. Reaction of Methylmagnesium Bromide with Benzophenone in Eth er at 259 ...... ...... 317

27. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° 318

28. Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25° ...... . . . 319

r LIST OF ILLUSTRATIONS

Figure Page

PART I

1. Stoichiometry of the Basic Hydrolysis of Chlorodifluoro- methane at 36° 25

2. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Azide at. 36° 38

3. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Pheroxide at 36° 43

4. The Basic Hydrolysis of Chlorodifluoromethane in 60 Percent Ethanol at 25° 65

5. The Basic Hydrolysis of Chlorodifluoromethane in 60 Percent Ethanol at 25° 66

6. Infrared . Spectrum of Mono(Difluoromethyl) Catechol (in Carbon Tetrachloride). Instrument Settings: Resolution, 990; Response, 2; Gain, 4; Suppression, 4. 93

7. Nuclear Magnetic Resonance Spectrum of Catechol Mono(Di- fluoromethyl) Ether (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.0. External TMS Reference. 93

8. Nuclear Magnetic Resonance Spectrum of Catechol Mono(Di- fluoromethyl) Ether (in Carbon Tetrachloride). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.0. External TMS Reference 94

9. Nuclear Magnetic Resonance Spectrum of Deuterodifluoromethyl Phenyl Ether (in Carbon Tetrachloride). Instrument Settings: Filter Band Width, 1.0 cps.; R.F. Field, 0.20 mG.; Sweep Width, 500 spc.; Sweep Time, 500 sec.; Spectrum Amplitude, 25. External TMS Reference. 94

10. Infrared Spectrum of Difluoromethyl Azide (in Carbon Tetra- chloride). Instrument Settings: Resolution, 927; Response, 2; Gain, 6 5 95 xix

List of Illustrations (continued)

Figure Page

11. Infrared Spectrum of Difluoromethyl Phenyl Ether (in Carbon Tetrachloride). Instrument Settings: Resolution, 927; Response, 2; Gain, 5; Suppression, 5 95

12. Nuclear Magnetic Resonance Spectrum of Difluoromethyl Phenyl Ether (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.,; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 0.63. External TMS Reference 96

13. Infrared Spectrum of Difluoromethyl p-Tolyl Sulfide (in Carbon Tetrachloride). Instrument Settings: Resolution, 927; Response, 2; Gain, 6 5 96

14. Nuclear Magnetic Resonance Spectrum of Difluoromethyl p-Tolyl Sulfide (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec. External TMS Reference. 97

15. Infrared Spectrum of Difluoromethyl m-Chlorophenyl Ether (in Carbon Disulfide). Instrument Settings: Resolution, 930; Response, 2; Gain, 5; Suppression, 4 97

16. Nuclear Magnetic Resonance Spectrum of Difluoromethyl m-Chlorophenyl Ether (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.25. External TMS Reference. 98

17. Infrared Spectrum of Difluoromethyl p-Chlorophenyl Ether (in Carbon Disulfide). Instrument Settings: Resolution, 930; Response, 2; Gain, 5; Suppression, 4 98

PART II

1. Attempted Determination of the Rate of Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d by Nuclear Magnetic Resonance Spectroscopy. 218

2. Infrared Spectrum of Diisopropyl Formal (Neat). 219

3. Infrared Spectrum of Di(tert-Butyl) Formal (Neat) 219 xx

List of Illustrations (continued)

Figure Page

4. Infrared Spectrum of Methyl Alcohol in Methyl Alcohol-d Determination of the Absorbance of the 0-H Absorption byl\the Empirical Ratio Method 22Q

5. The Effect of the Concentration of Sodium Methoxide on theExtinctionCoefficientsofMet 11371- A1c01101. 5m Methyl Alcohol-d 221

6. The Effect of the Concentration of Potassium Isopropoxide on the Extinction Coefficients of Isopropyl Alcohol in Isopropyl Alcohol-d 222

7. The Effect of the Concentratfti of Potassium tert-Butoxide on the Extinction Coefficient of tent-Butyl Alcohol in tert-Butyl Alcohol-d. , 223

8. Nuclear Magnetic Resonance Spectrum of Di(tert-Butyl Formal (Neat) 224

9. Nuclear Magnetic Resonance Spectrum of Diisopropyl Formal (Neat). 224

10. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d. Data Shown in- Table 43. 225

11. Calculation of the Rate of Constants for Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d. Data, Shown in Table 43. 226

12. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methanol-d. Data Shown in Table 43. 227

13. Calculation of the Rate Constants for Deuterium Exchange of. Methylene Bromide in Methanol-d. Data Shown in Table 43. . 228

14. Calculation of the. Rate Constants for Deuterium Exchange of Methylene Bromide in_Methanol-d. Data Shown in Table 43. . 229

15. Calculation of the RAte Constants for Deuterium Exchange of Methylene Bromide in Isopropyl Alcohol-d. Data shown in Table 48. 230 ••

xxi

List of Illustrations (continued)

Figure Page

PART III

1. Inert-Atmosphere Box and Recirculating Purification System. 254

2. Calculation of Pseudo-First-Order Rate Constants for the Reaction of Methyl Magnesium Bromide with Benzophenone. Determination of Initial Absorbance A 261 , 1, 3. The Effect of the Concentration of Methylmagnesium Bromide on the Degree of Association in Diethyl Ether (Courtesy of E. C. Ashby and D. White) 265

4. Graphical Test of, the Grignard Mechanism at High Grignard- to-Ketone Ratios by Equation (13). Data from Table 2. . . . 268

5. Graphical Test of the Meisenheimer Mechanism at High Grignard-to-Ketone Ratios by Equation (16). Data from Table 2 270

6. Graphical Test of the Dimer Mechanism at High Grignard-to- Ketone Ratios by Equation (19). Data from Table 2. . . . 274

7. Graphical Test of the Swain MetOlanism at High Grignard-to- Ketone Ratios by Equation (22).' Data from Table 2. . . . 278

8. Calculation of the Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at Low Grignard- to-Ketone Ratios with the Modified Swain Stoichiometry. Determination of the Relationship between C and P, Derivation 3, Step 3. Data from Table 24, Appendix 320 SUMMARY

PART I THE RELATIVE REACTIVITIES OF COMMON NUCLEOPHILIC REAGENTS TOWARD DIFLUOROMETHYLENE--A SEARCH FOR BIFUNCTIONAL CATALYSIS

The relative reactivities of azide, cyanide, hydroxide, fluoride, p-methylthiophenoxide, and several phenoxides toward difluoromethylene have been determined. The reaction was studied by generating difluoromethylene (by the basic hydrolysis of chlorodifluoromethane) in the presence of the various nucleophilic reagents and then determining the product distribution.

The results have shown that the relative reactivities do not follow the Swain-Scott nucleophilicity pattern and that the most reactive nucleophile, p-methylthiophenoxide, is only 30 to 40 times as reactive as the least reactive, fluoride. The small differences in reactivity no doubt refledt the high reactivity (low stability) of the methylene intermediate. The deviation from the Swain-Scott nucleophilicity pattern was unexpected since Hine and co-workers had previously shown that the halogens, chloride, bromide, and iodide, capture dichloromethylene in the same order as they perform other nucleophilic displacements at carbon.

The relative reactivities of bifunctional reagents, such as catechol monoanion, have also been determined. These reagents were investigated to ascertain if labile hydrogens on a site adjacent to a nucleophilic center would improve their reactivity toward difluoromethylene as compared to monofunctional reagents, such as phenoxide. The results have shown that the bifunctional reagents are no better at capturing the methylene intermediate than are the monofunctional reagents. The bifunctional capturing concept (as presented in the text), therefore, is either not operating or unobservable under the reaction conditions.

PART II THE REACTIONS OF THE METHYLENE HALIDES WITH ALKOXIDES IN ALCOHOLIC SOLVENTS--A SEARCH FOR AN a-ELIMINATION MECHANISM AND METHYLENE INTERMEDIATES

The reactions of methylene iodide, bromide, chloride, and chlorobromide with sodium methoxide and potassium isopropoxide and tert- butoxide in the corresponding protio- and deuteroalcohols have been studied to determine if the reactions proceed by the S N2 reaction mechanism or whether there is a measurable contribution from an a-elimination mechanism.

The rate constants for formal formation have been determined in both the protio- and deuteroalcohols. In all cases, the reactions yield the corresponding formals, i.e., methylal, diisopropyl formal, and di(tert-butyl) formal. The order of reactivity of the alkoxides toward a given methylene halide was found to be isopropoxide > tert-butoxide > methoxide. The relative reactivities were ,about 4:2:1, respectively.

The rate constants for the reaction of the alkoxides with chloroform, a substance known to react by an a-elimination mechanism, were also determined. The order of reactivity of the alkoxides toward the haloform was found to be tert-butoxide > isopropoxide > methoxide. The relative reactivities were about 10,000, 710, and 1.0, respectively. The observed order of reactivity of the alkoxides toward the methylene halides and xxiv

the small differences in reactivity suggest that the formals are formed largely by a nucleophilic displacement process.

The rate constants for deuterium exchange of the methylene halides in methyl alcohol-d, isopropyl alcohol-d, and tert-butyl alcohol-d, catalyzed by the conjugate bases of the alcohols, were also determined.

The order of reactivity of the alkoxides in promoting deuterium exchange was found to be tert-butoxide > isopropoxide > methoxide. The relative reactivities were about 7000, 300, and 1.0, respectively. The observed order is that expected from the basicities of the alkoxides. It will also be noted that the order of reactivity and the differences in the observed rate constants parallel those observed for the reaction8 of the alkoxides with chloroform. Since removal of a proton is the initial step in both reactions, agreement was expected.

The deuterium content of the methylene group of the methylal produced by the reaction'of methylene bromide with sodium methoxide in methyl alcohol-d was also examined for clues as to the reaction mechanism.

The results showed that the deuterium content of the methylal was more closely,related to that predicted assuming the methylene halide first undergoes deuterium exchange with the solvent and subsequently reacts with methoxide by an S N 2 reaction mechanism to give the formal. The deuterium content was not nearly as large as that predicted by the a-elimination mechanism.

The reaction of methylene iodide with sodium methoxide in methyl alcohol-d in the presence of added sodium iodide was studied to ascertain if there was a measurable decrease in the rate constants that could be ascribed to a mass-law effect. The results showed that no change in the xxv

rate constants could be detected indicating the absence of a mass-law effect. These results suggest that monoiodomethylene is not an inter-

Mediate in the reaction and add further support to our earlier conclusion that the reactions are largely S 2 in character. N

PART III THE REACTION OF METHYLMAGNESIUM BROMIDE WITH BENZOPHENONE—THE MECHANISM OF THE GRIGNARD REACTION.

The kinetics of the reaction of methylmagnesium bromide with benzophenone in diethyl ether have been studied by a new experimental technique which promises to be applicable to a wide range of organometallic reactions. The reaction was studied at Grignard-to-ketone ratios ranging from 1.4 to 152/1 which represent the broadest range reported to date. Analysis of the kinetic data has shown that the reaction is best described in terms of the Swain mechanism with the added stipulation that complex formation is governed by an equilibrium constant, K1 , of about 1000.

K 1 G+ K C

k 2 C.+ G P + G

There is also the distinct possibility that the stoichiometry of the reaction as originally proposed by Swain may be incorrect, i.e., the active Grignard reagent may not be regenerated in the rate-determining step of the reaction, but may remain complexed to the carbinolate. In this form it presumably has a lower reactivity toward the ketone; however, more work is necessary to verify this point. xxvi

The kinetic equations derived from the modified Swain mechanism

give consistent values for the rate constants which are independent of

the Grignard-to-ketone ratios. In addition, the rate equations correlate

the data throughout the entire region wherein meaningful kinetic data

can be derived, i.e., between 15 and 75 percent reaction. The rate

equations derived from other reaction mechanisms that have been proposed

fail at one or both of these points. These observations lead us to believe that the mechanism of the Grignard reaction is best described

in terms of the modified Swain mechanism. PART I

THE RELATIVE REACTIVITIES OF COMMON NUCLEOPHILIC REAGENTS TOWARD , DIFLUOROMETHYLENE--A SEARCH FOR BIFUNCTIONAL CATALYSIS - 4

CHAPTER I

INTRODUCTION

Background

The Basic Hydrolysis of the Haloforms

Largely due to the work of Hine and co-workers (1-6) the mechanism of the basic hydrolysis of the haloforms and the existence of divalent-carbon intermediates in solution is now well established.

- 4 CHXYZ + OH H2O + CXYZ

CXYZ CXY + Z

CXY Products

The dihalomethylenes so produced are rapidly converted to products depending upon the solvent and the nature of the nucleophilic reagents present. Briefly, mechanism (1) is supported by the following observations

(7a, 8a): [1] chloroform (where X, Y, and Z are all chlorine) undergoes base-catalyzed aldol condensation with carbonyl compounds, [2] all of the haloforms, except those containing two fluorine atoms, undergo base catalyzed deuterium exchange more rapidly than -they hydrolyze, and [3] capture of the intermediate dihalomethylenes by various nucleophilic reagents including halogen illustrating, among other things, a mass-slaw effect (2).

Arguments based on the relative reactivities of the mono-, di-, tri-, and tetrahalomethanes with regard to the accepted mechanisms for nucleophilic displacement at carbon, i.e.,, S N1 and SN2 reaction mechanisms, also support the a-elimination mechanism. 2

The Basic Hydrolysis of Difluorohalomethanes

Hine and Langford (3, 8b) have presented convincing evidence that difluorohalomethanes, HCF 2 X (where X is Cl, Br, or I), undergo a concerted a-dehydrohalogenation whereby the intermediate trihalomethyl anion is by-passed.

HO- + HCF X F [H0...H. , -CF —.X] H2O + CF + X 2 2 2 (2)

In agreement with mechanism (2), the base-catalyzed hydrolysis of deuterodifluorobromomethane has been found to proceed without deuterium exchange with the solvent, i.e., isolation of the haloform after a sub- stantial amount had been hydrolyzed showed that its deuterium content had not decreased. It was proposed that by-passing the trihalomethyl anion could be explained by examining the role of the halogens in stabilizing the two intermediates: CXYZ and CXY. The ability of the halogens to stabilize the trihalomethyl anion (as indicated by the ratio of the rates of deuterium exchange of the haloforms to the corresponding rates of hydrolysis) stands in the order Br •••• I > Cl >> F. This order roughly reflects the ability of the halogen to expand its octet and to stabilize the carbanion by resonance (7b, 8c).

The abilities of the halogens to facilitate hydrolysis of the haloforms and presumably their abilities to stabilize dihalomethylenes, stand in the order F >> Cl > Br > I. This order is believed to result in part from the relative abilities of the halogens to stabilize the electron-deficient carbon atom by resonance (3, 8c). _e 7(-c-71 1 -)-(=c-71 Lx-c=71

- - 3

The Principle of Microscopic Reversibility. Application of the principle of microscopic reversibility (8d) to the reaction mechanism proposed for the basic hydrolysis of HCF2 X demands that reaction of the intermediate difluoromethylene with chloride, bromide, or iodide must be accompanied by simultaneous protonation, i.e., addition of the elements of the hydrohalic acid must be a concerted process. If another nucleophilic reagent, such as an alkoxide, is employed as the capturing agent, the principle of microscopic reversibility does not demand that nucleophilic attack by the alkoxide and hydrogen transfer be concerted (since this is not the exact microscopic reverse of the formation of difluoromethylene).

CF2 ROCF2H + RO RO1- C1- 2 1-1-10RZ[RO— • • .H.. .0R] 2

However, due to the inability of fluorine to stabilize a negative charge on the carbon to which it is bonded (3, 8c, 8e), hydrogen transfer must either be simultaneous with or closely follow nucleophilic attack; otherwise, the alkoxydifluoromethyl anion may lose fluoride to yield products characteristic of an alkoxyfluoromethylene (9, 8f)

RO ROR + CO + F F - - ROCF ROOF 2 RO ROH (RO)CH + F

The Capture of Difluoromethylene by Nucleophilic Reagents. The basic hydrolysis or alcoholysis of chlorodifluoromethane in the presence of several nucleophilic reagents has been studied in some detail by Hine and co-workers (9-13). The products reported depend upon the nature of the solvent and the nucleophilic reagents present. The majority of these 4

studies have been qualitative and have served primarily to identify the products from these reactions. In some cases, however, attempts have been made to determine the product distribution.

Hine and Porter (11, 12) reported that the basic alcoholysis of

chlorodifluoromethane in the presence of sodium thiomethoxide gives

difluoromethyl methyl sulfide, trimethyl orthothioformate, and dimethoxy- methyl methyl sulfide in 48, 28, and 6 percent yields, respectively.

CH0- CH S7 HCF Cl HCF SCH + (CH.S) CH + (CH 0) CHSCH 3 2 CH OH "2 3 3 3 3 2 3 3

In the absence of sodium thiomethoxide, the reaction yields 48 percent difluoromethyl methyl ether, 32 percent trimethyl orthoformate, and an unspecified amount of carbon monoxide.

CH 0 - 3 HCF Cl HCF OCH + (CH 0) CH + CO 2 CH OH 2 3 3 3 3

Hine and Tanabe (9) reported that the reaction of potassium isopropoxide with chlorodifluoromethane in isopropyl alcohol gives 5 percent fluoroform, 30 percent isopropyl orthoformate, and 58 percent difluoromethyl isopropyl ether based on the amount of haloform reacted.

iso-PrO - HCF Cl HCF + (iso-PrO) CH + HCF OCH(CH ) 2 iso-PrO 3 3 2 3 2

Miller and Thanassi (13) have shown that the basic hydrolysis of chlorodifluoromethane in the presence of sodium phenoxide leads to the formation of difluoromethyl phenyl ether in yields as high as 65 percent based on the phenoxide. 5

Ph0 + HCF2 Cl --+ PhOCF H + (Ph0) CH OH0H 2 3 2

Triaryl orthoformates were produced also. Miller and Thanassi employed a two-phase, aqueous-dioxane solvent system and maintained that this system was necessary in order to obtain good yields of the aryl difluoromethyl ethers. When water was employed as the reaction medium, only small amounts of the desired ethers were obtained and as much as 70 percent of the haloform was recovered unreacted.

The Relative Reactivities of Nucleophilic Reagents Toward

Difluoromethylene. The previously mentioned studies are quite useful in defining the major products to be expected from the basic hydrolysis or alcoholysis of chlorodifluoromethane in the presence of thiomethoxide, phenoxide, etc.; however, for the most part., they are not sufficiently quantitative to estimate the relative reactivities of the various nucleophilic agents toward difluoromethylene. Since the latter was of particular interest, a study was undertaken to define the relative reactivities of several common nucleophilic reagents toward the divalent-carbon intermediate.

Bifunctional Capturing Agents

As previously discussed, the principle of microscopic reversibility demands that the addition of the elements of HX (where X is Cl,

Br, or I) to difluoromethylene, i.e., the reverse of reaction (2), must be a conserted process.. Furthermore, the logical extension of this principle suggests that the capture of the divalent-carbon intermediate by other nucleophilic reagents should be concerted also. With these prin- ciples in mind, it seemed probable that bifunctional reagents, such as catechol monoanion, which contain an acidic proton adjacent to a nucleo- F

philic site might be markedly superior capturing agents compared to other phenoxides which lack the appropriately situated acidic proton.

Assuming that the capture of difluoromethylene must be concerted, in the first case the reaction is bimolecular whereas in the second, i.e., capture by phenoxide, the proton must come from the solvent which would necessitate a concerted termolecular reaction. Because of the growing interest in bifunctional catalysis, which in recent years has produced numerous articles (13, 14, 8g, El) showing how reagents with appropriately oriented reactive sites can accelerate the rates of certain chemical transformations (particularly in biochemical systems), it was decided to compare the relative reactivities of several bifunctional reagents (such as catechol monoanion) toward difluoromethylene with those of simple monofunctional reagents (such as phenoxide).

Objectives

The primary objective of this study was to determine if certain bifunctional reagents, such as catechol monoanion, were superior to monofunctional reagents at capturing difluoromethylene. In order to have a basis for comparison, it was first necessary to establish the relative reactivities of the monofunctional reagents toward the divalent-carbon intermediate. As the study progressed, it also proved of interest to ascertain the relative reactivities of several common nucleophilic reagents, such as azide, fluoride, and cyanide, and to determine if these 7

reagents showed the same order of reactivity toward difluoromethylene as they do in SN2-type displacements at carbon.

Approach

The relative rates at which reactive intermediates, such as difluoromethylene, are transformed to products by reaction with two or more reagents are often determined by generating the intermediate in the presence of the capturing agents and subsequently determining the distribution of products. If the rate- and product-controlling step of each reaction is the combination of the intermediate with the capturing agent, then from a knowledge of the ratio of products and the concentrations of the capturing agents, a ratio of rate constants can be obtained. Skell and Garner (16) and Doering and Henderson (17) employed this technique to determine the relative reactivities of various olefins toward dichloro- and dibromomethylene.

In the present study, difluoromethylene (produced by the basic hydrolysis of chlorodifluoromethane) was generated in the presence of various pairs of nucleophilic reagents. The amounts of the various products were determined indirectly from the stoichiometry of each reaction by following the hydroxide consumed and the chloride and fluoride produced. From the product distribution, the relative reactivities of the nucleophilic reagents toward difluoromethylene,were calculated. 8

CHAPTER II

EXPERIMENTAL

Instrumentation pH Measurements and Potentiometric Titrations

A Beckman Zeromatic pH meter equipped with glass-calomel electrodes was employed for all pH measurements and the potentiometric titrations with standard acetic acid. The silver-glass electrode system and the millivolt scale of the instrument were employed to determine the endpoints in the chloride analysis.

Constant-Temperature Bath

A constant-temperature water bath (Sargent No. S-84845) in conjunction with a Sargent Thermonitor (Sargent No. S-82055) was used for controlling the temperature (* 0.01°).

Nuclear Magnetic Resonance Spectra

The nuclear magnetic resonance (nmr) spectra were recorded on a

Varian Nuclear Magnetic Resonance Spectrometer, Model A-60. The instrument settings are recorded on each spectrum. All of the spectra were recorded with external tetramethylsilane referencing.

Infrared Spectra

The infrared spectra were recorded on a Perkin-Elmer Infrared

Spectrophotometer, Model 21. 9

Chemicals

Acetic Acid. Baker Analyzed.

3-(2-Arsonophenylazo)-4,5-dihydroxy-2,7-naphthalenedisulfonic Acid Disodium Salt. Eastman.

Carbon Tetrachloride. W. H. Curtin.

Catechol. Eastman.

Cerous Nitrate, Ce(NO ), • 6H n O. Fisher Scientific Co.

Chlorodifluoromethane. Matheson, Coleman and Bell Co. m-Chlorophenol. Eastman. p-Chlorophenol: Eastman.

Deuterium Oxide (Heavy Water). U.S. Atomic Energy Commission.

2,4-Dichlorophenol. Eastman.

(Ethylenedinitrilo)tetraacetic Acid Disodium Salt, EDTA Eastman.

Hydrogen Peroxide, 30 percent. Baker Analyzed.

2-Mercaptoethanol. Carbide and Carbon Chemicals Corporation.

Phenol. Baker Analyzed.

Potassium Cyanide, Baker Analyzed.

Potassium Fluoride. Baker Analyzed.

Potassium Hydroxide. Baker Analyzed. 1 0

Pyrogallol. Eastman.

Sodium Azide. Eastman.

Sodium Chloride. Baker Analyzed.

Sodium Hydroxide. Baker Analyzed.

Silver Nitrate. Baker Analyzed.

`p-Toluenethiol. Eastman.

The Basic Hydrolysis of Chlorodifluoromethane in the

Presence of Various Nucleophilic Reagents

Apparatus and Procedure

A 500 ml., three necked flask equipped with a stirrer and a

fritted-glass, gas-addition tube was immersed (up to the necks) in a

constant-temperature bath. The flask was purged with nitrogen and the

basic solution containing the nucleophilic reagent to be studied added.

Potassium hydroxide was employed as base in preference to sodium hydroxide

due to the greqter solubility of potassium fluoride in aqueous solutions.

The solution was then analyzed for both hydroxide and the capturing

(nucleophilic) reagent.

After ascertaining the initial concentrations of both reactants,

chlorodifluoromethane was passed into the solution with good stirring at

a rate of 0.04 to 0.08 cubic feet per hour (according to a flowmeter

calibrated for air) for a period of 15 to 45 minutes depending upon the

amount of reaction desired. The solution was then purged with nitrogen

and allowed to stir for an additional hour or more to insure complete

reaction of any dissolved haloform. The solution was analyzed 1 1

subsequently for hydroxide, capturing agent, fluoride, and chloride.

This procedure was repeated until sufficient data were obtained (usually four to five additions of chlorodifluoromethane) to evaluate the relative rate constants.

Analyses

Hydroxide, Phenoxides, and Sulfides

An aliquot (usually 10 ml.) of the reaction mixture was transferred to a small beaker and titrated potentiometrically with standard acetic acid (ca,, 1.0 M) employing a Beckman Zeromatic pH meter equipped with '

glass - calomel electrodes. The solutions were magnetically stirred and, when necessary (catechol and pyrogallol studies), the titrations were carried out under nitrogen to prevent oxidation of the reactants. Com- plete titration curves were then plotted and the endpoints determined graphically (18a).

The titration of a mixture of hydroxide and phenoxide yields two distinct breaks from which the respective concentrations of the two reactive species can be ascertained. Direct titration proved satisfactory for determining hydroxide and the various phenols and sulfhydryl compounds employed. Other methods had to be devised to determine azide and cyanide.

Acetic acid was chosen as the titrating agent in preference to a stronger mineral acid to prevent interference from the formate and fluoride formed in the reaction. Being weak bases, formate and fluoride are included in the weak base portion of the potentiometric titration curve if strong acids are employed. The use of acetic acid eliminates this interference. The relatively high concentrations of the standard acetic acid solutions, ca., 1.0 M, were necessary to effect sharp breaks 12

in the titration curves which facilitated the graphical determination of the endpoints.

Chloride

The chloride concentration of the reaction mixtures was determined by direct potentiometric titration with standard silver nitrate employing the silver-glass electrode system. In most cases, the sample previously employed for the determination of hydroxide and phenoxide (or other capturing agent) was used since it was already neutralized; otherwise, a separate sample was neutralized with nitric acid prior to titration.

After sufficient familiarization with the method, it was not necessary to plot the titration curves in order to determine the endpoint; the latter could be ascertained by dead-stop titration.

Chloride in the Presence of Sulfhydryl Compounds. In several cases, the capturing agent either obscured the endpoint or reacted with the silver nitrate in such a manner as to give erroneous results. In these cases, special techniques had to be developed. For example, it was found that both p-toluenethiol and 2-mercaptoethanol reacted with silver nitrate. To eliminate this interference, a 10 ml. aliquot of the basic reaction mixture was treated with 4 ml. of 30 percent hydrogen peroxide prior to analysis. This treatment oxidized the thiols to water-soluble compounds (presumably the corresponding sulfonic acids) which caused no interference in the chloride determination.

Chloride in the Presence of Catechol and Pyrogallol. Catechol and pyrogallol were found to interfere also with the chloride titration.

The interferences due to these compounds could be eliminated by treating a 10 ml. aliquot of the basic reaction mixture with 2 ml. of 30 percent 13

hydrogen peroxide. After acidifying the reaction mixture to a pH of five,

titration with silver nitrate gave sharp, reproducible endpoints.

Chloride in the Presence of Azide and Cyanide. In the azide and

cyanide studies, the interference of these halogenoids was eliminated by

treating a 10 ml. aliquot of the reaction mixture with 20 ml. of 1:1

nitric acid. The high acid concentration converted the interfering

anions to the corresponding conjugate acids, i.e., hydrazoic and hydro-

cyanic acids, thus allowing the direct potentiometric determination of

chloride. This method also provided a means of determining the azide and

cyanide concentrations. By titrating another sample of the reaction mixture under neutral conditions,the total amount of chloride plus the

halogenoid was obtained. The difference between the two titrations provided a measure of the concentration of azide or cyanide.

Fluoride

The fluoride concentration of the reaction mixtures was determined by a modification of the method of Yamamura, Kussy, and Rein (19).

Basically, this method effects the precipitation of fluoride as

cerium(III) fluoride byaddition of an excess of cerium(III) nitrate under controlled conditions. The cerium(III) fluoride is separated by centrifugation and an aliquot of the supernatant liquid is then analyzed

for unreacted cerium(III) by titration with standard EDTA.

The basic procedure was to add two drops of 1 M chloroacetic acid

to an aliquot of the reaction mixture and then to adjust the pH to

3 ± 0.5 by addition of dilute nitric acid. An aliquot (usually 10 ml.) of a standard solution of cerium(III) nitrate, ca., 0.075 M, was added

and the pH of the resulting solution adjusted to 1.75 with dilute nitric 14

acid and ammonium hydroxide. All pH measurements were made with a pH meter. The mixture was transferred to a 50-m1. volumetric flask (not diluting to volume) using water adjusted to a pH of 1.75 with nitric acid for rinses. The mixture was then digested on a steam bath to precipitate the cerium(III) fluoride.

After the solution cooled to room temperature, it was diluted to volume with water adjusted to pH 1.75 and mixed well. About 35 ml. was - centrifuged to separate the cerium(III) fluoride. An aliquot (usually

25 ml.) of the clear,supernatant liquid was transferred to a 200-ml., tall-form beaker followed by addition of 50 ml. of water, two drops of

0.2 percent cresol red indicator, and a sufficient amount of

3-(2-arsonophenylazo)-4,5-dihydroxy-2,7-naphthalenedisulfonic acid disodium salt (added as a one percent solid mixture in sodium chloride) to produce a purple color. The mixture was titrated immediately with standard EDTA (ca., 0.05 M) to a change in color from purple to orange- yellow. A potentiometric method of determining the endpoint was employed in some of the analyses using the mercury-drop electrode as described by Flaschka (18c). The concentration of fluoride was then calculated from equation (3):

[F-] - 3(bcd - 50ef) ab (3) where:

a = ml. of reaction mixture.

b = ml. of centrifuged solution.

c = molarity of cerium(III) nitrate.

d = ml. of cerium(III) nitrate. 15

e = molarity of EDTA.

f = ml. of EDTA.

Catechol Mono( Difluoromethyl) Ether

Analysis of catechol mono(difluoromethyl) ether was effected by infrared spectroscopy employing the C-F absorption (13, 20)

(A 8.70-8.75 p). The extinction coefficient of this absorption was max determined by dissolving 0.6430 g. of the purified ether in carbon tetrachloride and diluting to 10 ml. (0.4018 M). The absorption spectrum was recorded (Figure 6, Appendix) on Perkin-Elmer semilog paper

(No. 021-6304) and the absorbance of the C-F band (most intense band in the spectrum) calculated by the empirical ratio method (21) by taking the difference between the base line absorbance between 4.0 and 6.0 p and the maximum absorbance at 8.70 to 8.75 p. A 0.025 mm. cell was employed.

The extinction coefficient, was then calculated from equation (4):

_ A Cl (4 ) where A is the absorbance, 1 is the cell-path length in centimeters,: and

C is the molar concentration of the difluoromethyl aryl ether. The value of £ obtained was 621.

Preparations

Preparation of Catechol Mono(Difluoromethyl) Ether

Catechol, 27 to 29 g. (ca., 0.25 mole), and potassium hydroxide,

80 g. (1.2 moles), were dissolved in about 300 ml. of water and transferred to the reaction apparatus. Chlorodifluoromethane was then passed into the mixture for about ten hours. The reaction mixture, which was 16

still strongly basic to litmus, was transferred to a separatory funnel

and extracted with petroleum ether. The extracts were combined and dried

over calcium chloride. Catechol bis(difluoromethyl) ether was isolated

later from this mixture. The aqueous solution was then adjusted to a pH

between five and six and extracted again with petroleum ether.

After removing the petroleum ether from the acidic extracts, the

product was distilled on a micro--Vigreux column collecting about 10 ml.

of a water-white liquid, b.p., 55° to 65° (5 mm.). This product was

subsequently redistilled on the same Vigreux column collecting about 5 ml.

of a heart cut, b.p., 61° (5 mm.).

The nmr spectrum of catechol mono(difluoromethyl) ether is shown

in Figure 7, Appendix. The absorptions at 2.93, 4.14, and 5.35 T are due

to the proton of the -CF 2H group (J = 72). The proton of the 0-H group

is superimposed on the -CF 2H absorption at 4.14 T. The aromatic hydro-

gens absorb between 3.45 and 4.0 T. The fact that the 0-H absorption

is coincident with the proton absorption of the -CF 2H group is illustrated

in the nmr spectrum of the monoether in carbon tetrachloride as shown, in

Figure 8, Appendix. The 0-H absorption is shifted to higher field as

expected, ca., 6.3 T, and the integrated areas of the absorptions take on the expected 4:1:1 ratio.

Reaction of Chlorodifluoromethane with Phenoxide in Deuterium Oxide

Phenol, 10.3 g. (0.110 moles), 10.7 ml. of 9.35 M NaOD (0.100 moles),

and 190 ml. of deuterium oxide were mixed in a 500 ml. flask. After

analyzing the reaction mixture for phenoxide, the flask was purged with

chlorodifluoromethane and connected to a gas manometer which was used to keep a positive pressure (ca., 250 mm. of Hg) of the haloform on the 17

reaction mixture. The reaction was allowed to proceed for several days at room temperature with continual mixing effected by means of a magnetic stirrer. The pressure declined from day to day as the haloform was consumed and the initial pressure was maintained by periodic additions of chlorodifluoromethane. After several days, an aliquot of the reaction mixture was removed and analyzed for phenoxide and chloride and the respective fractions of formate and phenyl difluoromethyl ether calculated. The same apparatus and procedure were employed for the study of the reaction of phenoxide with chlorodifluoromethane in water.

Identification of Phenyl Difluoromethyl Ether. The reaction mixture was treated with sufficient aqueous sodium hydroxide to make the solution strongly basic and extracted with small portions of carbon tetrachloride. The nmr spectrum of the combined extracts was then recorded (Figure 9, Appendix). The spectrum shows clearly the absorbance due to the aromatic protons but the triplet due to the -CF 2H group is conspicuously absent (compare with Figures 12 and 8, Appendix). The absorptions at 8.75 and 9.95 T are due to silicone stopcock grease which dissolved in the carbon tetrachloride during the extraction process.

Identification of Reaction Products

In all of the relative rate studies, an attempt was made to identify the reaction products by spectroscopic methods. This was effected by extracting the reaction mixture (sometimes after acidifi- cation) with carbon tetrachloride. After drying the extracts over anhydrous sodium sulfate or other desiccant, the infrared and nmr spectra were recorded. The combination of the two spectral methods usually sufficed to identify the product. - - - - -

CHAPTER III

DISCUSSION AND RESULTS

The Basic Hydrolysis of Chlorodifluoromethane

The basic hydrolysis of chlorodifluoromethane has been reported to yield formate, carbon monoxide, and fluoroform (3, 9). The formate and

carbon monoxide apparently arise from hydrolysis of the intermediate difluoromethylene and fluoroform from capture of the divalent-carbon species by fluoride [produced by reactions (5) and (6)].

k 1 HCF C1 + 2H + 2 20 HCOO + 2F + Cl + 4H

k} + HCF C1 + H2 + 3H 2 O CO+ 2F + Cl

k - 3 HCF C1 + F HCF + Cl 2 3

Determination of the Product Distribution from the Stoichiometry

It will be noted that reactions (5), (6), and (7) differ with respect to hydroxide consumed and chloride and fluoride produced. These differences in stoichiometry can be used to calculate the relative amounts of formate, carbon monoxide, and fluoroform produced as shown in equations

(8), (9), and (10):

Amount of Formate • A[DH] - A[c1] - A[Fl] (8)

Amount of Carbon Monoxide = (4A[C1 - ] + 4A[F] - 3A[OH])/3 (9)

Amount of Fluoroform • (2,6[C1 - ] -/[F- ])/3 (10) 19

where A[OH- ], A[Cl],- and A[F- ] are the changes in the concentrations of hydroxide, chloride, and fluoride, respectively. It will also be recognized that the total amount of chlorodifluoromethane reacted is given by

A [Cl].- As a result, the respective fractions of the haloform going to each product may be ascertained :by dividing equations (8), (9), and (10) by A[Cl]- as shown in equations (11), (12), and (13):

f = (A[OH- ] - A[C1 - ] - - ] (11) 1 0EF- D/A[C1 f = (4A[C1- ] 4A[F- ] - 3AE0H (12) 2 - 1)/3A[C1- ]

f = - (13) 3 (2A[C1 - ] A[F- ])/ 3A[C1 - ]

2, and f are the fractions of formate, carbon monoxide, and .Where fl , f 3 fluoroform, respectively.

The values of A[OH - ] were calculated from equation (14) and the values of A[C1 - ] and A[F- ] from equation (15):

- [OH - ]n = AC0H-1 [08- ]n_1 (14) = A[X-] [X ]n - [X ]n-1 (15) where n is the n th addition of chlorodifluoromethane and X - is chloride or fluoride. When other nucleophilic reagents, such as azide, were employed, the changes in concentration were calculated from a relationship similar to equation (14) since their concentration also decreased as the

reaction progressed. It will be noted that changes in the concentrations ' of the measured quantities are regarded as positive numbers, even though in the case of hydroxide the concentration actually decreases during the course of the reaction. These considerations were also employed in the derivation of the algebraic equations for calculating the respective 20

fractions of the reaction products.

Determination of the Relative Rate Constants

As shown in equations (16), (17), and (18), the fractions of formate, carbon monoxide, and fluoroform can be expressed in terms of the appropriate rate constant and concentration term leading to each product:

f = k [0H- ]/(k [OH - ] + k [HC00 - ] + k [F- ]) (16) 1 1 1 2 3

f = k [HC00- ]/(k [OH - ] + k [HC00 3 [F- ]) (17) 2 2 1 2 - ] + k

f = k [F- ]/(k [OH- ] + k [HC00- ] + k3 [F- ]) (18) 3 3 1 2 assuming that:

= k [0H (19) 1 - ][CF2 ]

= ][CF2 ] (20)

v = k [F- ][CF (21) 3 3 21 where v1, v2, and v are the rates of formation of formate, carbon 3 monoxide, and fluoroform, respectively. At present, assignment of the concentration terms to a particular rate constant may seem quite arbitrary; however, these parameters must be fitted from the experimental results as will be discussed later.

From equations (16), (17), and (18), a ratio of rate constants can be derived by dividing one equation by the other. Of interest to us were

and k1/k3 the values of k1/k2 which may be obtained by dividing equation (16) by equations (17) and (18), respectively, as shown in equations (22) and (23). 21

k /k = f [HC00]/f [0H- ] 1 2 1 2 (22) k /k = f [F- ]/f [OH- ] 1 3 1 3 (23)

The concentration term associated with each calculated fraction is the average value during the period which the fraction was determined.

By substituting equations (11), (12), and (13) into (22) and (23), equations (24) and (25) are obtained. With these equations, calculation

and ki/k3 , can be made directly of the relative rate constants, }el/k2 from the experimental data.

k 1 3(A[OH- ] - [Cl - ]A - A[F- ])[HC00- ] k (4A[C1 - ] + 4A[F ])[OH (24) 2 - ] - 3A[OH - - ]

k1 3(A[OH - ] - 4[C1 - ] - A[F- ])[F- ] k - (2A[C1- ] - A[F- - ] (2 5) 3 ])[OH

The Basic Hydrolysis of Chlorodifluoromethane

The results of a study of the basic hydrolysis of chlorodifluoromethane are shown in Table 1. The data (Table 21, Appendix) were analyzed assuming that three products were formed: formate, carbon monoxide and fluoroform. The relative rate constants, k i/k2 and ki/k 3 , were calculated from equations (24) and (25), respectively.

As the results show, hydroxide is about nine times as reactive as fluoride and three to four times as reactive as formate toward difluoromethylene. 22

Table 1. The Basic Hydrolysis of Chlorodifluoromethane at 36°

Add'n. of [OH-] [F- ] [HC00- ] e flb c f d k /k k1 g HCF2C1 Avg. Avg. Avg. f2 1 2 1-3 1 0.836 0.086 0.042 1.000 0.015 -0.015 3.17 h 2 0.608 0.199 0.097 0.933 0.033 0.033 4.47 9.2 3 0.444 0.283 0.136 0.862 0.080 0.057 3.28 9.6 4a 0.211 0.400 0.186 0.662 0.247 0.090 2.36f 13.9f

Avg. 3.64 ±0.55 9.4 ±0.2

a b HCF2C1 additions 4 and 5 combined for calculations. Calcul ated from equation (11). c0alculated from equation (12). dCalculated from equation eCalculated from equation (24). (Omitted from average ([0H-] (13). avg < 0.4 M). gCalculated from equation (25). hNegative value.

Comparison of the Basic Hydrolyses of Chlorodifluoromethane and Chloroform

Examination of the calculated fractions of formate, carbon monoxide, and fluoroform in Table 1 shows that during the first addition

of chlorodifluoromethane, essentially all of the intermediate difluoromethylene was converted to formate. Furthermore, in each subsequent

addition, the larger fraction is always formate. In contrast, Robinson (22) has reported that the sole product from the basic hydrolysis of chloroform is carbon monoxide.

k HCC1 3 + H2O CO + 3C1 - + 3H+

In addition, he has shown that the formate, often observed as a product in this reaction, is formed from a secondary and much slower reaction of

carbon monoxide with hydroxide. Surprisingly, this reaction occurs at 23

hydroxide concentrations as low as 0.01 M. Robinson also maintains that the only way that significant quantities of formate can be produced is by running the reaction in sealed tubes; otherwise, the carbon monoxide will escape and formate formation will be minimized.

In contrast to the work of Robinson, there seems to be no doubt that formate is the principal product from the basic hydrolysis of chlorodifluoromethane. Not only do the results summarized in Table 1 support this conclusion, but all of the data which we will subsequently submit point to this fact. It will also be noted that all of the experimental work was carried out in flasks open to the atmosphere and no effort was made to contain any carbon monoxide that might have been produced. Due to the limited solubility of carbon monoxide in aqueous solutions and its slow reaction with hydroxide, it seems unlikely that the formate produced in our studies could arise from the secondary reaction as described by Robinson; the logical conclusion is that it is a primary reaction product.

Of course, an indirect method of determining the product distribution, as employed in this study, is always subject to question and evidence that the calculated amounts of products are indeed the actual amounts must, be presented. A real test of the method would be to calculate the amount of a particular species from the stoichiometry relationships and then isolate or independently determine its concentration by a direct method. This approach has been employed in several instances during the course of this investigatift and excellent agreement between the calculated and experimentally determined amounts has been obtained. These experiments will be discussed in their proper 24

place. It seems adequate for the present to point out that no other reaction of chlorodifluoromethane with aqueous base consumes the amount of hydroxide that formate formation does; consequently, if the stoichiometry studies show that essentially four moles of hydroxide are consumed for'each mole of chloride produced, there can be no question but that the formation of formate is the principal reaction. A test of the formate stoichiometry can be made by plotting hydroxide versus chloride from the data in Table 21, Appendix, as shown in Figure 1. The best straight line through the origin and the first three points, representing the first three additions of chlorodifluoromethane, has a slope of 3.90.

Pure formate stoichiometry demands a slope of 4.00 while carbon monoxide formation necessitates a slope of 3.00. The observed slope is certainly closer to that predicted for formate formation and deviation from the predicted 4.00 value could be explained by the formation of small quantities of fluoroform.

Rate Constants and the Appropriate Concentration Terms

In order to calculate the relative rate constants, k l/k2 and kl/k3 , the concentration term or terms associated with each rate constant have to be ascertained. As previously mentioned, this can only be done by experiment. Perhaps the best way to proceed is to pick out the concentration term or terms which are most likely to be associated with the particular rate constant in question and then calculate the ratios of rate constants from the experimental data. The concentration term giving the more constant set of values is the logical choice.

Selecting the concentration terms associated with the rate constants was somewhat simplified by the fact that association of the fluoride 25

M ide, x dro Hy

0.1 0.2

Chloride, M

Figure 1. Stoichiometry of the Basic Hydrolysis of Chlorodifluoro- methane at 36° 26

seemed self-evident and indeed proved to be correct. concentration with k 3 was then investigated by The concentration term associated with kI evaluating ki/k 3 from equation (25). The concentration terms tested for association with k were hydroxide and water. The results using hydroxide I are shown in Table 1.

Examination of the results in Table 1 shows that the average value of k /k is 9.4 ± 0.2. The value of from the fourth addition of 1 3 k1/k3 the haloform was omitted from the average because the average hydroxide concentration dropped below 0.4 M. As will be shown later, water begins to compare successfully with hydroxide at and below this level. If the concentration term associated with k l is water rather than hydroxide, the values of are 0.140 and 0.077 for the second and third additions of k1/k3 the haloform, respectively. This represents a twofold change. At average hydroxide concentrations above 0.4 M, this trend toward steadily decreasing values of the relative rate constants was observed whenever the concentration of water was associated with k 1 and indicates that, under these conditions, hydroxide is the reagent largely responsible for the conversion of difluoromethylene to formate.

Again, the basic hydrolysis of chlorodifluoromethane differs with that of chloroform. Robinson (22) has reported that water rather than hydroxide is the reagent which leads to the conversion of dichloromethylene to carbon monoxide. Unfortunately, in this case, Robinson's results do not seem to warrant his conclusions. The ratio of rate constants, , obtained by Robinson ranged from +24.0 to -26.8 and the kOH/kH 0 2 average of nine determinations gave a value of -0.2 from which it was concluded that the value was actually zero, thus making k H 0 >> kOH. Although it 2 27

is admittedly difficult to obtain precise values for the ratio of rate constants by the method employed by Robinson, it seems equally unwise to draw such concise conclusions from such a wide spread in the experimental results.

It is not necessary to assume that a contradiction exists because our results differ from those of Robinson; it must be kept in mind that two distinctly different species are being dealt with here: dichloromethylene and difluoromethylene. The latter is markedly more stable than the former. This is illustrated by the fact that haloforms containing two fluorine atoms hydrolyze several orders of magnitude faster than those containing no fluorine. Hine (3, 8c) has attributed the greater reactivity of the haloforms containing fluorine to the added resonance stabilization of divalent-carbon intermediate.

11-7.c-7-1 IF-C=FI oe OG

The seleCtivity that an intermediate, such as one of the dihalomethylenes, shows toward a series of reagents of varying nucleophilicity is primarily a function of its stability; the more stable the intermediate, the more selective it is in its subsequent reactions. Large differences in nucleophilicity are often obscured when an unstable, highly reactive intermediate is employed because it will usually react with the first nucleophile that it encounters which is normally the solvent. Because of its greater stability, difluoromethylene should be the most selective of the dihalomethylenes in its subsequent reactions with reagents of varying nucleophilicity. This may account for the fact 28

4 that the results show that hydroxide, a reagent 10 times as reactive as water on the Swain-Scott nucleophilicity scale, is the principal capturing agent for difluoromethylene. Robinson's results might be explained by assuming that dichloromethylene is simply too unstable to show the desired selectivity and simply combines with water, which is in gross excess. This seems rather unlikely, however, in view of the work of Hine and Dowell (2) who showed that the basic hydrolysis of chloroform was slowed by the addition of the halides: chloride, bromide, and iodide.

This work was logically interpreted in terms of a mass-law effect, i.e., the halides capture the intermediate dichloromethylene converting it back to the trihalomethyl anion. The relative efficiencies with which the halide ions capture dichloromethylene were found to vary in the same manner as the relative rates at which they performed other typical nucleophilic displacement reactions, i.e., in the order of their corresponding Swain-Scott nucleophilicity constants.

The relative reactivities of the halides, water, and hydroxide on the Swain-Scott nucleophilicity scale are H 2 0(0.0), C1 - (3.04),

Br- (3.89), 0H - (4.20), and 1 - (5.04). Since Hine and Dowell were able to observe mass-law effects with chloride and bromide as well as with iodide, it seems unlikely that hydroxide, which is intermediate between bromide and iodide on the Swain--Scott nucleophilicity scale, would be completely unreactive as Robinson reports. These arguments suggest that the relative reactivities of hydroxide and water toward dichloromethylene ought to be reinvestigated.

The concentration term associated with k 2 was then investigated by evaluating k1/k [equation (24)]. Examination of the trends observed 2 29

in the calculated fractiowsof carbon monoxide (Table 1) shows that the amount of chlorodifluoromethane converted to this product increases as the reaction progresses. This suggests that one of the products formed in the reaction is responsible for converting difluoromethylene to carbon monoxide. Since it seemed unlikely that fluoride could lead to anything but fluoroform, the concentration of formate was associated with k 2 . The values of k /k (Table 1) were then calculated assuming that hydroxide 1 2 and formate were responsible for converting difluoromethylene to formate and carbon monoxide, respectively.

The values of k1/k2 show reasonably good precision and no trends toward rising or falling values are evident; however, the average value of k1/k2, i.e., 3.64 t 0.55, seems unusually high. Unfortunately, the relationship between k 2 and the concentration of formate was not dis- covered until after the study had been completed and the basic hydrolysis of chlorodifluoromethane in the presence of added formate was not studied.

Carbon Monoxide as a Reaction Product

Although the calculations in Table 1 point to the fact that carbon monoxide is indeed a product from the basic hydrolysis of chlorodifluoromethane and suggest that a relationship exists between k 2 and the concentration of formate, it could never be confirmed_that these correlations were not simply fortuitous, i.e., it could never be conclusively demonstrated that carbon monoxide was a reaction product. In the subsequent studies of the reaction of chlorodifluoromethane with phenoxide, direct evidence was sought for the presence of carbon monoxide because the calculated fractions showed that even during the first addition of the haloform, a large and definitely measurable fraction (ca., 0.3) of 30

chlorodifluoromethane was converted to carbon monoxide. To verify this,

the reaction flask was arranged so that the effluent gases would pass

into the gas-sampling loop of a gas chromatograph equipped with a column

(silica gel) capable of separating nitrogen, carbon monoxide, and

chlorodifluoromethane. The - haloform was then passed into the basic- phenolate mixture. After each addition, the atmosphere above the liquid phase was swept into the= sampling loop of the gas chromatograph with a slow stream of nitrogen and analyzed; however, no trace carbon monoxide could be detected. By repeating this procedure several times (after which the concentrations of formate and fluoride were comparable to the runs reported in Table 21, Appendix, we still could not detect carbon monoxide. These observations led to the conclusion that the calculated fractions of carbon monoxide (Table 1) were simply the result of the method of following the reaction and that, in fact, carbon monoxide was not a major reaction product. In all probability, it is never formed in sufficient quantities to warrant consideration in the stoichiometric relationships.

The reason why the calculated fractions of carbon monoxide were large in the phenoxide study and usually significant in the other studies is probably related to the indirect method of following the reaction, i.e., the inability of the experimental technique to distinguish between reactions of closely related stoichiometries. It will be noted that the formation of carbon monoxide and Formate differ only in the amount of hydroxide consumed. In addition, the difference is small, i.e., three equivalents to four, respectively. The changes in the concentrations of hydroxide, chloride, and fluoride are usually small and in many instances result in numbers of only two significant figures. Furthermore, 31

examination of the equations for calculating the respective fractions reveals that the concentration changes are often multiplied by integers ranging from two to four. As a result, small experimental errors are often magnified in the subsequent calculations. Due to the limitations of the experimental method for differentiating between reactions of closely related stoichiometries and the lack of conclusive evidence for the presence of carbon monoxide as a reaction product, it was elected to analyze the data neglecting carbon monoxide formation. However, it is evident that even if carbon monoxide formation is real, the amount produced is minimal at low concentration of formate. As a result, every effort was made in the subsequent investigations to study the reactions at low formate concentrations so that the fraction of carbon monoxide could be justifiably neglected, assuming that it is a reaction product.

Analysis of the data (Table 21, Appendix) from the basic hydrolysis of chlorodifluoromethane (assuming that only formate and fluoroform were formed) is shown in Table 2. The values of the relative rate constants, k1/k 3, were calculated from either of two independent equations derived from the stoichiometry relationships. The first was derived in terms of

A[OH - ] and A[C1 - ] and the second in terms of A[F - ] and [Cl]- as shown in equations (26) and (27).

k /k = (A[OH - - 1 3 ])CF- ]/(4A[C1 ] - ACOH - ])[OH- ] (26)

k /k = (ACF- 1 3 ] + ACC1 - DEF- 7/(2.6[C1- ] - ACF- DEOH- ] (27)

As the results show, the values of obtained k1/k3 from equations (26) and (27) were 7.7 ±0.1 and 10.1 ± 0.5, respectively. These values are in good agreement with that obtained when formation of carbon monoxide was 32

included, e.g., 9.4 ± 0.2. Hydroxide, therefore, is about an order of magnitude better than fluoride at capturing '`difluoromethylene.

Table 2. The Basic Hydrolysis of Chlordifluoromethane at 36°

Add'n. of [OH- ] [F- ] HCF C1 Avg. Avg. k1/k k,/k,e 2 --3b J_ 0 1 0.836 0.086 c c

2 0.608 0.200 7. . 7 9.6 3 0.444 0.283 7.6 10.6 4a 0.211 0.400 10.6d 9.1 d a HCF b 2Cl additions 4 and 5 combined for calculations. Calculated from c equation (26). Calculated fraction of HCF was negative. dOmitted 3 e from average ([0H-] < 0.4 M). Calculated from equation (27). avg

The Basic Hydrolysis of Chlorodifluoromethane in

the Presence of Monofunctional Reagents

In order to determine if bifunctional reagents, such as catechol monoanion, are more reactive toward difluoromethylene than monofunctional reagents, such as phenoxide, it was first necessary to establish the relative reactivities of the latter. For this reason, the basic hydrolysis of chlorodifluoromethane in the presence of phenoxide and several common nucleophilic reagents was studied.

The Reaction of Difluoromethylene with Fluoride

In order to firmly establish the relative reactivities of hydroxide and fluoride toward difluoromethylene, a study of the basic hydrolysis of 33

chlorodifluoromethane was conducted in the presence of added fluoride.

Table 3. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Fluoride at 36°

Add'n. [OH- ] [F- ] of b HCF C1 Avg. Avg. k,/k, k,/k 0 c 2 _1-3 ..1- -l)--.- 1 0.829 0.561 13.1 12.0 2a 0.600 0.669 8.5 8.1 3 0.432 0.746 5.7 7.3

9.1 i2.7 9.1 ±1.9 a b HCF2 Cl additions 2 and 3 combined for calculations. Calculated from equation (26). cCalculated from equation (27),

The results are shown in Table 3. The data (Table 22, Appendix) were analyzed assuming that only two products were formed: formate and fluoroform. The values of k1/k3 were calculated from equations (26) and (27).

As the results show, the values of k i/k3 decline as the reaction progresses. No explanation can be offered for this trend. However, despite the observed trend and the rather large average deviation, the results show clearly that hydroxide is about an order of magnitude more efficient at capturing difluoromethylene than is fluoride. This is in good agreement with the values obtained in the absence of added fluoride

(Tables 1 and 2). Because of its low reactivity toward difluoromethylene and the fact that in the presence of more reactive capturing agents the

3 4

necessary data could be obtained before significant quantities of fluoride were produced, the formation of fluoroform was neglected in subsequent calculations.

The Reaction of Difluoromethylene with Azide

The results of study of the basic hydrolysis of chlorodifluoromethane in the presence of azide are shown in Table 4. The data

(Table 23, Appendix) were analyzed assuming that only two products were formed: formate and difluoromethyl azide.

HCF C1 + N HCF N + Cl - 2 2 3

The values of were calculated from equations (28) and (29): k1/k5 k /k = (A[OH- - ]/(46.[C1- ] - A[OH - 1 5 ])[N ])[OH - ] (28) k /k =.(A[F - ] - A[F- - ] 1 5 - ])[N- ]/(2A[C1 ])[OH (29) where N is the nucleophilic reagent employed.

Table 4. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Azide at 36°

Add'n. of [OH- ] [N 3 - ] b k a/k k /k HCF C1 Avg. Avg. 1 5 1 5 2 1 1.111 0.459 0.343 0.292 2 1.004 0%424 0.345 0.384 3 0.893 0.403 0.340 0.419 4 0.794 0.373 0.314 0.319 5 0.686 0.337 0.403 0.377

Avg. 0.349 1 0.020 0.358 1 0.042

a b Calculated from equation (28). Calculated from equation (29).

As the results show, the average values of k l/k 5 calculated from both equations agree closely, e.g., 0.349 and 0.358. Azide, therefore, is about three times as efficient as hydroxide and about 27 times as efficient as fluoride at capturing difluoromethylene.

In cases where the concentration of the capturing agent (azide in this case) is followed directly, the fractions of the reaction product can be determined directly and compared with the calculated fractions for an additional check on the indirect method of following the reaction. , In this case, direct calculation of the fractions of difluoromethyl azide formed during each addition of chlorodifluoromethane can be obtained.from equation (30):

f = A[N 5 3- ]/A[C1 - ] (30) where f is the fraction of difluoromethyl azide. The indirectly deter- 5 mined fractions may be determined from equations (31) and (32).

f = (4A[C1 - ] - A[OH 5 - ])/4A[C1 - ] (31) f = (2A[C1- 5 ] - A[F- ])/2A[C1 - ] (32)

Comparison of the directly and indirectly determined fractions are shown in Table 5.

As the results show, good agieement between the two methods is observed. 36

Table 5. Comparison of Directly and Indirectly Determined Fractions of Difluoromethyl Azide

Add'n. of a b c HCF2 C1 f5 f5 f5 1 0.55 0.55 0.60 2 0.53 0.53 0.59 3 0.60 0.59 0.47 4 0.55 0.57 0.65 a b Calculated from equation (30). Calculated from equation (31). c Calculated from equation (32).

The data from the reaction of difluoromethylene with azide can also be analyzed with an integrated form of the rate equations. This additional method of checking the values of the relative rate constants can be employed whenever the concentrations of the capturing agent are followed directly and the concentration changes observed are large enough to permit graphical representation of the data. Derivation of the appropriate equations is shown below:

d[OH - ]/dt = 4k1 C0H- HCF2 ] (33) d[N 3- ]/dt = k 5 EN 3- HCF2 ] (34)

Dividing equation (33) by ( -34) gives equation (35).

d[OH- ]/d[N - ] = 4k COH- Yk [11 - ] 3 1 5 3 (35) 37

Separating the variables and integrating equation (35) gives equation (36):

log [0H - ] o /[0H- ] = 4k1 /k 5 log [143],/[N3] (36) where [OH - ] 0 and Ely () are the initial concentrations of hydroxide and azide, respectively.

As demanded by equation (36), a plot of log [0H - ] 0 /[0H- ] versus log [1\13] 0 /[N3] should give a straight line of slope 4k 1 /k 5 passing through the origin. Evaluation of k 1 /k 5 by the graphical method is shown in Figure 2.

As the results show, the graphical method gives a value of ki /k 5

of 0.348 which is in good agreement with the values, i.e., 0.349 and

0.358, obtained by the differential method.

Identification of Difluoromethyl Azide. Following the study of the reaction of difluoromethylene with azide, the basic reaction mixture was extracted with carbon tetrachloride. The combined extracts were dried and the infrared spectrum (Figure 10, Appendix) recorded. The absorptions due to the azide (20b) and C-F linkages (13, 20a) at 4.5 to

4.7 p and 8.0 to 8.2 p, respectively, leave no doubt that difluoromethyl azide was indeed the reaction product. The physical properties of difluoromethyl azide have not been reported; however, due to the toxicity and explosive character of low molecular weight azides, it was decided not to determine them.

The Reaction of Difluoromethylene with Cyanide

The results of a study of the basic hydrolysis of chlorodifluoromethane in the presence of cyanide are shown in Table 6. The data 38

0.3

0 . -)

x. ' 0

U 0

0.1

0.1 0.2

log By oliNV

Figure 2. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Azide at 36°

(Table 24, Appendix) were analyzed assuming that only two products were

formed: formate and difluoromethyl cyanide.

k HCF C1 + ON - 6 HCF CN + Cl- 2 2

The values of were calculated from equations (28) and (29). k1/k6

Table 6. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Cyanide at 36°

Add'n. of [OH- ] [CN - ] b HCF Cl Av•. A . a/k k, /kc 2 l 6 J.-U*- 1 0.654 0.464 7 .14 7.63

a b Calculated from equation (28). Calculated from equation (29)

As the results show, cyanide is only slightly better than fluoride

at capturing difluoromethylene. This is, of course, markedly out of line

with its expected reactivity. On the Swain-Scott nucleophilicity scale,

cyanide is three to four orders of magnitude more reactive than fluoride

in SN2-type displacements at carbon. Identification of Difluoromethyl Cyanide. Attempts were made to

isolate and obtain the infrared and nmr spectra of difluoromethyl

cyanide; unfortunately, these attempts proved futile, perhaps because of

the small amount of nitrile that was formed. 40

The Relative Reactivities of Azide, Cyanide, Hydroxide, and Fluoride

Summary

As shown in Table 7, the relative reactivities of azide, cyanide, hydroxide, and fluoride toward difluoromethylene are markedly different from their corresponding nucleophilicities in S N 2-type displacements at carbon (8, 23). It will be noted that not only is the order of reactivity different but the differences in reactivity are not even of the same order of magnitude. In contrast, Hine and Dowell (2), have shown that the halogens, chloride, bromide, and iodide, capture dichloromethylene in the same order in which they perform other displacements at carbon, i.e., in the same order as their respective Swain-Scott nucleophilicity constants. The differences in the order of reactivity

of , difluoromethylene and dichloromethylene with the various nucleophiles

(assuming that dichloromethylene consistently follows the Swain-Scott pattern) may be due to the instability of the resulting anion in the case of difluoromethylene.

N- + CF 2 N - CF

Alternatively, it may be related to the necessity for simultaneous protonation of the intermediate in accord with the principle of microscopic reversibility. Lengthy comparisons, however, cannot be made due to the uncertainties about the order of reactivity of dichloromethylene toward these same nucleophilic reagents.

Table 7. The Relative Reactivity of Azide, Cyanide, Hydroxide, and Fluoride toward Difluoromethylene

NuEleophile Difluorometh lene S 2 a

CN 1.2 1000 OH - 9.3 159 N - 26 100 2 F 1.0 1 a From the Swain-Scott nucleophilicity constants.

Although the reason for the change in the order of reactivity of the various nucleophiles toward difluoromethylene cannot be ascertained with any degree of certainty, it should be pointed out that many apparently simple nucleophilic displacements at carbon do nbt follow the order predicted from the Swain-Scott nucleophilicity constants

(8i, 8j, 24). The results, therefore, are not anomalous and, indeed, might even have been expected.

The Reaction of Difluoromethylene with Phenoxide

The results of a study of the basic hydrolysis of chlorodifluoromethane in the presence of phenoxide are shown in Table 8. The data

(Table 25, Appendix) were analyzed assuming that only two products were formed: formate and difluoromethyl phenyl ether.

HCF2 C1 + CF H + Cl - O- 2

The values of k1/k7 were calculated from equations (28) and (29).

Table 8. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide at 36°

Add'n. of [0H- ] [Ph0 - ] a HCF Avg. Avg. k /1 k /k c 2 C1 1 7 1 7 1 1.032 0.505 1.56 1.48 2 0.818 0.485 1.22 1.42 3 0.662 0.464 1.58 1.25 4 0.513 0.449 1.49 1.36 b b 5 0.373 0.434 1.65 1.81

Avg. 1.46 ±0.12 1.38 ±0.07

a Calculated from equation (28). b Omitted from average ([0H-] < 0.4 M). avg c Calculated from equation (29).

As the results show, the average values of kl /k 7 were found to be

1.46 and 1.38. Phenoxide, therefore, is only slightly less reactive than

hydroxide and about six times as reactive as fluoride toward

difluoromethylene.

As in the case of azide, the value of can also be graphically k1/k7 determined from equation (3 7 ) as shown in Figure 3.

log [0H - ] 0 /[0H- ] = 4k1/k 7 log [Ph0- ] 0 /[Ph0 - ] (37)

The value of ki /k 7 obtained in this manner, e.g., 1.48, is again in good

agreement with the values obtained by the differential method. 43

0.5

0. 14

-] H CO / 0.3 -]o [OH log

0.2

0.1

0 0.02 0.04 0.06 0.08 log [Ph0 - ] 0 /[Ph0- ]

Figure 3. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide at 36° 44

The Reaction of Difluoromethylene with 2,4-Dichlorophenoxide

The results of a study off the basic hydrolysis of chlorodifluoromethane in the presence of 2,4-dichlorophenoxide are shown in Table 9.

The data (Table 26, Appendix) were analyzed assuming that only two products were formed: formate and difluoromethyl 2,4-dichlorophenyl ether.

k 8 HCF C1 + Cl Cl OCF2H t Cl 2 Cl Cl The values of were calculated from equations (28) and (29). k1/k8

Table 9. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2,4-Dichlorophenoxide at 36°

Add i n. of [Ar0 [014- ] - ] b HCF C1 Avg. Avg. k/k a k /k 2 1-0 1-8 1 1.029 0.431 1.96 2.21 2 0.901 0.423 2.08 2.19 3 0.725 0.408 1.94 2.27 4 0.519 0.400 2.37 1.70

Avg. 2.09 10.14 2.09 1 0.20

a b Calculated from equation (28). Calculated from equation (29).

As the results show, the average values of ki/k 8 calculated from both equations were found to be identical, i.e., 2.09. Phenoxide, 45

therefore, is about 1.5 times as reactive as 2,4-dichlorophenoxide at capturing difluoromethylene. The basicity of these two phenoxides

differs by two to three orders of magnitude, 2,4-dichlorophenox 'ide being the weaker of the two. Although the order of reactivity is that expected, i.e., the weaker base being the weaker nucleophile, the difference in reactivity by no means reflects the difference in basicity.

The Reaction of Difluoromethylene with p-Methylthiophenoxide

The results of a study of the basic hydrolysis of chlorodifluoromethane in the presence of p-methylthiophenoxide are shown in Table 10.

The data (Table 27, Appendix) were analyzed assuming that only two products were formed: formate and difluoromethyl p-tolyl sulfide.

k 9 CH + HCF C1 CH CF H + Cl- 2 2

The values of kl/k9 were calculated from equations (28) and (29).

Table 10. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of .p-Methylthiophenoxide at 36°

Add'n. of [OH- ] [PhS - ] b HCF„Cl Avg. Avg. k /k a k,/k, 1 9 1 1.042 0.283 0.197 0.288 2 0.860 0.230 0.250 0.287 3 0.583 0.158 0.269 0.277

Avg. 0.238 ±0.026 0.284 ±0.005 a b Calculated from equation (28). Calculated from equation (29). 46

As the results show, the, average values of k l /k9 were found to be

0.235 and 0.284. p-Methylthiophenoxide, therefore, is about six times as reactive as phenoxide and about 37 times as reactive as fluoride toward difluoromethylene.

Identification of Difluoromethyl p-Tolyl Sulfide. Difluoromethyl p-tolyl sulfide was isolated from the reaction mixture by extraction with carbon tetrachloride. The infrared and nmr spectra are shown in

Figures 13 and 14, Appendix.

The Relative Reactivities of Some Nucleophilic

Reagents Toward Difluoromethylene

A list of the monofunctional nucleophilic reagents studied and their corresponding nucleophilicities toward difluoromethylene (relative to fluoride), kN/kF , are shown in Table 11.

Table 11. The Relative Reactivities of Some Nucleophilic Reagents Toward Difluoromethylene

Nucleophile k /k N F

Fluoride 1,0 Cyanide 1.2 2,4-Dichlorophenoxide 4.4 Phenoxide 6.0 Hydroxide 9.1 Azide p-Methylthiophenoxide 37

As the results show, the most reactive reagent studied, p-methylthiophenoxide, is only 35 to 40 times as reactive as the least reactive: fluoride. Much larger differences in reactivity would have been predicted from the Swain-Scott nucleophilicity constants. For example, cyanide is about 1000 times as reactive as fluoride in SN2-type displacements at carbon whereas it is less than twice as reactive as the latter toward difluoromethylene. Also surprising is the reactivity of the organic reagents toward difluoromethylene. As the results show, phenoxide is comparable to hydroxide and p-methylthiophenoxide is more reactive than azide. Unfortunately, these organic reagents are not included on the Swain-Scott nucleophilicity scale so direct comparison cannot be made; however, it is doubtful that phenoxides would be comparable to hydroxide in typical S N2-type displacements at carbon.

The Relative Reactivities of Water and

Hydroxide toward Difluoromethylene

As previously mentioned, it was recognized quite early that both water and hydroxide could attack difluoromethylene and convert it to formate.

v = k c0H—] + k.TH 0] I 2

However, our earlier results indicated that at strong base concentrations above 0.5 M, hydroxide was largely responsible for this transformation.

As the study progressed, it proved possible to evaluate the relative reactivities of hydroxide and water toward difluoromethylene.

The Reaction of Difluoromethylene with Phenoxide

By studying the basic hydrolysis of chlorodifluoromethane in the

it [ 48

presence of phenoxide under certain conditions, the relative reactivities of hydroxide and water toward the divalent-carbon intermediate can be ascertained. Considering that both water and hydroxide convert difluoromethylene to formate, the ratio of the fractions of formate and difluoromethyl phenyl ether can be calculated from equation (38).

=1[0H-] + d)/k [Ph0- ] (38) fl /f7 12 7

The relative reactivities of water and phenoxide toward difluoromethylene can be obtained by solving equation (38) for yk 7 , as shown in equation (39).

kc/k7 (Ph0 - ]/CH 01)"(fl/f7 - y0li - 7/k 7 M0- 1) (39)

To solve equation (39), the average concentration of phenoxide can be obtained from the titration data (Table 28, Appendix). The concentration of water was assumed to. be 55 M. The average concentration of hydroxide can be calculated from the average concentration of phenol and 10-10. the dissociation constant of phenol, i.e., 1 x The value of k /k 1 7 employed, i.e., 1.4, was that determined in our previous study of the basic hydrolysis of chlorodifluoromethane in the presence of phenoxide where the average concentration of hydroxide was always above 0.5 M.

The reaction of phenoxide with chlorodifluoromethane was studied under conditions where only equilibrium quantities of hydroxide were present, i.e., insufficient hydroxide was added to convert all of the phenol to phenoxide. Under these conditions, the fraction of difluoromethylene converted to formate from attack by water is at a maximum.

The data (Table 28, Appendix) were evaluated assuming that only two

4 9

products were formed: formate and phenyl difluoromethyl ether. The fractions of the two products were calculated from equations (40) and

(41), respectively.

f =(A[Ph0 - ] - A[C11V3A[C1 (40) 1 - ] f = (4.6[C1 - ] - A[Ph01)/3.6[C1 - ] 7 (41)

Calculation of f and f from equations (40) and (41) gave values 1 7 of 0.15 and 0.85, respectively. Substituting fl and f7 into equation (39) and solving for ki/k 7 gave a value of 1.3 x 10 -3 . Dividing the previously

/k gives a determined value of k 1/k 7, i.e., 1.4, by that obtained for k'1 7 value for k /k' of 1080. Hydroxide, therefore, is at least 1000 times as 11 reactive as water toward difluoromethylene.

The fraction of difluoromethylene converted to formate by attack of hydroxide, fi (OH - ), and water, f3'(H2 0), can be evaluated from equation (42).

f (OH - )/f - (H 0) = k [OH - ]/k"[H 11 2 11 2 0] (42)

At hydroxide concentrations of about 0.5 M, the value of fi (OH - )/fl (H20) calculated from equation (42) is about ten. At lower hydroxide concentrations, water begins to compete significantly. This explains why our earlier studies showed that at hydroxide concentrations above 0.5 M, hydroxide appeared to be the concentration term associated with k It 1. is also apparent why the relative rate constants tended to increase when the hydroxide concentrations fell below 0.5 M, i.e., because the total fraction of formate was employed in the calculations, not that due 50

only to capture of difluoromethylene by hydroxide. Because of the competition of hydroxide and water for difluoromethylene, we elected to discard any values of the relative rate constants which were determined at average hydroxide concentrations below 0.4 M.

The Reaction of Difluoromethylene with Phenoxide in Deuterium Oxide

The foregoing study of the hydrolysis of chlorodifluoromethane in the presence of phenoxide where only equilibrium quantities of hydroxide were present led to results which were useful in evaluating the relative reactivities of water and hydroxide toward difluoromethylene.

The reaction was slow, ca., 10 days, under these conditions. It seemed advisable, therefore, to ascertain whether the reaction was still proceeding by an a-elimination process since, in the absence of strong base, a simple nucleophilic displacement process might be the predominant route to the difluoromethyl phenyl ether. If this were the case, the values of

/k' would be in error. As a result, the reaction of chlorodifluoro- k11 methane with phenoxide was carried out under identical conditions except that deuterium oxide was employed as solvent and sodium deuteroxide as base. If the reaction is still proceeding by an a-elimination mechanism, then the difluoromethyl phenyl ether formed will contain carbon-bound deuterium.

Base + HCF C1 Base•H+ + CF + C1 - 2 2

Ph0- + CF + D 0 PhOCF D + OD- 2 2 2

On the other hand, if the reaction is proceeding by a nucleophilic displacement process, the product will not contain deuterium. 51

S 2 N . RE-10 - fiu,c1 ---4.PhOCF2H + Cl-

Analysis of the reaction product by nmr spectroscopy would, of course, reveal whether deuterium or protium was incorporated in the difluoromethyl group of the ether.

The reaction was carried out under conditions identical to those employed in our earlier study, i.e., insufficient sodium deuteroxide was added to convert all of the phenol to phenoxide. Analysis of the data

(Table 29, Appendix) with equations (40) and (41) showed that the fractions of formate and difluoromethyl phenyl ether were almost identical to those obtained in water, i.e., 0.18 and 0.82, respectively. Calcula- tion of k /k' for the reaction in heavy water showed that deuteroxide 11 is about 1000 times as reactive as deuterium oxide toward difluoromethylene.

Identification of Deuterodifluoromethyl Phenyl Ether. The nmr spectrum (Figure 9, Appendix) of the difluoromethyl phenyl ether obtained by extraction of the basic reaction mixture with carbon tetrachloride showed the expected absorption due to the aromatic protons but the distinct triplet of the difluoromethyl group (Figures 12 and 8, Appendix) was absent indicating that the product was deuterodtifluoromethyl phenyl ether and that the reaction was still proceeding by an a-elimination mechanism.

The Basic Hydrolysis of Chlorodifluoromethane

in the Presence of Bifunctional Reagents

After determining the relative reactivities of several monofunctional reagents (Table 11) toward difluoromethylene, a study of bifunctional reagents was undertaken. 52

The Reaction of Difluoromethylene with Catechol Monoanion

The basic hydrolysis of chlorodifluoromethane in the presence of

catechol was undertaken to ascertain if catechol monoanion, in accord

with the bifunctional catalysis concept, was superior to phenoxide at

capturing difluoromethylene.

OCF H k 2 HCF 2 C1 + 10 + H+ + Cl

OH 0- The relative rate constants, k 1/k10, were calculated from equations (43)

and (44):

k /k = (a0H - ] - A[Cl (43) 1 10 - ])[M]/(4A[C1 - ] - ACOH - ])'![OH - ] k /k10 = (A[F- ])[M]/(2A[C1 - ] - A[F - ])[OH - ] (4 4)

where M is catechol monoanion. Equations (43) and (44) were derived

assuming that only two products were formed: formate and catechol mono-

(difluoromethyl) ether. Capture of difluoromethylene by catechol dianion

and the anion of catechol mono(difluoromethyl) ether was neglected.

In order to calculate the relative rate constants, the concentra-

tions of the capturing agents, hydroxide and , catechol monoanion, must be

ascertained. Unfortunately, neither can be determined directly.

Potentiometric titration of a basic solution containing catechol with

standard acetic acid shows two breaks in the titration curve. The first

includes catechol dianion and hydroxide and the second, catechol mono-

anion. If the reaction mixture is analyzed after addition of chlorodi-

fluoromethane, the anion of catechol mono(difluoromethyl) ether is also present. This moiety behaves like catechol monoanion and is included in

the second break of the titration curve. 53

Although the concentrations of the reactive catechol species cannot be determined directly, they can be ascertained from the stoichiometry relationships. Equation (45) relates the initial catechol concentration to the various catechol species in solution:

[Catechol] o = [D] + [M] + [E] (45) where D and E are catechol dianion and the anion of catechol mono(difluoromethyl) ether, respectively. The concentration of catechol mono(difluoromethyl) ether can be calculated from either equation (46) or equation (47).

[E] = (4A[C1 - ] - A[Strong Base])/3 (46)

[E] = (2A[C1 - ] - A[F- ])/8 (47)

Letting,

C = [Catechol], - [E] and rearranging equation (45), equation (48) is obtained.

C = - [M] + [D] (48)

The stoichiometry of the strong base portion of the potentiometric titration curve is given by equation (49):

b = [D] + [OH - ] (49) where b is strong base.

The equilibrium expression relating the various reactive species in solution is given by equation (50): 54

K2 = [0][1.1+]/[M] = [D]Kw/[M][0H - ] (50)

where K2 is the second dissociation constant, of catechol. The value of K2 has been reported by several workers (26-28). Jameson and Neillie

(27) and Timberlake (28) agree that the value of K 2 is 1.5 x 10-13 .

This value was employed in our calculations.

Solving the three simultaneous equations, (48), (49), and (50), for hydroxide yields equation (51):

- + 4bK Cotf_ ] _ (K + C - b) t i(K + C - b) 2 (51) 2 where K is Kw/K2 . The concentrations employed in these calculations are the average values during which the respective fractions were determined.

After solving for the average hydroxide concentration [equation (51)], the average concentrations of catechol di- and monoanions can be ascertained from equations (49) and (48), respectively.

The reaction of chlorodifluoromethane with catechol was studied at strong to weak base ratios of (2.8 to 1.2)/1 and (0.88 to 0.42)/1.

Under the first set of conditions, the principal organic moiety in solution is catechol dianion; under the second set of conditions, approximately equal concentrations of both catechol mono- and dianions were present. The results of the study at strong to weak base ratios of (2.8 to 1.2)/1 (Table 30, Appendix) are shown in Tables 12 and 13.

As the results show, catechol monoanion appears to be about an order of magnitude more efficient at, capturing difluoromethylene than is phenoxide or about as efficient as azide. Although the trend is in the direction supporting the bifunctional capturing concept, the magnitude 55

of the change is rather small.

The results of the study at strong to weak base ratios of

(0.88 to 0.42)/1 are shown in Tables 14 and 15. (Table 31, Appendix).

As the results show s the values of k1 P510 are about an order of

magnitude larger than those determined (Tables 12 and 13) at the higher

strong to weak base ratios. Under, these conditions, catechol monoanion

appears to be comparable to 2,4-dichlorophenoxide at capturing difluoromethylene. This study, however, is complicated by the fact that the hydroxide concentration is so low that water now competes successfully with hydroxide for difluoromethylene. The calculated fractions of formate, therefore, are not all due to attack by hydroxide as demanded by equations (43) and (44). Substituting the fraction of formate due to

attack by hydroxide, as calculated from equation (42), however, does not

give values of which are identical to those obtained in Tables 12 k1/k10 and 13. This seems to indicate that catechol dianion is also capturing

difluoromethylene.

i /k Despite the uncertainties as to the exact value of k ' 10' it is evident that catechol monoanion is no more than ten times as reactive as phenoxide toward difluoromethylene and in all probability, it is even less, since catechol dianion undoubtedly captures to some extent.

Assuming that catechol monoanion is, an order of magnitude more reactive than phenoxide, it is questionable whether this small difference in reactivity can be attributed to bifunctional catalysis, particularly in view of the experimental uncertainties. In fact, this small difference

in reactivity is rather insignificant in the light of the effects of other, bifunctional reagents that have been studied. For example, Swain

Table 12. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36°

Add'n. of [0H- ] [D] [M] HCF C1 Avg. Avg. Avg. a 2 k1/k10 1 0.804 0.451 0.037 0.304 2 0.581 0.423 0.048 0.183 3 0.408 0.384 0.063 0.255 4 0.300 0.351 0.077 0.415

Avg. 0.289 ±0.93

a Calculated from equation (43).

Table 13. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36°

Add'n. of [OH -] [D] [m] HCF2 C1 Avg. Avg. Avg. k /k a 1 10 1 0.810 0.445 0.036 0.249 2 0.602 0.402 0.044 0.148 3 0.448 0.344 0.051 0.143 4 0.348 0.303 0.058 0.230

Avg. 0.193 ±0.47

a Calculated from equation (44).

Table 14. The BaSid Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36°

Add'n. of [OH - ] [D] [M] HCF C1 Avg. Avg. a 2 Avg. k1 /k10 1 0.146 0.670 0.304 1.95 2 0.109 0.587 0.357 2.11 3 0.078 0.470 0.42 3.30 4 0.053 0.382 0.482 3.00

Avg. 2.59 1 0.56

a Calculated from equation (43).

Table 15. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol at 36°

Add'n. of [OH- ] [D] [m] HCF Cl Avg. Avg. Avg. k /k a 2 1 10 1 0.148 0.668 0.300 5.52 2 0.111 0.585 0.353 4.00 3 0.075 0.473 0.427 3.20 4 0.053 0.382 0.484 4.32

Avg. 4.32 ±0.72

a Calculated from equation (44). 58

and Brown (8g, 14) showed that at a concentration of 0.001 M, 2-hydroxypyridine was 7000 times as effective in promoting the mutarotation of tetramethylglucose (in benzene) as a mixture of 0.001 M phenol and 0.001 M pyridine. The small differences in reactivity observed in our studies

(if indeed they exist) are perhaps best interpreted in terms of polar effects.

Alternatively, the bifunctional capturing mechanism may be operating but is simply obscured by the ease of intermolecular proton transfer facilitated by the large concentration of water. In support of this concept, Swain and Brown found that 2-hydroxypyridine was not nearly as effective at promoting the mutarotation of glucose in water as it was at promoting the reaction of tetramethylglucose in benzene.

Direct Determination of Catechol Mono(Difluoromethyl) Ether. As previously mentioned, in a study of this type, where the amounts of the various products are determined in an indirect manner, the question naturally arises as to whether the calculated values are correct, i.e., do they correspond to the amounts of the respective products as determined, by direct analysis. In several instances, we have described ways of checking the, calculated fractions or amounts of products with those determined directly. These methods have been contingent upon following the disappearance of the capturing agent directly. In the catechol study, it proved impossible to follow'the disappearance of catechol directly; consequently, an analytical technique was developed for following the appearance of the reaction product, i.e., catechol mono(difluoromethyl) ether.

Catechol mono(difluoromethyl) ether was prepared by the basic hydrolysis of chlorodifluoromethane in the presence of catechol

-If 59

(Table 32, Appendix). An aliquot of the reaction mixture (300 ml.) was transferred to a beaker and adjusted to a pH of 4.5 with dilute hydro- chloric acid. The acidic mixture was then transferred to a separatory funnel and extracted a total of twelve times with 15 to 20 ml. portions of carbon tetrachloride. The first six extracts were combined, as were the last six, and each combination analyzed separately for catechol mono(difluoromethyl) ether by an infrared method. The results are shown in Table 16.

Table 16. Direct Determination of Catechol Mono( difluoromethyl) Ether

Combined C7H60 2 F2 Vol. of Extract C 7H202F2 Extract Absorbance M ml. moles

1-6 0.365 0.235 113 0.0266 7-12 0.070 0.045 117 0.0053

Total 0.0319

The amount of catechol mono(difluoromethyl) ether in the 300 ml. aliquot was calculated with equations (46) and (47) from the data

(Table 32, Appendix). The average of the two calculated values was 0.0343 moles., As shown in Table 16, 0.0319 moles were isolated by extraction which corresponds to a 93 percent recovery. Further extraction would no doubt have improved this figure somewhat; however, it was evident that the calculated fraction (0.660) was indeed quite close to the directly determined value (0.620). 60

It was recognized that any catechol bis(difluoromethyl) ether that might have been formed in the course of the reaction would also absorb at the wavelength employed for analysis of catechol mono(difluoromethyl) ether; however, since only about 15 percent of the available catechol had been reacted and since it seeme&unlikely. that the anion of catechol mono(difluoromethyl) ether should be markedly more reactive than catechol monoanion, we neglected this contribution.

The Reaction of Difluoromethylene with Pyrogallol

The basic hydrolysis of chlorodifluoromethane in the presence of pyrogallol was also studied in a search for a bifunctional capturing agent.

OH OH

HCF2 C1 + + H+ + Cl -

OCF H 2

The data (Table 33, Appendix) were analyzed' assuming that only two products were formed: formate and pyrogallol mono(difluoromethyl) ether.

The fractions of the reaction products were calculated from equations

(52) and (53).

f = (A[Strong Base] - A[C1 - ])/3A[C1 - ] 1 (5 2) f = (40[C1 -] - A[Strong Base])/3A[C1 - ] (53) 11

An attempt was made to analyze the data from the pyrogallol study in a manner similar to that employed in the catechol study; however, the second dissociation constant, K for pyrogallol is an order of 2' 61

magnitude greater than that for catechol (26), i.e., 2.3 x 10 -12 . Under the conditiOns of the experiment,theconcentration of hydroxide was less -2 than 10 M. Water, therefore, was the principal capturing agent leading to the conversion of difluoromethylene to formate.

Because of the low hydroxide concentration, the only way to compare the relative reactivities of the anion of pyrogallol with 'phenoxide was to employ the phenoxide data (Table 28, Appendix) where insufficient hydroxide was added to convert all of the phenol to phenoxide. In this study, water is also the principal reagent leading to conversion of difluoromethylene to formate. It will also be noted that both studies were carried out at phenol and pyrogallol concentrations of 0.5 M; the concentrations of the two organic species, therefore, may be assumed to be equal. With these assumptions, the ratio of the fractions of the difluoromethyl aryl ethers providesa rough measure of the relative reactivities of the two organic species toward difluoromethylene.

The calculated fraction of difluoromethyl phenyl ether (Table 28,

Appendix) was 0.85. The calculated fraction of mono(difluoromethyl) pyrogallol (Table 33, Appendix) was 0.78. Other things being equal, phenoxide is actually a better capturing agent than is the anion of pyrogallol. These results support those obtained in the catechol study in that the bifunctional reagent did not show an accelerated rate of capture. Bifunctional catalysis, therefore, either is not occurring or is not observable under the reaction conditions.

The , Reaction of Difluoromethylene with 2-Mercaptoethanol

In order to eliminate part of the uncertainty from the study of the reactions of difluoromethylene with catechol and pyrogallol 62

monoanions and to extend our study of bifunctional capturing agents, the basic hydrolysis of chlorodifluoromethane in the presence of the anion of

2-mercaptoethanol was investigated.

HO - + HCF C1 H 2O + CF + Cl- 2 2

...... ,-- OH o-•, ,.... OH CH k H2 O CH 1 2 12 I 2 1 2 + CF ---- H 2 2H + OH - CH CH e'! F CH CF 2 2 N . • 2 2N. / 2 S- S .

Unfortunately, even in this system it is necessary to make assumptions about certain dissociation constants in order to solve for the relative rate constants; however, for the most part, the assumptions are much more tenable. For example, if it is assumed that the second-dissociation constant of 2-mercaptoethanol is several orders of magnitude lower than water, then when the teaction is carried Out in the presence, of excess of hydroxide the amount of 2-mercaptoethanol in the, form of the dianion may be neglected.

K d

- S-CH -CH -OH 4-- - S-CH -CH -C+ HI" 2 2 2 2

As a result, the first and second breaksin the titration curve may be regarded as hydroxide and the anion of 2-mercaptoethanol, e.g.,

- S-CH -CH -OH. In this case, the assumption that K > K is probably 2 2 w d quite good. The dissociation constants for alcohols, are about two pK units higher than water and the inductive effect of the full negative charge already on the anion of 2--mercaptoethanol will certainly impede further ionization. 63

Unfortunately, those changes in the structure of the organic capturing agent which lead to less ambiguity about the nature of the reactive species in solution, also tend to make it (at least from a theoretical viewpoint) a less likely bifunctional capturing agent. For example, the anion of 2-mercaptoethanol would be expected to be a poorer bifunctional capturing agent than either catechol or pyrogallol monoanions because the conformation of the reacting centers is not fixed as it is with catechol monoanion due to free rotation about the carbon-carbon bond. In addition, the hydroxylic proton is alcoholic in nature rather than phenolic and not nearly as acidic.

The results of a study of the basic hydrolysis of chlorodifluoromethane in the presence of the anion of 2-mercaptoethanol (Table 34,

Appendix) are shown in Table 17.

Table 17. The Basic Hydrolysis of Chlorodiflupromethane in the Presence of 2-Mercaptoethanol at 36°

Add'n. of [OH - ] [I6- ] b HCF C1 Avg. Avg. k /k a k /k 2 1 12 1 12 1 1.224 0.494 0.258 0.390 2 1.053 0.447 0.312 0.691 3 0.796 0.378 0.326 0.846 4 0.567 0.320 0.456 0.550

Avg. 0.338 ±0.059 0.619 ±0.149 a b Calculated from equation (28). Calculated from equation (29). 64

The data (Table 34, Appendix) were analyzed assuming that only two products were formed: formate and 2-hydroxyethyl difluoromethyl sulfide.

The values of were calculated from equations (28) and (29). k1/k12 As the results show, the relative rate constants exhibit a tendency, toward rising values as the reaction progresses and the precision is not as good as in several of the other cases that have been studied. However, the order of magnitude of the relative rate constants is not in question and shows that the bifunctional reagent is not quite as efficient as p-methylthiophenoxide at capturing difluoromethylene. As a result, there is no measurable rate-accelerating effect that could be attributed to bifunctional capture.

The Basic Hydrolysis of Chlorodifluoromethane in

60 Percent Ethanol

An investigation of the basic hydrolysis of chlorodifluoromethylene in 60 percent ethanol was undertaken to determine how solvent changes affected the course of the reaction.

The Reaction of Difluoromethylene with Hydroxide in 60 Percent Ethanol

The data from the basic hydrolysis of chlorodifluoromethylene methane in 60 percent ethanol are shown in Table 35, Appendix. Plots of hydroxide versus chloride and fluoride versus chloride are shown in

Figures 4 and 5, respectively. The stoichiometry of the reaction in this solvent is given by the slopes of these plots which may be summarized by equation (54).

k + HCF C1 + H 2 O + EtOH 2.51 H + 1.02 F- + C1 - 2 OH- (54) 65

1.0

0.8 M ide, drox

Hy 0.6

0.4

0.1 0. :2 0.3 0.4

Chloride, M

Figure.4. The Basic Hydrolysis of Chlorodifluoromethane in 60 Percent Ethanol at 25° 66

0.4

0.3 M ide,

Fluor 0.2

0.1

0 0. 1 0.2 0.3 0.4

Chloride, M

Figure 5. The Basic Hydrolysis of Chlorodifluoromethane in 60 Percent Ethanol at 25° 67

This stoichiometry is not characteristic of any single reaction but undoubtedly represents a combination of several reactions.

Determination of the fractions of the reaction products by the, indirect method that we have employed throughout this study is dependent upon exact knowledge of the stoichiometries of the principal reactions.

In the, basic hydrolysis of chlorodifluoromethane (in water), earlier workers had already defined the reaction products, i.e., formate, carbon monoxide, and fluoroform. In the presence of ethanol, it is quite probable that triethyl orthoformate and difluoromethyl ethyl ether are also formed; however, without a detailed analysis of the reaction products, application of the indirect method of determining the product distribution would be meaningless.

Fortunately, it is not necessary to have ,a detailed knowledge of the reactions of the solvent with difluoromethylene in order to compare the relative reactivities of other nucleophilic reagents with that of the solvent. For this purpose, it is only necessary to know the stoichiometry of the solvent reaction as defined in equation (54). For example, if one wishes to compare the relative reactivities of phenoxide with hydroxide in 60 percent ethanol employing the indirect approach, the algebraic equation for calculating the relative rate constants may be derived from the balanced chemical equations representing the reactions of chlorodifluoromethane with the solvent and phenoxide as shown in. equations (55) and (56).

k/k 7 = ( A[OH- ])[Ph0 - ]/(2.51 A[C1 - ] - ACCIFI - MOH- ] (55) k /k ] - 1.02 A[F- (56) s 7 = ( A[F- ])[Ph0 - ]/(A[C1 - ])[°H- ] 68

For the sake of simplicity, it has been assumed that the concentration term associated with k is hydroxide, i.e., strong base, although it is s recognized that part of the base is in the form of ethoxide. This approach of analyzing the data is also predicted on the assumption that the products are determined by the nature of the species attacking difluoromethylene and that addition of another nucleophilic reagent does change the relative reactivities of the solvent species toward the divalent-carbon intermediate.

The Reaction of Difluoromethylene with Phenokide

The results of a study of the basic hydrolysis of chlorodifluoromethane in 60 percent ethanol in the presence of phenoxide (Table 36,

Appendix) are shown in Table 18. The values of were calculated ks/k7 from equation (55).

k 7 + HCF7 C1 —4- )-0CF H + Cl L. 2

As the results show, hydroxide, i.e., strong base, is three to five times as effective as phenoxide at capturing difluoromethylene. In water, k i/k 7 was about 1.4 which indicates that phenoxide is about twice as efficient at capturing the divalent-carbon intermediate in aqueous solution as it is in 60 percent ethanol.

The Reaction of Difluoromethylene with m-Chlorophenoxide

The results of a study of the basic hydrolysis of chlorodifluoromethane in 60 percent ethanol in the presence of m-chlorophenoxide

(Table 37, Appendix) are shown in Table 19. The values of were ks/k13 69

Table 18. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide in 60 Percent Ethanol at 25 °

Add'n. of [OH- ] [Ph0 - ] HCF C1 Avg. k/k a 2 Avg. o 7 1 0.875 0.495 3.07 2 0.698 0.473 2.67 3 0.550 0.456 3,17 4 0.440 0.438 4.52 5 0.259 0.415 3.43

Avg. 3.37 ±0.28 a Calculated from equation (55).

Table 19. The Basic Hydrolysis of Chlorodifluoromethane in the Presence.of m-Chlorophenoxide in 60 Percent Ethanol at 25°

Add'n. of [OH- ] [Ph0- ] HCF C1 Avg. Avg. k/k a 2 s 7 1 0.981 0.477 2.82 2 0.829 0.460 2.59 3 0.672 0.443 2.56 4b 0.521 0.426 2.34 5 0.386 0.409 2.64

Avg. 2.59 ±0.11 a . Calculated from equation (55). b AdditIons. 4 and 5 mb inede'for calculation.

calculated from equation (55).

Cl Cl

As the results show, m-chlorophenoxide is more efficient at capturing difluoromethylene than is phenoxide although the latter is more basic (and assumedly more nucleophilic) by several orders of magnitude.

Identification of Difluoromethyl m-Chlorophenyl Ether. Difluoro- methyl m- chlorophenyl ether was isolated from the reaction mixture by extraction with carbon tetrachloride. The infrared and nmr spectra are shown in Figures 15 and 16, Appendix.

The Reaction of Difluoromethylene with p-Chlorophenoxide

The results of a study of the basic hydrolysis of chlorodifluoromethane in 60 percent ethanol in the presence of p-chlorophenoxide

(Table 38, Appendix) are shown in Table 20. The values of were ks/k14 calculated from equation (55).

k14 + HCF2 OCF Cl Cl Cl 2H + Cl 71

Table 20. The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Chlorophenoxide in 60 Percent Ethanol at 25° ,

Addition of [0H- ] [Ph0 - ] HCF C1 Avg. Avg. k/k 2 c 14 l a 0.942 0.517 3.39 2 0.76() 0.493 2.54 3 0.576 0.465 3.71 4 0.414 0.449 2.715

Avg. 3.10 ±0.45 a b Additions land 2 combined-for calculation. Celculated from equation (55).

As the results show, p-chlorophenoxide is more effective at capturing difluoromethylene than is phenoxide although the latter is more basic by several orders of magnitude.

Identification of Difluoromethyl p-Chlorophenyl Ether. Difluoro- methyl p-chlorophenyl ether was isolated from the reaction mixture by extraction with carbon tetrachloride. The infrared spectrum of this

compound is shown , in Figure 17, Appendix. 72

CHAPTER IV

CONCLUSIONS

The primary objective of this study was to ascertain if bifunctional

reagents were more reactive toward difluoromethylene than monofunctional

reagents. The results have shown that they are not. In conclusion,

therefore, the bifunctional catalysis theory, as previously proposed, is

either not operating or not observable under the reaction conditions.

The secondary objective, i.e., to establish the relative reac-

tivities of several nucleophilic reagents toward difluoromethylene, has

also been completed. The results have shown that common nucleophilic

reagents, such as azide, cyanide hydroxide, etc., do not follow the

Swain-Scott nucleophilicity pattern and that there is little difference

in the reactivities of a variety of nucleophilic reagents (Table 11)

toward difluoromethylene. In conclusion, therefore, it seems evident

that difluoromethylene is not selective enough (due to its high reactivity)

to show marked differences in reactivity.

The effect of changing the solvent from water to 60 percent

ethanol has been investigated. The results show that very little change

in the reactivity of difluoromethylene was effected by changing the

solvent. In conclusion, therefore, the reaction is not markedly

solvent dependent.

The relative reactivities cf water and hydroxide toward difluoro-

methylene have been established. The results show that hydroxide is

11 73

about 1000 times as reactive as water toward the divalent-carbon intermediate. This is the largest difference in reactivity that has been observed. 74

AFIENDI X Table 21

The Basic Hydrolysis of Chlorodifluoromethane at 36°

Strong Base Chloride Fluoride a d Add'n. of AcOH OH- AgNO 3 c Cl- Sample Ce(NO ) Titration EDTAe F- 3 3 b HCF C1 ml. 2 b ml. M ml. ml. ml. ml. Initial 9.94 1.006

1 6.62 0.670 2.72 0.084 10.0 10.0 25.0 1.70 0.172 6.54 0.662 2.72 0.084 10.0 10.0 25.0 1.67 0.173

2 5.44 0.551 3.68 0.114 5.0 10.0 25.0 3.65 0.229 5.0 10.0 25.0 3.62 0.230

3 3.33 0.337 5.57 0.172 5.0 10.0 25.0 1.84 0.336 5.57 0.172 5.0 10.0 25.0 1.85 0.335

4 1.95 0.197 6.84 0.212 5.0 10.0 25.0 0.87 0.394 5.0 10.0 25.0 0.88 0.393

5 0.85 0.086 7.95 0.246 2.0 10.0 25.0 4.37 0.464 2.0 10.0 25.0 4.38 0.463 a b c d e 1.012 M. Per 10 ml. sample. 0.3096 M. 0.07423 M. 0.04949 M.

Table 22

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Fluoride at 36°

Strong Base Chloride Fluoride a d e Add'n. of AcOH OH AgNO 3 c Cl Sample Ce(NO ) Titration EDTA 3 F - b 3 HCF C1 2 ml. M ml.b M ml. ml. ml. ml. M Initial 9.30 0.941 2.0 10.0 25.0 4.07 0.509 9.21 0.932 2.0 10.0 25.0 4.05 0.512

1 7.14 0.723 1.81 0.056 2.0 10.0 25.0 3.37 0.613 - 2.0 10.0 25.0 3.37 0.613

2 5.75 0.582 3.07 0.095 2.0 10.0 25.0 3.02 0.665 2.0 10.0 25.0 3.01 0.666

3 4.72 0.478 4.03 0.125 2.0 10.0. 25.0 2.61 0.726 2.0 10.0 25.0 2.61 0.726

4 3.81 0.386 4.93 0.153 2.0 10.0 25.0 2.34 0.766 2.0 10.0 25.0 2.34 0.766

a b c d e 1.012 M. Per 10 ml. sample. 0.3096 M. 0.07423 M. 0.04949 M.

Table 23

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Azide at 36°

Strong Base Chloride Azide Fluoride Add'n. a e AcOH OH AgNO3 c Cl AgNO 3 c N3 - + Cl - N Sample Ce(NO 3 ) 3 d Titration EDTA F of 3 b HCF Cl ml. M ml.b .M ml.b M ml. ml. ml. ml. M 2 M Initial 11.46 1.160 15.85 0.491 0.491 11.47 1.161

1 10.50 1.063 1.71 0.0 5 3 15.55 0.481 0.428 5.0 10.0 25.0 6.75 0.044 L 5.0 10.0 25.0 6.75 0.044 2 9.35 0.946 3.80 0.118 17.38 0.538 0.420 5.0 10.0 25.0 5.73 0.105 9.35 0.946 5.0 10.0 25.0 5.70 0.107

3 8.31 0.841 5.63 0.174 18.12 0.561 0.387 5.0 10.0 25.0 4.81 0.160 5.0 10.0 25.0 4.83 0.159

4 7.39 0.748 7.50 0.232 19.10 0.592 0.359 5.0 10.0 25.0 4.02 0.207

5 6.17 0.624 9.71 0.301 19.91 0.616 0.315 5.0 10.0 25.0 3.02 0.266 5.0 10.0 25.0 2.98 0.268

a b c d e 1.012 M. Per 10 ml. sample. 0.3096 M. 0.0742 M. 0.04949 M.

Table 24

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Cyanide at 36°

Strong Base Chloride Cyanide Fluoride Add'n. of Ac0Ha OH- AgNO 3 c Cl- AgNO 8 c CN - Sample Ce(NO 3 ) 3e Titration EDTAf F- b HCF Cl ml. d 2 M ml,b M ml.M ml. ml. . ml. ml. M Initial 9.82 0.994 7.60 0.470 9.86 0.998 7.56 0.468

1 3.08 0.312 6.06 0.188 7.40 0.458 5.0 10.0 25.0 1.69 0.345 6.12 0.189 7.40 0.458 5.0 10.0 25.0 1.71 0.344

a b c d e f 1.012 M. Per 10 ml. sample. 0.3096 M. Per 5 ml. sample. 0.07423 M. 0.04949 M. Table 25

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide at 36°

Strong Base Weak Base Chloride iluorlde Add f n. of Ac0Ha OH Ac0Ha PhO AgNO Cl Sample Ce(NO 3 )1 Titration EDTAe F b b b HCF2 C1 ml. M ml. M ml. M ml. ml. ml. ml. M

Initial 10.85 1.165 4.80 0.515 10.84 1.164 4.77 0.512

1 8.38 0.900 4.64 0.498 3.56 0.087 5.0 10.() 25.0 5.29 0.131 5.0 ln.n 25.0 5.30 0.131

2 6.85 0.736 4.40 0.473 6.04 0.148 5.0 10.0 25.0 3.87 0.216 5.0 10.0 25.0 3.83 0.218

3 5.48 0.589 4.25 0.456 8.21 0.201 5.0 10.0 25.0 2.73 0.283 5.0 10.0 25.0 2.68 0.286

4 4.08 0.438 4.12 0.442 10.65 0.261 5.0 10.0 25.0 1.47 0.358 5.0 10.0 25.0 1.47 0.358

5 2.88 0.309 3.98 0.427 12.86 0.316 4.0 10.0 25.0 1.78 0.425 4.0 10.0 25.0 1.78 0.425 a c d e 1.074 M. bPer 10 ml. sample. 0.2455 M. 0.07423 M. 0.04949 M. Table 26

The Basic Hydrolysis of Chlorodifluoromethane in the . Presence of 2,4-Dichlorophenoxide at 36°

Strong Base Weak Base. Chloride Fluoride e Add'n. of Ac0H a OH Ac0Ha PhO AgNO Cl Sample Ce(NO 3 )3 Titration EDTA F b b ml b, HCF C1 ml. M ml. 11114. M ml. ml. ml. ml. M 2 M Initial 10.26 1.102 4.04 0.434 10.26 1.102 4.04 0.434

1 8.91 0.957 3.99 0.429 1.80 0.044 5.0 10.0 25.0 6.25 0.074 5.0 10.0 25.0 6.25 0.074

2 7.88 0.846 3.88 0.417 3.18 0.078 5.0 10.0 25.0 5.33 0.129 5.0 10.0 25.0 5.30 0.130

3 5.63 0.604 3.72 0.399 6.34 0.156 5.0 10.0 25.0 3.20 0.255

If 4.05 0.435 3.73 0.401 8.63 0.212 5.0 10.0 25.0 1.90 0.332 5.0 10.0 25.0 1.90 0.332 a b c d e 1.074 M. Per 10 ml. sample. 0.2455 M. 0.07423 M. 0.04949 M. Table 27

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Methylthiophenoxidea at 36°

Strong Base Chloride Fluoride b d Add'n. of AcOH OH AgNO Cl- Sample Ce(NO ) a Titration EDTAf F- 3 3 3 HCF C1 ml. c M ml.c ml. ml. ml. 2 M ml. Initial 10.21 1.096

1 9.20 0.988 2.61 0.064 4.0 10.0 25.0 6.61 0.066 2.62 0.064 4.0 10.0 25.0 6.62 0.065

2 6.82 0.732 8.00 0.196 4.0 10.0 25.0 4.76 0.203 8.00 0.196 4.0 10.0 25.0 4.77 0.203

3 4.03 0.433 14.04 0.345 4.0 10.0 25.0 2.72 0.355 14.15 0.347 4.0 10.0 25.0 2.71 0.356

b c d e aInitial concentration 0.301 M. 1.074 M. Per 10 ml. sample. 0.2455 M. 0.07423 M. f0.04949 M.

Table 28

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide a at 36°

Weak Base Chloride b d Add'n. of AcOH Ph0 AgNO Cl 3 HCF Cl ml. M ml. c 2 Initial 4.72- 0.478 4.71 0.477

1 3.51 0.355 2.78 0.086 3.49 0.353 2.76 0.085

a c d Phenol, 0.50 M. b1.012 M. Per 10 ml. sample. 0.3096 M.

Table 29

The Basic Hydrolysis of Chlorodifluoromethane in Deuterium Oxide a in the Presence of Phenoxide at 36°

Weak Base Chloride b d _ Add'n. of AcOH Ph0 AgNO Cl 3 c HCF2 C1 ml. M ml. c

Initial 4.66 0.472 4.66 0.472

1 3.85 0.391 1.70 0.053 1.72 0.053

a c d Phenol, 0.50 M. b1.012 M. Per 10 ml. sample. 0.3096 M. Table 30

The Basic Hydrolysis of Chlorodifluoromethane in _the Presence of Catechol Dianion at 36°

Strong Base Weak Base Chloride Fluoride a e Add'n. of AcOH OH Ac0Ha Ph0 AgNO3 c Cl Sample Ce(NO ) Titration EDTA F 3 3 HCF2 C1 ml. ml. b ml. ml. ml. ml.

Initial 13.76 1.392 4.86 0.492 13.76 1.392 4.89 0.495

1 11.06 1.119 2.46 0.076 5.0 10.0 25.0 5.31 0.130 11.04 1.117 2.46 0.076 5.0 10.0 25.0 5.34 0.128

2 8.81 0.891 4.84 0.150 5.0 10.0 25.0 3.66 0.228 4.84 0.150 5.0 10.0 25.0 3.63 0.230

3 6.85 0.693 7.10 0.220 5.0 10.0 25.0 2.20 0.314 7.05 0.218 5.0 10.0 25.0 2.18 0.316

4 6.02 0.609 7.95 0.246 5.0 10.0 25.0 1.53 0.354 8.07 0.250 5.0 10.0 25.0 1.52 0.355 a b c d e 1.012 M. Per 10 ml. Sample. 0.3096 M. 0.07423 M. 0.04949 M. Table 31

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol Monoanion at 36°

Strong Base Weak Base Chloride Fluoride e Add'n. of Ac0H OH Ac0H a PhO AgNO3 Cl Sample Ce(NO 3 )3 Titration EDTA F b a b b HCF 2C1 ml. M ml. M ml. M ml. ml. ml. ml. M

Initial 8.58 0.868 9.73 0.985

1 7.56 0.765 9.77 0.989 1.35 0.042 10.0 q 10.0 25.0 6.88 0.018 1.37 0.042 10.0 10.0 25.0 6.92 0.017

2 6.21 0.628 9.64 0.976 3.38 0.105 10.0 10.0 25.0 4.54 0.088 10.0 10.0 25.0 4.55 0.088

3 4.63 0.469 9.57 0.968 5.83 0.180 10.0 10.0 25.0 2.70 0.143 10.0 10.0 25.0 2.80 0.140

4 3.96 0.401 7.10 0.220 5.0 10.0 25.0 4.65 0.169 7.04 0.218 5.0 10.0 25.0 4.67 0.168

b c d e a1.012 M. Per 10 ml. sample. 0.3096 M. 0.07423 M. 0.04949 M.

- -

Table 32

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Catechol Monoanion at 36°

Strong Base Weak Base Chloride Fluoride c e Add'n. of AcOH a OH AcOHa Ph0 AgNO 3 Cl Sample Ce(NO 3 ) c13 Titration EDTA F b HCF Cl ml. M 2 ml. M ml. M ml. ml. ml. ml. M Initial 4.29 0.434 4.90 0.496

1 0.093 0.094 4.81 0.487 5.58 0.173 10.0 10.0 25.0 3.33 0.124 5.b0 0.1/3 10.0 10.0 25.0 3.32 0.124

a c d e 1.012 M. bPer 10 ml. sample. 0.3096 M. 0.07423M. 0.04949 M.

CO Cr)

Table 33

The Basic Hydrolysis 'of Chlorodifluoromethane in the Presence of Pyrogallola Mono- and.Dianions at 36°

Total Base Chloride b d AcOH Ph0- AgNO Cl - Add'n. of 3 c c HCF C1 ml. ml. 2

Initial 9.38 0.949 9.37 0.948 1 7.86 0.795 2.83 0.088 7.84 0.793 2.81 0.087

a . b c Initial concentration, 0.50 M. 1.012 M. Per 10 ml. sample. d 0.3096 M. Table 34

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 2-Mercaptoethanol at 36°

Stron Base Weak Base Chloride Fluoride a a f Add'n. of AcOH OH AcOH PhO AgN0c, Cl- Sample Ce(NO 3 ): Titration EDTA F b b b° HCF2 C1 ml. ml. ml. ml. ml. ml. ml.

Initial 11.83 1.270 4.77 0.512

1 10.97 1.178 4.43 0.476 2.41 0.059 5.0 10.0 25.0 6.51 0.059 2.40 0.059 5.0 10.0 25.0 6.53 0.057

L 8.64 0.928 3.88 0.417 6.94 0.170 5.0 10.0 25.0 4.15 0.199 1.05 0.173 5.0 10.0 25.0 4.20 0.196

3 6.18 0.664 3.16 0.339 6.11 0.300 5.0 10.0 25.0 1.37 0.364 5.0 10.0 25.0 1.45 0.359 d 4.38 0.470 2.79 0.300 7.75 0.380 2.0 10.0 25.0 4.55 0.438 d 7.75 0.380 2.0 10.0 25.0 4.51 0.444 a b c d e f 1.074 M. Per 10 ml. sample. 0.2455 M. Per 5 ml. sample. 0.07423 M. 0..94949 M. Table 35

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of 60 Percent Ethanol at 25°

Strong Base Chloride Fluoride a - d e Add'n. of AcOH0H AgNO 3c Cl Sample Ce(NO ) Titration EDTA F 3 3 b HCF2 Cl ml. M ml.b M ml. ml. ml. ml. M Initial 10.73 1.152 10.74 1.153

1 10.04 1.078 1.20 0.029 10.0 10.0 25.0 6.86 0.028 10.04 1.078 1.17 0.029 10.0 10.0 25.0 6.90 0.027

2 6.84 0.735 6.83 0.168 10.0 10.0 25.0 1.84 0.178 6.85 0.736 6.80 0.167 10.0 10.0 25.0 1.98 0.173

3 5.75 0.617 8.68 0.213 5.0 10.0 25.0 4.18 0.216 5.74 0.616 8.75 0.215 5.0 10.0 25.0 4.20 0.215

4 3.73 0.401 12.41 0.305 5.0 10.0 25.0 2.84 0.296 3.65 0.392 12.44 0.305 5.0 10.0 25.0 2.82 0.297

5 0.54 0.058 18.03 0.443 4.0 10.0 25.0 2.03 0.430 18.15 0.446 4.0 10.0 25.0 2.02 0.430

a b c e 1.074 M. Per 10 ml. sample. 0.2455 M. 0.07740 M. 0.04949 M. Table-36

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of Phenoxide in 60 Percent Ethanol at 25°

Strong Base Weak Base Chloride Fluoride e Add'n. of Ac0H a OH- Ac0Ha Ph0 - AgNO c Cl - Sample Ce(NO 3 ) 3 d Titration EDTA F- b b b HCF Cl ml. ml. ml. 2 ml. ml. ml. Initial 9.40 0.998 4.79 0.503 c 9.34 0.981 4.86 0.510

1 7.30 0.767 4.63 0.486 4.06 0.100 10.0 10.0 29.0 5.56 0.067 7.34 0.771 4.59 0.482 4.08 0.100 10.0 10.0 29.0 5.62 0.065

2 5.96 0.626 4.40 0.462 7.00 0.172 10.0 10.0 37.0 3.25 0.138 5.98 0.628 4.39 0.461 6.92 0.170 10.0 10.0 25.0 2.48 0.159f

3 4.50 0.473 4.28 0.450 10:10 0.248 10.0 15.0 25..0 4.57 0.213 4.52 0.475 4.26 0.448 10.12 0.248 10.0 15.0 25.0 4.44 0.216

4 3.13 0.328 4.07 0.428 13.03 0.320 10.0 15.0 29.0 2.30 0.289 3.11 0.327 4.07 0.428 13.02 0.320 10.0 15.0 25.0 2.45 0.275

5 1.84 0.193 3.81 0.400 16.30 0.400 5.0 10.0 25.0 2.25 0.328 1.80 0.189 3.84 0.403 16.24 0.399 5.0 10.0 25.0 2.31 0.321 a c d e 1.0506 M. bPer 10 ml. sample. 0.2455 M. 0.07740 M. 0.04949 M. (Value used in calculations. Table 37

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of m-Chlorophenoxide in 60 Percent Ethanol at 25°

Strong Base Weak Base Chloride Fluoride a Add'n. of Ac0Ha OH AcOH PhO AgNOc Cl Sample Ce(NO 3 ) d Titration EDTAf b HCF Cl ml. ml. ml. ml. ml. ml. ml.

Initial 9.69 1.041 4.52 0.485 9.71 1.043 4.51 0.484

1 8.57 0.920 4.36 0.468 2.35 0.058 10.0 10.0 30.0 7.25 0.053 8.57 0.920 4.38 0.470 2.33 0.057 10.0 10.0 25.0 6.05 0.053

2 6.87 0.738 4.21 0.452 5.85 0.144 10.0 10.0 25.0 3.55 0.127 6.86 0.737 4.22 0.453 5.90 0.145 10.0 10.0 25.0 3.55 0.127

3 5.64 0.606 4.05 0.435 8.57 0.210 10.0 10.0 25.0 1.53 0.187 5.64 0.606 4.05 0.435 8.60 0.211 10.0 10.0 25.0 1.67 0.183

4 4.75 0.510 3.97 0.426 10.65 0.261 10.0 15.0 25.0 3.54 0.243 4.75 0.510 3.98 0.427 10.65.. 0.261 10.0 15.0 25.0 3.56 0.243

5 4.08 0.438 3.88 0.417 12.20 0.300 5.0 10.0e 25.0 3.25 0.252 4.06 0.436 3.91 0.420 12.28 0.301 5.0 10.0e 25.0 3.26 0.252

6 3.13 0.336 3.74 0.402 14.53 0.357 5.0 10.0 25.0 2.51 0.296 e 3.12 0.335 3.73 0.401 14.61 0.359 5.0 10.025.0 2.53 0.295

a b c d e f 1.0738 M. Per 10 ml. sample. 0.2455 M. 0.07740 M. 0.07423 M. 0.04949 M. Table 38

The Basic Hydrolysis of Chlorodifluoromethane in the Presence of p-Chlorophenoxide in 60 Percent Ethanol at 25°

Strong -BaseL Weak Base Chloride Fluoride a d e Add'n. of AcOHa OH AcOH PhO AgNO c Cl Sample Ce(NO ) Titration EDTA F 3 b HCF Cl ml. ml. 2 ml. ml. ml. ml. ml. Initial 9.41 1.010 4.87 0.523 9.41 1.010 4.92 0.528

1 8.99 0.965 4.85 0.521 0.83 0.020 10.0 10.0 35.0 10.16 0.017 9.02 0.968 4.87 0.523 0.83 0.020 10.0 10.0 25.0 7.27 0.016

8.14 0.874 4.76 0.511 2.54 0.062 10.0 10.0 25.0 5.80 0.060 8.14 0.874 4.75 0.510 2.60 0.064 10.0 10.0 30.0 7.03 0.058

3 5.80 0.623 4.46 0.479 7.57 0.186 10.0 10.0 10.0 0.93 0.163 5.87 0.630 4.40 0.472 7.60 0.187 10.0 10.0 25.0 2.28 0.164

4 4.90 0.526 4.33 0.465 9.60 0.236 5.0 10.0 25.0 4.65 0.188 4.90 0.526 4.35 0.467 9.61 0.236 5.0 10.0 25.0 4.50 0.197

5 2.84 0.305 4.00 0.430 14.65 0.360 5.0 10.0 25.0 3.50 0.256 2.80 0.301 4.03 0.433 14.62 0.359 a c d e 1.0738M. bPer 10 ml. sample. 0.2455 M. 0.07740 M. 0.04949M.

(.0

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ma...mmram - immimi■ Imormi MILICKU EN II' MOM 11111111111111

I MIN 111111111 gionegm =now wwww Milli Ishompimmussanim Ming IMINIIIIMEIMIREMMINEMIMIEWER■RE1111 o mmn moismi ommajN ■ 1111111 NOMMEMINIIIIMINIEW11111 111•1111IMIEWENIIIIIIMMIIMINEMNIMMI

Figure 6. Infrared Spectrum of Mono(Difluoromethyl) Catechol (in Carbon Tetrachloride). Instrument Settings: Resolution, 990; Response, 2; Gain, 4; Suppression, 4.

7A T T 10 400 300

-rrimome44frionirpori

0 PM (01

Figure 7. Nuclear Magentic Resonance Spectrum of Catechol Mono(Di- fluoromethyl) Ether (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.0. External TMS Reference. 94

0-0* 0 CPS 100 300

Figure 8. Nuclear Magnetic Resonance Spectrum of Catechol Mono(Di- fluoromethyl) Ether (in Carbon Tetrachloride). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.0. External TMS Reference.

T )-2* 000 ICA 0 CPS

PPM ) I)

Figure 9. Nuclear Magnetic Resonance Spectrum of Deuterodifluoromethyl Phenyl Ether (in Carbon Tetrachloride). Instrument Settings: Filter Band Width, 1.0 cps.; R.F. Field, 0.20 mG.; Sweep Width, 500 spc.; Sweep Time, 500 sec.; Spectrum Amplitude, 25. External TMS Ref.trence.

100 1111111 1 I III If i 11111 Mt. 1 11111111h 1111111111 Ik1 I 111111111 111.111.111 111 111111 111111111111111111 11111111M11111111 III If :4+1111E al I11111111 III DIEM 1111111 1111111 _'1 -- 111111 IN III 11111111 I M 111111 II II Ili H11111 101 1101111 I 1111111.11E 11111 11111 II II 11110111111 1 1 I 11111111 1111F11111111111i1iii UM !kir nn8111111 iM PI I 11 Ii1111111111 111111011 1111 I 11111 II II 111111 11111 111111110110 I IIIIII III 11111111 1 11111 11111 1 1111 1111. 11E, I!11111111,P11 !011111111111F1 P111111,11111 Mk' rig 11.1111 I ' 1P 111111 so 1111111111 I 1 1 I I I II II 1 I III I I 1911 III 11111111111 II 11111 1 0111 m mmon mum mum muuma m iMmoomm au"' I mpf1111111111 1 I I 1111 1111 111 110 I 1111 11111 II I 111111111 1111 1.1 I MI II 0111 MI ' II 1111111 I 111 1 III 1111111111111 4 II 11111111111111 1 8i1111 Fp u, : m1 1 11111 I I 11111111111 III I MI MI 10111111 II 11111111111 11101 I 1 111 IIII 1111 1 mu Iii11111 1 II I 111 111111111111111MM 11 01111111! 11 III ii IIIIIIIIIii 4'0111 1111 1111111111 II I 1 1111 III 111111 MIMI 11I II 1 111 1 1 1 1111 111111111 1 11111 11111111111111 III 11 M11111111 II 1 1011111111 II 111111;111111111111111 I HM1111111E11 1011 11111101 111 III I 11111 1 I III III 1111111 I 1111 I 11 III 1 111 11 II 111111 II 111111E 111111111111111111 MI 1111111110110 III I I IIIMIE 111111111 11111 11111 1111111 1 I 1 11111 III 111 II III M0 111111 I 0111111 IOW I 11 MI I 1 111111 II I 1111111 11 II I I II 111111 111111 1111101 III I 1111111E11 11111 I 1 IIII'MME .1 111111/1 III 1111 Oil 111 111111 II 1111111111M111 1111 1111111 1 111111111111 I I II II 1011 II 111111111 1 111 Hull 111111 1111 1111111 111111111 11111 111111111 I. II III 1111111111 III 111111119 011111111101110 I II 11111111111 II III 1111 11 11 II III 0 11001 II 1111 1E1E11 I MIEWIMIll I11111 100111111 1 1111111111111 1111 1111 I II I 1111111111 I 1111 HI 11111 111111111111111111 111 11111111 I 40 E HUE El 11 1111 111111 11 111111111 1111111 min ma 011 11111 11 III 119.11 III II I UHMEUM IN 11 MI III 1111111 I 1 1 1111111111 Rom III 111111 II 111111 1111 11111111 11111111 11111111 mmumummu 111111111 11111111 11111111111MM 1111 1111 III II 11 11 11111111 MI 1111111 I 11 11II 1111 I 1 111 I II II 11 I 11 II 111111111111 1 1111E111 11 11 11111 MED 11 11111E111111111101.MEN1 111111111 11 I III I II II 1111110 MEE 111111111 I 1 II 101111 III 1111111 MI II 11111 1111 II I 11 1 11011 I 11.1111111111111 111111111111 11111MMI11.111 1111 MI01111111 11.11E1REMEMEN III 1111111111 11111111 11111 fillEMIU 1 I 1111 11111 MI 1 111 111111111MEEMEM 11111111 11 11 I IF III II I III 111 I 11111111 MM11 110 1111111.1111 111111111111 1111111 INNEN II MIENIM III III 1111 III 111111 111111 1111 III 1 Mil 11.11 III 2 111111 1111 II MI 1111111111111MM 11 0 1 II 1 II 111 El 11 1101 I II 111 1111 II 1111111111111 1111111111111111111111011111111111111 I I IMM111 11 1111111111111 10001 II '1111E11 11111 1E1 1 II II 1 III II II I I II II 1 010011 11111 11111111 II 11 III II 11111111 1111111111 II 111 NEER II 11111 1111 1111111111 11 on III I J1.11111". III 11111 111111E111110 1 11 11111 MI III MI II III 1 11111 1111.1411 III. II WM 1111111 I 1111111 1 II 1 III 11111 11E101 II 111111111 1 II 1111 III 11111 11 III I 1 1111 'ME 11 1 , 1111 1111111 11111EIMI NIE 111111 111111 1 DI III 110111 I II II 1111110111 11111 11111 11E111.1111 I 111111111 111 I III 111111 II 1111111111 II 10111E1111 1111 II 11111 111111111 I I 111111111 11 ! IIII1 1 •

Figure 1 0 . Infrared Spectrum of Difluoromethyl Azi de ( in Carbon Tetra- chloride ) . Instrument Settings : Resolution, 9 2 7 ; Response, 2; Gain, 6.5.

I NI 0.0 III 11111 111 11.1111 ... 11111.1111,1 11111 1111111 1111111111111111 1111 1 1111 11111111111111.111111111 111111111111 ki 111[1111 0111111111101 111 00 111 02 101 111111111N11 111 11111 111111 OA 111111i 1011111 111 I It HI If 11 110111 HE I MEE 11111 111,11 1111r1P, 911111111111111111 11117,110110 HP 111111111a 0.6 0.. 11111111 II ' IIIMMEE111111111111111EM III 11 ENID 1111 E11111 III III III 1111111 I 11E11 11111I II EMWEE III II 11111111 MINIM 111M11211111111111E1111021111111111111111M1111MIIIIIMII I 1111111111111111INKIIIMIIIIIIMIIIIIEW11111 III MI EWE 0.8 11111111111E1111111111111111=11111=1111111111111MOIME111111111111M1111 111111811011011111111 MEW11111111111MIKWINI111111111111 11111 1111111111E111111111111 1111Thi lE1 IIII 11 .0 11111 1.5 RINUMR1111111111111111111111111111111111111111111111MW1M11111111111111MOMM1111111 11111111 11111111•111111 11111 1111111 1 111111111 111111111111111111111111 11 1 1111 1 11111111111 15

Figure 11. Infrared Spectrum of Difluoromethyl Phenyl Ether (in Carbon Tetrachloride). Instrument Settings: Resolution, 927; Response, 2; Gain, 5; Suppression, 5.

al• 7.0 0 NW (I

1-11* 010 270 00 OM

Figure 12. Nuclear Magnetic Resonance Spectrum of Difluoromethyl Phenyl Ether (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 0.63. External TMS Reference.

100 100 III 1111111 I II III 11 RI II u 1 I u.1 P 1 IWO I .. men raii ill lumoull uunuerumumuu 1u2uum mu muulloqiummuu..i won mull ul:III II mow 11 1 mum 111111111 III 1 II I li 11 I , III II II 1 1111 11111 1: 11 1 111111 III I 11111 4 III 1 11111 Inv ;.■ 1 11111 11111111 1 i.1111111fililidel 1111111111111111111111111111 I 1 1 II 1111111 III IN I Ilion11111! . 1111111 Ill I 111 11111111111W MI I I I R 1111 I 1 1111 111111 I l'il MI I 11 1 , 11 1 1111 111111 I II 111 1111 111111 I 111111111 1111 111111111111 1111111111 111111 11111111111111 I 111111111111111 III 1 1111111111 11111111 III 1111111111 III 1111 III I P I 11111111 11111111 1 1 I I I 11111110 111 III II 1111111 I 1 111111 RH 11111 I 111111111111 1111 1111111M11111111111111111111 III 111111 III I 11111111111 80 1111111 111 1111111111111111 11111 II I III II I 1 111111111 MI 1111 1 11111111 1111111 I 111111 1111111 1 I IR I 111111111 MI11111111111111 MI111111111111 1111111111 11 1 1111 II 1111 11111111111111 1111111111 1111111E 111111111111111111111111 111111111 111111 111111 1111 1111 11 I 1 I 1110E11 1 III I limuluill II ii II u 11111111 nu 1 NI 11110111 1111 111111111111 11111111111111 III 11111 111111 11 II 1111111111 RI 1111 11111111 111111 11111111 11111 II I 1 111111111111111111111 1111 1111 1111111111111111111 111 1 II 111111 111111111111 111111 11111111111 .A1 1111111 I 11111111101 11111111111111111111111111 1 HIIII III II 1111111111 1111111111111M 11111111 1111011 111111111 I I I I I 11111 II 11111111111151111 01111111 1111111111 I 1 I 1 111111 MI 1111111111111 II III 11111 11111111 1111111111111111 111111 111111111 11 111 II III 11 111 III II I 01111111 1111111 111111111 1111111 I 1111 III I I 1111111 11111 111111 1111 111111111111111111 RR 1 1 IR 1111111101111 111111 11111111111111111111 1 111111111 11111111111111 1111111111111111111 1 10E II MIMI I 1111 II REM 1111111 1111111111111111111 I I 011 111 II 1111111 111111111 11111111111111111111111111111 1111111 111111111111 1111 111 1111111111111111111111111011111111111111111111 11111111111 111 111111 1111111111 101111111 I 11111 R EMO 11111111111.11 111111111111 1111 1 1 1111111111111 1 1 1111 Mil I MI III I III MI 111 1 11111111111 11111011111111 1 II 11111 HI 11 I II II I 1111111111111 II 111111 1111111111111 111111111011 1111 1111111 11110 111111111 III 11111/111111 11111111 11111111 1 MI 111111 hIIII NOVI N 1111111111111 111111 1111111111RM MI 1111111111 111111 11 11111 I II 111111111111 111 40 11111 II 1111111111111111 II 1 111111111PM III 1111111 III 111111111111111111 HI 1111 11111111111 1111 III II MI 11111111111111111 1 11111111111111111 IMP111111111111! 1111 111111111111111 111 111111 I 11 II 111111 111111111111101111. III 1111111111 11111E1 III 11111 11110111 1101111111111111111111 11 III 11111 1111111 1111 I 11111111111111111111111/11, 11111111111111 III 1 110 III 11 111111 11 111111111111111111111111111 111111111111 011111111111111111111 111111111 RIM 1111 1111 II 1111111 NI 1111111111111 1111 111111 1111 I II 1 M111111101' 11111 11111111111111111111111 11111111 11111 11 111 II 11 11 HIM 111111111111 11111 II 111 I III 1111111 1 11111111111111111111 III 11111111111111 III 1111111111111111 11 111111111 1101 III 1111111 111111111111111111111111111111111111111111111111111101 1111111111111 1111 MI 11111 I 1 I 1111111111 2 111 111 .1 in 11!111 II umumunimu 11 i ill i 1 111111 III 11111111111111111 II 11 11111111 III III 111111 III 1111 11111111111111111111111111111101011111111111q4MP.P111 III 1 MI 1111111111111 1111111111111 IIU I 11 111 ' 1 ' III Ill 1 11111111111111111111 II 1 III I 1 .11111111 1111 11 NNE II 111 MI II I 1111 III 11111 11110111 1111111111111111111111r —"" " ""- ' ' 1111111101111 11 HEM III r 11111111111111 1 11 Ill /11111111 I 11111111111111111111 11 11111 1 1 11 11111101111 111111:1111111 1 II IIIII0O li 1110!11 11111 111111111111 1111111111 11111111. 1 11111111111 1111111 1111.111.1.11111111 1 :11 1 11 11L. = 11111111111111111111 111111P1111111111111 II 11111111 1.111111111 111100 11 IN 111111§101 111111111 IIIIIM 11:::::::i iicliliiIIIhI NM 1101i1 111111111111111M1111111 1111

Figure 13. Infrared Spectrum of Difluoromethyl p-Tolyl Sulfide (in Carbon Tetrachloride). Instrument Settings: Resolution, 927; Response, 2; Gain, 6.5.

Figure 14. Nuclear Magnetic Resonance Spectrum of Difluoromethyl p-Tolyl Sulfide (Neat). Instrument Settings: Filter Band Width, 4.0 cps.; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec. External TMS Reference.

lll UM= llll M ll M lllllll M lll . lll p II llllll I m 1111 1 llllllll m l 111111;hllii llll I 110"11 l 1111 ,P191111111111. l I 1111 111110101111 1111 11111'llllll I;"1 ill i I 111111111111111 11111 .111.1a111 1111 llll 11 1 .111 11 11 1 „I .11111 1 n 11 111 11 n1 11111111111,111.111 111. 111 ll 0111111111111 Nil I I hi OTII 11111111111 II I 111111R111100 III I 111111 111111 1 IN 11. 1111 NM 1111111110111illiiiii iii

1 0 111 Mil limo hi Huai u/ lir 11 11 1111111211 01111 1111111 11011! I MI iliN 1,11101 111111111011111111111111111111111r,110 II 1101111111111111111111 111111 11010011111111,111110 11111011111111 111111110111111Plli 111111 111u milmiloominimmienr1111111111111 111111'11 1111111111111111111111hililitlilliiiiiii 11 I 1111111111111111111111110111 11111111111111111111111111101111111111011111 rill hi IR 1111111111 11111 1111111E11 1 1 111E1111101111111 11111111111111111111 ■111 111111111111111111111111111111 M II I 111111110111 000111 11111111111111 111111011111101111111110 11 11111 111111111111111111111 11111111111111111 101111 11111 11 1111 1111111 summumilo 110111101111111P' 1111 101 11110111 1110111111011111111111110111001111 1' Ni 1111111111111111INIMNIMOMM111111111111111111 111111M11111111 MUM 111111111,1M111111MNI HI II 11111 M11,111111/11111111111WINIMI MIIMMH MINIM 111111111WIMIAMIIIIIIIJIM111111111111111111u1 III 1 M11 111111110M1111111111 111111111111111111MMII11111111111111111111111111111 1111111 11111111111111M11 mifigums miamoui 1111111111111111111111111111111111111111111 IMMO' 11111111111111N111111M1111011111111111111111 M11110. 11011111111111111111111111111111111111111111111 II 111111 11111 11111111111111 11111111111111111111111 11111111 1111111111111 If di! 11111111111M111111110MINNW1111W11 HMI HI 111111 Hum 1111111111111111111111111 11111110111111111111111111111111111111 1111111111 111111111111111111110111111111111MM 11111111 11 IMIMMUI1111111111111111111111 111 1111111 1111911111111111111111111111g1 kl1111111111 1111 MAN unniumminimmusimuutuutignmffimuill11111MERMINIMMIffill1111 111111 u 1111111N1111111111111111 I I 111 11 111 II 1111 1 1 1 1111 11111E1 111111 1111111111 P11 1111111 1110111 IN 1101 I III 110 11111111111111110111111011111111 110'1 011111 11111101111111 110 Y01110110111110111111110121911111 10011110111 000111 111111111110111 nommounu NI Hop orifflommi m lo I holutuumminimu ti 1 moimmimu ImuumuM111110111%1111111111111111111111ffiffill1111111111111iiilili 11111111111111111111111111111111111 1111111111111111111 11111 1111111111111 MIEN 11 11111111P 1 III 1111111 111111 WHO 111M1111111111 1111 IM11111111W111111111 1111 1 1 1 1 111111 I 1111111111111111111111111 11111 M1111111111111111111/11111111111111111111111 11111 1111111111.0111111111111111'11 Iroymnprzin roman 11111111111 111 11 111111 I 1111 muumuu I 11 1111f/1111 1111111i1 1.1111121111111q1 1111 11111111111111 11 1 111,111 prommmr1111mmunou 1111111111111111111MP 11111111H i 0111 im mill NOON I; 1111111. "" iii11111

Figure 15. Infrared Spectrum of Difluoromethyl m-Chlorophenyl Ether (in Carbon Disulfide). Instrument Settings: Resolution, 930; Response, 2; Gain, 5; Suppression, 4. 98

Figure 16. Nuclear Magnetic Resonance Spectrum of Difluoromethyl m-Chlorophenyl Ether (Neat). Instrument Settings: Filter

Band Width, 4.0 cps ..; R.F. Field, 0.02 mG.; Sweep Width, 500 cps.; Sweep Time, 500 sec.; Spectrum Amplitude, 1.25. External TMS Reference.

min 1 llllll Ilnla Ill II I lll II 1111M11611111111111111111111 Me IIIII 1 Inir1=1111111•111MISMNII MUMMY WWI mum= ...... 1 111111J 11111111111111 11111111111111111 1 11111111111111111'11111111111 11111111111111111111111111111111 111111 1111 NMI 1 I 1 III 11 1 111 I 11111111 1 11Nf111111 11 1 11111 1 11111 111111111111111111111 111111111 11111111M milMIM ImilAIMMINI 1111111111111AMIIIIFIIIIIVIRRINNI 0010111 mmwomm son I I MI NM Pill NV NI NE 1110 11111111111111111 Milo INNUMINEMINE111111 I I' 11111 UMMINIMEHIMONME NOM I 11111111 INTHANN MOOMIIN INERINIMMINIMMI IMMEMOCHIMMMEMOOMMEHM MNIIIMII MAIM MINMUNOHMEMMEA NEM AMMOOMMINE 111111111111111111MINIMEMUME I I 1111111 NH EMI HIM ENE mom now 1111111111MMINIMMI 111111101111011Elin own mil IIMMINIMPIMMUMMM1111111 NM 11111 MINIM= MIN I M I MI MDR IIIIIIMMIMIMMUMMIIMMUMMIIIIMAIMMEMEMMIMM ill mm mmummEMMEEMIIMArmr4MMMI a mMuOMMIAM m 1m malle 511p 1 UNN111111151111111mollilillismom IllUE11111111111111611MEMMUMMMIIMMMUMME MUM ELHNI MI MUNUMMENBRESEM MEMMEMMOOMM 11111111111111 IN normri Lome NOM MI VEIN . Iu 11111111 jniurilififil 1 1 1 e.....1...... 1111211"11 '1'152 mallingS12 1". 112. 5 111121151..... II No so Nom Iron 1 on u ME 11 MII N HI II NM= 111110110 11 110111101111111111011111111 IIIVERMEN II .1.1101 ii 911111 III fill 111111 II 111111 WW1 1 11 111111111 111111 1.9111M11 II Nal ! Kil II IIIWIIIIIMI MMIIP9 IMMIWAMH M 111111111111 1111111111ffil MOM= MI - II 1111111101111111 III

Figure 17. Infrared Spectrum of Difluoromethyl p-Chlorophenyl Ether (in Carbon Disulfide). Instrument Settings : Resolution, 930; Response, 2; Cain, 5; Suppression, 4. 99

LITERATURE CITED

1. J. Hine, Journal of the American'Chemical Society, 72, 2438 (1950).

2. J. Hine and A. M. Dowell, Jr., J. Am. Chem. Soc., 76, 2688 (1954).

3. J. Hine and P. B. Langford, J. Am. Chem. Soc., 80, 6010 (1958).

4. J. Hine and S. J. Ehrenson, J. Am, Chem. Soc., 80, 824 (1958).

5. J. Hine and F. P. Prosser, J. Am. Chem. Soc., 80, 4282 (1958).

6. J. Hine, N. W. Burske, M. Hine, and P. B. Langford, J. Am. Chem. Soc., 79, 1406 (1957).

7. J. Hine, Divalent Carbon, The Ronald Press Co., New York, N.Y., 1964, pp: (a) 36-38-, (b) 38-41.

8. J. Hine, Physical Organic Chemistry, John. Wiley and Sons, New York, N.Y., 1962, pp: (a) 484-486, (b) 488, (c) 486-488, (d) 69-70, (e) 230-231, (f) 491-492, (g) 253, (h) 269, (i) 161, (j) 294.

9. J. Hine and K. Tanabe, J. Am. Chem. Soc., 79, 2654 (1957).

10. J. Hine and K. Tanabe, J. Am. Chem. Soc., 80, 3002 (1958).

11. J. Hine and J. J. Porter, J. Am. Chem. Soc., 79, 5493 (1957).

12. J. Hine and J. J. Porter, J. Am. Chem. Soc., 82, 6118 (1960).

13. T. G. Miller and J. W. Thanassi, Journal of Organic Chemistry, 25, 2009 (1960).

14. C. G. Swain and J. F. Brown, Jr., J. Am. Chem. Soc., 74, 2534, 2538 (1952).

15. V. Franzen, Chemische Berichte, 88, 1361 (1955); 89, 1020 (1956); 90, 623, 2036 (1957).

16. P. S. Skell and Garner, J. Am. Chem. Soc., 78, 5430 (1956).

17. W. v. E. Doering and W.'A. Henderson, Jr., J. Am. Chem. Soc., 80, 5274 (1958).

18. H.A. Flaschka, EDTA Titrations, Pergamon Press, New York, N.Y., 1959, pp: (a) 117, (b) 63, (c) 113-115. 100

1g. S. S. Yamamura, M. E. Kussy, and J. E. Rein, Analytical Chemistry, 33, 1655 (1961).

20. L. G. Bellamy, The Infrared Spectra of Complex Molecules, Methuen and Co., London, 1960, pp: (a) 32'8.-9, (b) 273-4.

21., C. E. Meloan, Elementary Infrared Spectroscopy, The Macmillan Co., New York, N.Y., 1963, p. 169.

22. E. A. Robinson, Journal of the Chemical Society, 1663 (1961).

23. C. G. Swain and C. B. Scott, J. Am. Chem. Soc., 75, 141 (1963).

24. M. Green and R. F. Hudson, Proceedinglof the Chemical Society of London, 149 (1959).

25. C. T. Abichandani and S. K. K. Jatkar, Journal of the Indian Institute of Science, 21A, 417-41 (1938).

26. R. F. Jameson and W. F. S. Neillie, J. Chem. Soc., 2391 (1965)

27. C. F. Timberlake, J. Chem. Soc., 4887 (1957). PART II

THE REACTIONS OF THE METHYLENE HALIDES WITH ALKOXIDES IN ALCOHOLIC SOLVENTS--A SEARCH FOR AN&-ELIMINATION MECHANISM AND METHYLENE INTERMEDIATES 101

CHAPTER I

INTRODUCTION

Background

The Reactions of Mono-, Di, and Trihalomethanes with Strong Bases

Modern physiCal organic chemistry is based to a large extent on studies relating the structure and reactivity of organic compounds. One of the interesting aspects of such studies lies in determining when, in the course of such structural alterations, a change in the reaction mechanism is effected. A classical example of such a change is found in the work of Hughes, Ingold, and co-workers (1-3) who found that the basic hydrolysis of methyl bromide in 80 percent aqueous ethanol yields methyl alcohol. The rate of the reaction was found to be first-order in hydroxide, first-order in methyl bromide, second-order overall. If tert- butyl bromide is substituted for the methyl halide and the reaction carried out under identical conditions, the rate is found to depend solely upon the halide concentration. The change in the order of the reaction, indicating a change in the reaction mechanism, i.e., S N2 to S N1, illus- trates how suitable structural alterations can effect a change in the reaction mechanism.

As the hydrogens of methane are successively replaced by halogen yielding first the methylene halides, then the haloforms, and finally the tetrahalomethanes, the question logically arises as to whether the basic hydrolysis of each series of compounds proceeds by the same nucleophilic displacement mechanism found in the monohalides. This question has been 102

partially answered by Hine and co-workers (4-12) who found that the basic hydrolysis of the haloforms proceeds by an a-elimination mechanism (1).

HCXYZ + OH- ÷ H2O + CXYZ - (1 )

CXYZ- F CXY + Z -

CXY + H 2O -4 CO + HC00 - + X- + Y-

The a-elimination mechanism involving a dihalomethylene as an intermediate is supported by the following observations (4, 5): 1) chloroform (where

X, Y, and Z are all chlorine) undergoes base-catalyzed, aldol-type con- densations with carbonyl compounds; 2) all of the haloforms, except those containing two fluorine atoms, undergo base-catalyzed deuterium exchange more rapidly than they hydrolyze; and 3) capture of the intermediate dihalomethylenes by various nucleophilic reagents including halogen illustrating a mass-law effect. The latter is perhaps the strongest single piece of evidence for the dihalomethylene intermediate.

Although the reactions of the monohalo- and trihalomethanes seem to be rather cleanly divided into two distinct mechanistic classes, evidence is available that the dihalomethanes (depending upon the experimental conditions) may react by either a nucleophilic displacement process or an a-elimination mechanism. Hine, Thomas, and Ehrenson (13) determined the rates of reaction of several methylene halides with iodide in acetone and with methoxide in methanol. These reactions were discussed in terms of the S 2 reaction mechanism and were carried out N primarily to determine the effect of an a-halogen on the rates of nucleophilic displacement at carbon. As shown in Table 1, the relative rates of reaction of the methylene halides toward iodide in acetone 103

parallel those toward methoxide in methanol; however, the differences are not as large in the alkoxide series as they are in the halide series.

Table 1. Relative Rates,of Reaction of the Methylene Halides with Iodide in Acetone and Methoxide in Methanol at 50°

Methylene Halide k /k k /k 1 o 2 o

CHIC1 0.38 77.5 2 CHCl 1.0 1.0 CH022Bro 327 13.0 CH`BrI 236 17.9 2 CHClBr 515 41.7 2 CHFBr 3200 3960 2 CH 2 I 2 18.9

CH • k iodide in acetone; k 2 , methoxide in methanol. ko' 2 Cl2' 1 , '

These differences in the relative reactivities may be indicative of a change in reaction mechanism. Due to the low nucleophilicity of iodide toward hydrogen, i.e., basicity, it seems unlikely that this reagent could effect removal of a proton to initiate an a-elimination mechanism. On the other hand, an a-elimination mechanism and perhaps a methylene intermediate is plausible for the reaction of methoxide with the methylene halides, due to the greater nucleophilicity of the alkoxides.toward hydrogen. The differences in the relative reactivities of the two nucleophilic reagents toward the methylene halides could, therefore, be explained by a change in reaction mechanism. For this 104

reason and others which will be discussed shortly, an investigation of the reactions of the methylene halides with alkoxides was undertaken.

Several workers have obtained evidence which indicates that the reaction of strong bases with the methylene halides proceeds by an a-elimination mechanism. In addition, divalent-carbon intermediates have been suggested as intermediates in these reactions. Kursanov and co-workers (14) found that when a mixture of benzene and methylene chloride, bromide or iodide was treated with dry potassium tert-butoxide, tropilium tert-butoxide was formed in yields ranging from 0.1 to 1.4 percent. A plausible reaction mechanism is shown below:

tert-BuO - + CH 4- tert-BuOH + CHM:. (2) 2 C12 2

CHC12 CHC1 + Cl -

tert-BuO -

Although several workers have seemingly regarded the formation of cyclopropane derivatives or products derived therefrom as proof for the formation of methylene intermediates (15-17), more recent work has shown that this is by no means necessitated (18-21); an organometallic reagent may be employed to explain the results in an equally satisfactory manner.

In Kursanov's work, the reactive intermediate may be NaCHC1 rather than 2 monochloromethylene. The extent to which react,itn (2) proceeds through a free methylene intermediate is not known; it may be that the reaction proceeds entirely through a methylene intermediate and that benzene, being a poor methylene capturing agent, converts only a small fraction of 105

the divalent-carbon species to the observed product. On the other hand, benzene may be a good capturing agent and only 0.1 to 1.4 percent of the reaction is actually proceeding through a methylene intermediate

However, despite the lack of evidence for a methylene intermediate, it is clear that the reaction is proceeding by an a-elimination mechanism.

Closs and co-workers (15-17) have obtained evidence for an a-elimination mechanism in the reactions of several methylene halides with organolithium compounds and have suggested that these reactions proceed largely or entirely through a methylene intermediate. These workers found that when a mixture of methylene chloride and anlolefin was treated with an alkyl lithium compounda halocyclopropanes were obtained. The insertions were found to be stereospecific suggesting a singlet monohalomethylene intermediate. They also found that the treatment of methylene chloride in benzene with methyl lithium leads to a

20-percent yield of methylcycloheptatriene (22). Again, the products were rationalized in terms of a monohalomethylene intermediate. However, it must be reiterated that although these reactions are clearly proceeding by an a-elimination mechanism, the evidence for a free methylene intermediate is not as strong as some people would like.

Further evidence for an au-elimination mechanism in the reactions of the methylene halides with strong bases comes from the recent studies of Blanchard and Simmons (18). In their investigations of the reaction of methylene iodide with the zinc-copper couple, the need arose for dideuteromethylene iodide. It was found that this compound could be prepared by direct deuterium exchange of the methylene iodide in

C H OD-D 0-CH ONa. The fact that the dideutero-compound could be 2 5 2 3

if I, 106

prepared in this manner indicates that its rate of deuterium exchange is faster than its corresponding rate of hydrolysis and/or alcoholysis. In addition, Blanchard and Simmons state that the deuterium exchange was carried out in the presence of added sodium iodide to "suppress carbene hydrolysis." This, of course, is simply a synthetic application of the mass-law effect previously demonstrated by Hine and Dowell (7) in the basic hydrolysis of the haloforms.

Hine and Dowell showed that added chloride ions slowed the basic hydrolysis of chloroform. The role of chloride in retarding the hydrolysis rates is best understood in terms of mechanism (1), i.e., the chloride ions capture the intermediate dichloromethylene converting it to the trichloromethyl anion. Bromide and iodide were found to behave similarly but more effectively.

The mass-law effect constitutes some of the most convincing evidence for the presence of a dihalomethylene in the basic hydrolysis of the haloforms; however, the absence of such effects is not equally satisfactory evidence that divalent-carbon intermediates do not exist.

In order for a mass-law effect to be observable, the intermediate must be sufficiently stable to be selective in its reactions with the various nucleophilic reagents present; otherwise, it will combine with one of the first molecules it meets, very probably a solvent molecule. For example, there appears to be no direct mass-law evidence that simple aliphatic tertiary halides react by the S N1-reaction mechanism in hydroxylic solvents, although there is considerable other evidence that they do. 107

Purpose

The purpose of this investigation was to study the reactions of several methylene halides with alkoxides in alcoholic solvents and to determine if the reactions are entirely S N2 in character as previously

assumed (13), or whether they proceed either entirely or in part by an , a-elimination mechanism.

Approach

To realize the above mentioned objective, the basic approach previously employed by Hine and co-workers (6-12) in establishing the mechanism of the basic hydrolysis of the haloforms was followed. This approach, with minor modifications, may be outlined as follows:

1) determine the rate constants for formal formation; 2) determine the rate constants for deuterium exchange, 3) determine if a mass-law effect is observable, and 4) determine the deuterium content of the formal produced. With these data, qualitative and quantitative arguments can be presented to support either the S 2 or the a-elimination mechanism. N 108

CHAPTER II

EXPERIMENTAL

Instrumentation

Ultraviolet Spectra

The ultraviolet spectra were recorded on a Cary, Model 14,

Recording Spectrophotometer using 10 mm. Corex cells (Beckman).

Nuclear Magnetic Resonance Spectra

The nmr spectra were recorded on a Varian, Model A - 60

Spectrometer.

Infrared Spectra

The infrared spectra were recorded on a Perkin-Elmer, Model 21,

Recording Spectrophotometer.

Potentiometric Titrations

The potentiometric titrations were performed with the aid of a

Beckman Zeromatic pH Meter using the glass-calomel electrode system.

The potentiometric determinations of halogen were performed with the same instrument using the silver-glass electrode system.

Constant-Temperature Bath

A constant-temperature water bath (Sargent No. S-84845) and a

Sargent Thermonitor (Sargent No. S-82055) were used to control the temperature, 0.010_ Chemicals

Methylene Bromide Eastman.

Methylene Chlorobromide Eastman.

Methylene Chloride Eastman:

Methylene Iodide Eastman.

Deuterium Oxide U.S Atomic Energy Commission.

Methyl Alcohol Baker Analyzed.

Isopropyl Alcohol Baker Analyzed. tert-Butyl Alcohol Baker Analyzed.

Sodium Baker Analyzed.

Potassium Baker Analyzed.

Mannitol Baker Analyzed.

Methyl Borate Anderson Chemical Co.

Potassium Carbonate Baker Analyzed.

Sodium Hydroxide Baker Analyzed.

Perchloric Acid Baker Analyzed. 110

Reaction of the Methylene Halides with Alkoxides

Determination of the Rates of Formal Formation

The methylene halide to be studied was weighed into a dry, nitrogen-purged, volumetric flask and dissolved in the alcoholic solvent to be employed. An aliquot of a standard solution of the conjugate base of the solvent (as either the sodium or potassium salt) in the corresponding alcohol was then added and the reaction mixture diluted to volume with the alcoholic solvent. After thorough mixing, an aliquot of the reaction mixture was withdrawn by pipet and transferred to a 100-ml. beaker containing about 20 ml. of distilled water and subsequently titrated with standard perchloric acid to a phenolphthalein endpoint.

From this analysis, the initial concentration of the alkoxide was obtained. Since the methylene halides are essentially insoluble in water, addition of the reaction mixture to water effectively quenched the reaction. In some of the studies, the initial base, concentration was not determined by direct analysis. Instead, several analyses were performed before the reaction reached five percent completion. The initial alkoxide concentration was then determined by, plotting the alkoxide concentration versus time and extrapolating to zero time.

The remainder of the reaction mixture was transferred to a dry, nitrogen-purged, serum bottle. The serum bottle was then capped and the reaction mixture placed in a constant-temperature bath. Periodically, samples were withdrawn by means of a nitrogen-filled syringe and an aliquot analyzed by titration with standard perchloric acid.

The serum caps employed in these studies were previously extracted with methylene chloride to remove substances which often caused color to 111

develop in the solutions if this precaution was not taken. The caps were

subsequently heated to 70Dto 80° to remove the absorbed methylene

chloride and then stored in a desiccator along with the serum bottles.

Timing the Reactions. The initial time in these studies was

recorded when the serum bottle was placed in the constant-temperature

bath. The entire procedure for preparing the reaction mixtures, from the

time the alkoxide was first added to alcoholic solution of the methylene

halide, to the time the serum bottle was finally placed in the constant-

temperature bath, required only about three minutes. Since the prepara-

tion of the solutions was carried out at room temperature, ca., 25°, and

since it took a comparable amount of time for the reaction mixtures to

come to reaction temperature, i.e., 36°, recording of the initial time when the serum bottles were placed in the constant-temperature bath

seemed to offer the best reference point. However, all of the reactions were sufficiently long, i.e., two to eight days, so that an error of as much as two to three minutes constituted an error of less than one percent.

Homogeneity of the Reaction Mixtures. Only those reactions

carried out in methyl alcohol and methyl alcohol-d remained homogeneous

throughout the entire study. The alkali-metal halides begin to precipitate almost immediately (before one percent reaction) when either

isopropyl or tert-butyl alcohol is employed. In these studies, samples were removed from the clear, supernatant liquid without disturbing the

salt precipitating on the bottom of the serum bottle.

Reaction of Methylene Iodide with Sodium Methoxide in the Presence

of Sodium Iodide and Sodium Perchlorate. The basic procedure for

■•■ 112

determining the rate constants in these reactions was the same as

previously described. The sodium iodide and sodium perchlorate employed were dried overnight at 110° and then stored in a desiccator over

phosphorous pentoxide. The salts were weighed into dry, 50-ml. volumetric

flasks which had been previously tared (with stoppers). The volumetric

flasks containing the salts were again placed in an oven and dried overnight at 110°. After the second drying, the flasks were transferred to

a desiccator and allowed to cool and then reweighed to determine the

exact amount of salt added. This rather lengthy procedure was necessitated by the extremely hygroscopic nature of sodium perchlorate which

visibly absorbed moisture before the anhydrous salt could be transferred

from one flask to another by conventional weighing techniques.

Preparation of Di(tert-Butyl)- and Diisopropyl Formal

The preparation of di(tert-butyl)- and diisopropyl formal, by the

reaction of potassium tert-butoxide and isopropoxide with the methylene halides has not been reported; consequently, these reactions were carried out on a preparatory scale to ascertain whether the expected formals were

indeed the principal reaction products and to determine if any other reaction products could be isolated.

Preparation of Di(tert-Butyl) Form'al. Potassium, 42 g. (1.07 moles), was dissolved in 1000 ml. of dry, tert-butyl alcohol in a 2000-ml.

flask equipped with a condenser and , protected from the atmosphere with a

drying tube containing Drierite. After the reaction between tert-butyl

alcohol and the alkali metal was complete, methylene bromide, 100 g.

(0.574 moles), was dissolved in the alkoxide solution. The mixture was stoppered and allowed to stand for about 30 days at room temperature. 113

The clear, supernatant liquid was decanted from the salts that had precipitated during the course of the reaction and then distilled through a Claisen head to a pot temperature of 98°. Upon cooling, more potassium bromide precipitated from the residue remaining in the distillation flask. The salt was dissolved in a minimum amount of water and the mixture transferred to a separatory funnel where the aqueous layer was separated. The organic layer was then dried over calcium chloride and subsequently distilled on a spinning-band column.

The initial boiling point of the mixture was 80° to 82°, the boiling point of tert-butyl alcohol. After removal of 1 to 3 ml. of the alcohol, the temperature rose to 96° to 98°, the boiling point of methylene broMide. After,removal of a small amount of the methylene halide, the temperature rose rapidly to - 145°; the boiling point of di(tert-butyl) formal. The distillation went to a dry pot at this temperature.

As the results of the distillation show, the reaction of tert- butoxide with methylene bromide gives only the expected formal; no other products except unreacted starting materials were recovered.

Preparation of Diisopropyl Formal. Diisopropyl formal was prepared in much the same manner as di(tert-butyl) formal except that methylene chlorobromide was used in the place of methylene bromide. The results of the distillation were also much the same in that only unreacted starting material and the expected diisopropyl formal were recovered.

The boiling point of the formal was 117°.

III 114

Determination of the Stoichiometry of the Reactions of Potassium tert-

Butoxide and Isopropoxide with the. Methylene Halides

Determination of, the stoichiometry of the reactions of potassium

isopropoxide and tert-butoxide with the methylene halides, the

ratio of alkoxide consumed to halide produced, was effected by adding a

25 ml. aliquot of a standard solution of the alkoxide to a 50- or 100-m1. volumetric flask followed by an aliquot, ca., 1 ml., of the methylene halide to be studied. The reactants were mixed, the flask stoppered,

and the mixture placed in a constant-temperature bath at 36°. After

50 percent reaction or more, the flask was removed from the bath and

diluted to volume with distilled water. This served to dissolve the salts that had precipitated during the course of the reaction and resulted in a homogeneous solution from which aliquots were taken for analysis of strong base and halide.

The strong base concentration was determined by titration with standard perchloric acid to a phenolphthalein endpoint. The halide con-

centration was determined by potentiometric titration with standard silver nitrate. From the changes in the concentration of the base and the halide, the ratio, A[R0 - ]/ ,6[X- ], was calculated.

Deuterium Exchange of the Methylene Halides

The Nuclear Magnetic Resonance Method

Nuclear magnetic resonance (nmr) spectroscopy was the first analytical tool investigated for following the rates of deuterium exchange of_the methylene halides. This technique has been employed by several workers (31, 32) to ascertain both the position and the rate of deuterium 115

exchange. Ideally, this method can be applied by sealing the reactants in an nmr tube and periodically recording and integrating the spectrum.

The exchangeable protons in a given reaction mixture are constant. If the deuterium exchange of the methylene halides is studied in deuteroalcohols, the exchangeable protons of the parent compound are subsequently distri- buted between the solvent and the formal produced. By obtaining the areas of the absorptions from the protons of the methylene halide, the methylene group of the formal and the hydroxyl group of the solvent, the fraction of protium in each form can be determined and the rate constants for deuterium exchange calculated. Alternatively, the rate of deuterium exchange can be determined by following either the decrease in the absorption of the hydrogens of the methylene halide or the increase in the absorption of the hydroxyl group of the alcoholic solvent. An attempt was made to determine the rate of deuterium exchange of methylene bromide in methyl alcohol-d by nmr spectroscopy. The results are shown in

Figure 1, Appendix.

The spectrum of methyl alcohol-d is shown in Figure la. To obtain measurable signals from the protons of the methylene halides at concentrations between 0.5 and 1.0 H, it was necessary to record the spectrum at high amplitude. Under these conditions, the methyl hydrogens of methyl alcohol-d are completely off scale. The absorption at 5.58 T is due to the protium content of methyl alcohol-d, i.e., the 0-H absorption.

The spectrum is further complicated (because of the high amplitude) by 13 H-C absorptions at 5.85 and 8.21 T and spinning side bands at 6.12,

6.58, 7.48, and 7.92 T.

1 116

The addition of methylal to the alcohol is shown in Figure lb.

The methyl-group absorptions of methylal are coincident with those of the

methyl group of methyl alcohol and are thus obscured. The absorption of

the hydrogens of the methylene group of the formal at 5.82 T is almost 13 superimposed on the H-C satelite peak at 5.85 T.

Addition of methylene bromide to the mixture is shown in Figure lc.

The absorption of the, hydrogens of the methylene halide at 5.20 T is

clearly defined and suggests that a method for following the rate of

deuterium exchange might be developed based on this signal; however, the

probability of obtaining accurate values for the fraction of protium in both the reactants and products appears unlikely due to the inability to

separate the absorptions of the methylene group of the formal from that 13 of the H-C satelite of the solvent.

Addition of sodium methoxide to the reaction mixture is shown in

Figure ld. Examination of the spectrum shows that the absorption due to

the 0-H group of methyl alcohol is shifted downfield to 4.2 T. The fact

that this is indeed the 0-H absorption was verified by adding more methylene bromide to the reaction mixture, as shown in Figure le. It will be noted that addition of the methylene halide enhances the absorption

at 5.2 T due to the hydrogens of the halide. Furthermore, a shift in the

0-H absorption from 4.2 to 4.4 T is also effected. It is evident,

therefore, that the 0-H absorption is dependent upon the methoxide con-

centration.

In actual attempts to follow the rate of deuterium exchange by nmr spectroscopy, the absorptions due to the hydrogens of methylene bromide and the 0-H group of methyl alcohol were even closer together 117

initially than shown in Figure ld because lower methoxide concentrations

were employed. As a result, the spectrum was difficult to integrate even

when substantially narrower sweep widths were employed. Furthermore, as

the spectra indicate, the 0-H absorption moves progressively upfield as

the reaction proceeds, i.e., as the methoxide is consumed; consequently,

the two absorptions first move, together and then apart and at times are

coincident which renders quantitative analysis impossible.

The only alternate method of following the rate of deuterium

exchange by nmr spectroscopy was to quench an aliquot of the reaction

mixture, isolate the unreacted methylene halide, and subsequently

determine its deuterium content. This method could probably have been

employed but would undoubtedly have been quite time-consuming. Fortun-

ately, a much simpler method presented itself, was developed, and

subsequently employed.

The Near-Infrared Method

The use of the 0-H overtone absorptions in the near-infrared

region was also investigated as a possible method for determining the

rates of deuterium exchange of the methylene halides. The advantages of this technique as compared to employing the infrared absorptions

(Amax 2.9 to 3.0 u) lay in the longer cell-path lengths available and the inertness of the cells (silica) to the basic alcoholic solvents employed.

A study of the absorption spectra of methyl alcohol and methyl

alcohol-d (10-mm. cell) in the 1.0 to 2.0 u region showed that methyl

alcohol had an absorption maximum at 1.58 u while methyl alcohol-d

exhibited only base-line absorption in this region. Methylene bromide

also showed only negligible absorption at this wavelength. Determination 118

of the extinction coefficient of the 0-H absorption at 1.58 u showed that it was only 1.4 to 1.7. This was much too low to make the method useful as an analytical tool; consequently, work with this technique was abandoned.

The Infrared Method

Two basic approaches employing infrared spectroscopy were considered for following the rates of deuterium exchange of the methylene halides. The first involved isolating the unreacted methylene halide and ascertaining the extent of deuteration. Hine and co-workers (4-12) employed this method in determining the rates of deuterium exchange of the haloforms. Although a simple extension of their technique appears to be applicable to the study of the rates of deuterium exchange of the methylene halides, it nevertheless proved too complex to pursue. The problem with this method lies in the fact that the exchange of the methylene halides yields two products, CHDX and CD2 X2, rather than just 2 one as in the case of the haloforms. This requires that CHDX be 2 independently synthesized and that a method be devised for quantitatively determining both CHDX 2 and CD2 X2 simultaneously. The synthesis of the dideuteromethylene halides is feasible either by direct exchange of the methylene halide (18) or by reaction of the corresponding deuterohaloform

(DCX ) with sodium arsenate in deuterium oxide (33). Dideuteromethylene 3 bromide was prepared by this technique in excellent yields (34). Unfor- tunately, an unequivocal synthesis of the monodeuteromethylene halides was not readily available norh could one be devised. Furthermore, even if

CHDX and were available, there was no assurance that a satisfactory 2 CD2X2 analytical method based on infrared spectroscopy could be developed.

Since the synthetic aspects of this analytical approach were so complex, 119

the method was not pursued.

The second analytical technique (employing infrared spectroscopy) considered for ascertaining the rates of deuterium exchange of the methylene halides involved following the rate of appearance of protium in the solvent. This method proved to be both convenient and reliable and was employed to ascertain the rates of deuterium exchange of all of the methylene halides studied.

Absorption Spectra of the Reactants and Products. A study of the infrared absorption spectra of methyl, isopropyl, and tert-butyl alcohols in the corresponding deuteroalcohols showed that the 0-H absorption

(x 2.9 to 3.0 p) followed Beer's law with extinction coefficients max ranging from 135 to 175. The 0•H absorption was easily separated from the 0-D absorption between 4.0 and 4.2 p and the C-H absorptions in

3.1 to 3.5 p region. Furthermore, a study of the absorption spectra of the methylene halides (35a) and the corresponding formals (35b),

Figures 2 and 3, Appendix, formed in the course of the reactions showed that at the concentrations employed, these substances exhibited only background absorption in the 0-H absorption region; consequently, no corrections for ether the reactants or products were necessary.

The Intensity of the 0-H Absorption. The intensity of the 0-H absorption was determined by recording the spectrum of the reaction mixture between. 2.4 and 3.2 p (with air in the reference beam) on a

Perkin-Elmer, Model 21, recording spectrophotometer using Perkin-Elmer semilog paper (No. 021-6304). The intensity of the 0-H absorption

(absorbance) was determined by taking the difference (in log units)

E I, 120

between the minimum in the absorption spectrum between 2.65 and 2.75 p

and the maximum between 2.9 and 3.0 p as illustrated in Figure 4. This

method of determining the absorbance is known as the empirical ratio

method (36). Calculating the absorbance by the base-line technique (36)

did not give a linear correspondence with Beer's law. Alternatively, the

technique of recording differences in optical density at a fixed wavelength

was not employed because this method depends upon the characteristics of

the cell remaining constant throughout a given set of measurements. In

practice, the cells gradually grew opaque (especially when employing potassium isopropoxide or potassium tert-butoxide solutions). In addition,

the cell-path length also increased due to the solubility of the sodium

chloride windows of the cells in the alcoholic solvents.

The increasing opaqueness of the cells was internally compensated

for by the method of determining the absorbance; i.e., as the opaqueness

increased, the base-line measurement between 2.65 and 2.75 p and the 0-H maximum between 2.9 and 3.0 p increased correspondingly. Since the

absorbance was taken as the difference between these two values, this

variable tended to cancel itself. Unfortunately, no internal method for

compensating for the changing cell-path length could be devised; conse-

quently, it was necessary to periodically determine this parameter.

Variations in the . Cell-Path Length. As previously mentioned, periodic monitoring of the cell-path length was necessary due to the

solubility of the sodium chloride windows in the alcoholic solvents.

When isopropyl and tert-butyl alcohols were employed, determination of

the cell-path length was necessary only at the beginning and at the end of a run. These values were usually within experimental error and an 121

average of the two was employed to calculate the ROH concentrations. On the other hand, due to the greater solubility of sodium chloride in methyl alcohol, studies in methyl alcohol-d required more frequent monitoring of this variable. Indeed, when methyl alcohol was employed, it was found that, during the course of a single, deuterium-exchange study in which 10 to 15 individual measurements were made, the cell-path length often increased by 0.005 to 0.010 mm. Under these circumstances, the cell-path length was measured several times in the course of the study. In all cases, determination of the cell-path length was made at the beginning and at the end of a deuterium-exchange study.

Determination of the Cell-Path Length. To determine the cell-path length (37), the infrared cell was filled with spectral grade benzene and its spectrum recorded between 4.5 and 5.5 p with air in the reference beam. The optical density of the absorption at 5.1 p was then measured by the base-line technique and the cell-path length calculated from equation (3).

Cell-Path Length (mm.) = 0.100 x Absorbance (3)

The instrument settings employed were as follows: Resolution, 990;

Gain, 4; Response 2; Suppression, 5.

It was recognized that the values of the cell-path length obtained by this procedure were not as exact as those obtained by the interference method (37); however, the extinction coefficients and all of the subsequent analyses are based on cell-path lengths determined by this technique. The entire analytical scheme, therefore, is internally

122

consistent.

Determination of the Extinction Coefficients for ROH in ROD. The extinction coefficients for ROH in ROD were determined by first recording the absorption spectrum of the deuteroalcohol to be studied between 2.4 and 3.2 p with air in the reference beam. This gave a measure of the residual ROH content of the ROD. A known amount of ROH was then added to the ROD and the spectrum recorded again. The difference in the intensity of the two absorptions was then ascertained and the extinction coefficient calculated from equation (4):

E = A lc (4) where A is the absorbance, i.e., optical density, 1 is the cell-path length in centimeters, and c is the molar concentration of added ROH.

The instrument settings employed were as follows: Resolution, 990;

Gain, 4; Response 2; Suppression, 5. The results of these studies are shown in Tables 10, 11, and 12, Appendix.

Dependence of the Extinction Coefficients Upon the. Alkoxide

Concentration. It was also found that the extinction coefficients were dependent upon the alkoxide concentration; the coefficients decreasing with increasing base concentration. As a result, the extinction coefficients were determined at various alkoxide concentrations and a plot of e versus the alkoxide concentration for each alcohol employed constructed. The results of these studies are shown in Figures 5, 6, and 7, Appendix. With these plots and a knowledge of how the base concentration is changing throughout the course of a given deuterium exchange study, the appropriate value of the extinction coefficient 123

can be obtained.

Determination of the Rates of Deuterium Exchange of the Methylene

Halides. The methylene halide to be studied was weighed into a dry,

nitrogen-purged volumetric flask and dissolved in the alcoholic solvent

to be employed. An aliquot of a standard solution of the conjugate base

of the solvent (as either the sodium or potassium salt) in the corres-

ponding deuteroalcohol was then added and the mixture diluted to volume.

After thorough mixing, an aliquot of the reaction mixture was withdrawn

by pipet and transferred to a l00-ml. beaker containing about 20 ml. of

distilled water and subsequently titrated with standard perchloric acid to

a phenolphthalein endpoint. This gave a measure of the initial alkoxide

concentration. Another sample was withdrawn by means of a dry, nitrogen-

filled syringe. After rinsing the syringe several times with small

portions of the reaction mixture, a sample was transferred to an infrared

cell and the spectrum recorded between 2.4 and 3.2 11.

The remainder of the reaction mixture was transferred to a dry,

nitrogen-filled serum bottle, capped, and placed in a constant-temperature

bath. Periodically, samples were withdrawn by means of nitrogen-filled

syringes to follow the deuterium content of the solvent by the infrared

technique and the alkoxide concentration by titration with standard

perchloric acid.

As in the study of the rates of formal formation, the initial

time in the deuterium exchange studies was recorded when the serum

bottles were placed in the constant-temperature bath. With practice,

the entire procedure for preparing the reaction mixtures, including

taking the samples for analysis, required only about three minutes.

I ii 124

Since the preparation of the solutions was carried out at room temperature, ca 25°, and since it took a comparable amount of time for the reaction mixtures to come to reaction temperature, i.e., 36°, recording

of the , initial time when the serum bottles were placed in the constant- temperature bath seemed to be the best reference point. Unfortunately, the reaction times in the deuterium exchange studies were not as long as those observed in the study of the rates of formal formation and in some cases an error of two to three minutes constitutes an error of greater than one percent.

One of the shortcomings of the infrared method is that the con- centratdon of the methylene halide and the cell-path length cannot be varied substantially. In order to have enough protium present (in the methylene halide) to give changes in the optical density of 0.5 to 0.8 units (before the reaction reaches 40 percent completion), the initial concentration of the methylene halide must be about 0.5 M. Longer cell- path lengths cannot be employed because the 0-H absorption (due to the initial protium content of the solvent) would be off-scale. Therefore, as progressively stronger bases are employed and the rate constants for deuterium exchange increase, the only variable that can be manipulated is the alkoxide concentration.

Varying the alkoxide concentration in order to obtain reasonable rates of deuterium exchange in methyl alcohol-d and isopropyl alcohol-d proved to be no particular problem from an analytical point of view although it took some probing to ascertain optimum concentrations to obtain reasonable rates. On the other hand, the rate constants for

deuterium exchange in tert-butyl alcohol-d were so large, ca., k Ai 0.1, 125

-3 that concentrations of potassium tert-butoxide of, the order of 10 to -4 10 M were necessitated. At these concentrations, it proved difficult to obtain reliable values of the alkoxide concentration.

The problem lies in the fact that, at such low concentrations, the base is largely converted to carbonate in the course of performing the necessary transfers prior to titration. This led to erratic results. We experimented for some time with not only protecting the reaction mixture from air, as was customarily done, but with transferring and titrating the reaction mixtures in a nitrogen atmosphere using nitrogen-purged, carbonate-free water for preparing all of the standard

solutions and , as solvent for the titrations. However, even with the most exacting care, erratic results were still obtained.

As a solution to the problem, it was found that by allowing the reaction mixtures (after dilution with water) to stand in the atmosphere until the alkoxide (i.e., hydrokide) was completely converted to carbonate and then analyzing for carbonate rather than hydroXidelp(py potentiometric titration of the samples with standard perchloric acid) gave consistent sets of values which could be employed to detarmine the rates of deuterium exchange. As a result, this basic procedure was employed

in the, studies of the deuterium exchange of the methylene halides in tert-butyl alcohol-d.

The procedure adopted was to take an aliquot of the reaction mixture in the usual manner with a dry, nitrogen-filled syringe and to dissolve an aliquot in about 20 ml. of distilled water. The mixture was then allowed to stand overnight to assure, complete conversion to carbonate. On the following day, the samples were titrated 126

potentiometrically with standard perchloric acid and the endpoints determined graphically (38).

It is recognized that this method has shortcomings which in turn lead to uncertainties as to the absolute values of the rate constants.

Furthermore, it is impossible to ascertain how much of the base in the reaction mixture is there as alkoxide and how much is there as carbonate.

Nevertheless, the theoretical approach is sound and the fact that a literature search revealed no procedure for the titration of alkoxides in this concentration range suggests that this may present an analytical problem of some magnitude.

Determination of the Deuterium Content of Exchanged Methylal

Preparation of Deuterium-Exchanged Methylal

Methylene bromide, 12.50 g. (0.0719 moles), was added to a mixture of 25 ml. of methyl alcohol-d and 25 ml. of 2.314 M sodium methoxide

(in methyl alcohol-d) in a 60-m1. serum bottle. The bottle was stoppered, the reactants mixed, and the bottle placed in a constant-temperature bath at 36°. After 5 to 15 percent reaction, the reaction mixture was removed from the constant-temperature bath and immersed in an ice bath.

After cooling, it was poured into a separatory funnel containing an ice- water slurry. The mixture was then extracted four times with 4 to 5 ml. portions of carbon tetrachloride. The extracts were combined and dried over calcium chloride for several days while the sample was stored in a refrigerator to prevent loss of the volatile methylal (b.p. 44°).

Isolation of Deuterium-Exchanged Methylal

Separation and isolation of the partially-exchanged methylal from the combined extracts were effected by preparative gas-liquid 127

chromatography using a Wilkens Model A-700, Autoprep, and a 4' x 3/8"

Apiezon column. A two-way syringe valire (B-D, MS08) was connected to the

outlet of the gas chromatograph by a short piece of tubing. One arm of

the valve was connected to a 10" hypodermic needle which was inserted

into a small-bore nmr tube; the latter was immersed in a dry-ice bath.

The other arm, by means of a piece of tubing, led to a refuse bottle.

Attempts to analyze the dilute solutions of methylal in carbon

tetrachloride prior to separation by gas chromatography proved unsatis-

factory. Not only was there interference from the proton signals of the

unreacted methylene bromide but the nmr spectrum had to be recorded at

such high amplitude that reproducible results were difficult to obtain.

Procedure. Samples of the dried, combined extracts, ca 0.25 to

0.50 ml., were charged to the gas chromattograph at 30° to 40°. The

methylal and residual methanol were separated at this temperature. In normal operation, the effluent gases from the detector of the gas

chromatograph were directed into the refuse bottle. As the methylal

began to elute from the column, the two-way valve Was turned so that the

exit gases were directed into the nmr tube where the methylal was con-

densed. As soon as the elution of the methylal was complete, the two-way

valve was reversed and the temperature of the column raised as rapidly as

possible to 100° to 120° in order to facilitate removal of the carbon

tetrachloride and methylene bromide from the column. No attempt was made

to isolate this fraction. By processing the entire 15 to 20 ml. of

combined extracts, 0.2 to 0.5 ml. of the partially exchanged methylal

could be recovered. In the small-bore, nmr tubes, this was sufficient for nmr analysis. After complete processing of the combined extracts, the, 128

nmr tubes were sealed to prevent loss of the methylal.

In the course of developing the separation technique, it was found that if the chromatographic separation was attempted on extracts which had been dried over calcium chloride for only a few hours, small but measurable amounts of methanol could be detected; however, upon drying for 24 to 48 hours, the methanol peak was absent. Since the elution times of methanol and methylal were close together, it was advantageous to allow the methyl alcohol to be removed by the calcium chloride treatment prior to chromatographic separation; consequently, extended drying periods were employed in both of the reported studies and no methyl alcohol was observed in the chromatograms.

Determination of the CH /CH Ratio of Deuterium-Exchanged Methylal 2 3 The nmr spectrum of the partially-exchanged methylal was recorded on a Varian Associates iModel A-60 Spectrometer at a spectrum amplitude so that the two peaks (singlets for both the methyl and, methylene hydrogens) were both on the scale of the chart paper. The resulting spectrum was then integrated ten or more times and the ratio of methylene to methyl protons determined for each integration. The average ratio and average deviation were then calculated. Immediately following the analysis of the exchanged methylal, a sample of pure methylal (purified by the same gas-liquid chromatography procedure) was analyzed at the same instrument settings that had been previously employed to determine the ratio of methyl to methylene protons for the partially-exchanged methylal. The spectrum of pure methylal was also integrated ten or: more times and the ratio of methylene to methyl protons determined.

Although this ratio should theoretically be 1:3, it was found to be 129

slightly higher than this value. This experimentally-determined value

was used to correct the ratio of methylene to methyl protons in the

partially-exchanged methylal to the theoretical value.

Preparation of Methyl Alcohol-d

Methyl alcohol-d was prepared by the hydrolysis of methyl borate

with deuterium oxide.

B(OCH + 3D 0 -÷ 3CH OD + B(OD) 3 ) 3 2 3 3

Procedure

To a dry, nitrogen-purged, 5000-ml., 3-necked flask, equipped with

a stirrer, addition funnel, and condenser (protected from the atmosphere

with a drying tube filled with Drierite) were added 2300 g. (22.1 moles)

of methyl borate and 695 g. (6.67 moles) of anhydrous sodium carbonate.

The reaction mixture was brought to reflux followed by the addition of

deuterium oxide in 10-m1. increments. After addition of about 40 ml.

of deuterium oxide, the reaction initiated. It is important that the

reaction be initiated with a minimum amount of deuterium oxide due to the highly exothermic nature of the hydrolysis reaction and the extreme

volatility of methyl borate; otherwise, it is impossible to control the

reflux when the reaction initiates.

After the initial exotherm had subsided, the remainder of the

deuterium oxide, 1271 g. (63.5 moles), was added over a one-hour period

at a rate sufficient, to keep the reaction mixture at a gentle reflux.

As the reaction proceeds, sodium borate precipitates from the reaction

mixture and markedly impedes mixing. After several hours, the solids 130 content of the flask is usually so high that stirring becomes impossible; however, a gentle reflux was maintained for 16 hours after the addition of the deuterium oxide was complete.

A Claisen head was then attached to the flask and the crude methyl alcohol-d distilled collecting over 2000 ml., b.p., 60° to 90b.. An additional 326 g. of a material (largely D 20), b.p., 90° to 105°, was also recovered.

An aliquot of the crude methyl alcohol-d was analyzed for unreacted methyl borate by treating an aliquot with standard sodium hydroxide in the presence of mannitol and then back titrating with standard perchloric acid. The results showed that the crude alcohol contained 0.113 M methyl borate.

The crude methyl alcohol-d was treated with 20 g. of sodium, 20 ml. of deuterium oxide, and then the mixture was subsequently fractionated on a five-foot column packed with glass helices at a 10:1 reflux ratio, collecting between 1800 and 1900 ml. of methyl alcohol-d, b.p. 64.5°.

An aliquot of the redistilled alcohol was again analyzed for -4 methyl borate and found to contain less than 3 x 10 M of the borate ester. Analysis of the alcohol by the infrared method showed that it contained 0.174 M methyl alcohol which was sufficiently pure for our purposes.

Preparation of Isopropyl and.tert-Butyl Alcohol-d

Isopropyl and tert-butyl alcohol-d were prepared by direct exchange of the protioalcohols.

2R0H + D 0 2ROD + H 2O 2 131

Procedure

About 1000 ml. of a dry, commercial-grade alcohol was mixed with about 500 ml. of deuterium oxide in a dry, nitrogen-filled flask. After several hours, potassium carbonate was added to salt out the alcohol.

Addition of potassium carbonate was continued until the aqueous layer was saturated after which the mixture was transferred to a separatory funnel, and the aqueous layer separated. The aqueous phase was subsequently distilled to recover the partially-exchanged deuterium oxide and this was used as the initial exchange medium for subsequent preparations of the deuteroalcohols.

The alcoholic phase was mixed again with deuterium oxide, ca., 99 percent, and the entire procedure repeated. The exchange process was repeated until the protioalcohol content of the deuteroalcohol was between 0.1 and 0.3 M; this usually took four exchanges. The entire batch of the crude deuteroalcohol was then dried over anhydrous Drierite.

Purification. Isopropyl alcohol-d was purified by fractionation on a five-foot column packed with glass helices at a 10:1 reflux ratio.

The entire distillation assembly was dried and purged with nitrogen prior to use and the fractionation was carried out in a system protected from the atmosphere by a drying tube qontaining Drierite. The boiling point of the alcohol was 82.0° (uncorrected).

tert-Butyl alcohol-d was purified by simple distillation through a Claisen head after reacting it with several grams of potassium to insure that it was anhydrous. This procedure was repeated until the freezing point of the distillate was constant, i.e., 25.35°; this seldom took more than two distillations. The boiling point of the pure deuteroalcohol was 132

82.1 to 82.5 ° .

Yields. The yields in the preparation of the deuteroalcohols by the direct exchange method varied between 60 and 75 percent based oh the starting alcohol; the major losses occurred in the salting-out step. 133

CHAPTER III

DISCUSSION AND RESULTS

Reactions of the Methylene Halides with Alkoxides

The Rates of Formal Formation

The rate constants for the reactions of the methylene halides

with sodium methoxide and potassium isopropoxide and tert-butoxide (in

the corresponding protio- and deuteroalcohols) are shown in Table 2

(Tables 13 through 39, Appendix).

2R0- + -0- (RO)- CH + 2X- CH2' X2 2 2

The rate constants were calculated from the second-order rate

equation (5):

.2.303 b(a-x) k - to g (2a-YTT a(b-2x) (5) where a and b are the initial concentrations of the methylene halide and

alkoxide, respectively, and x is the concentration of formal as calculated

from equation (6).

x = ([R0 - ], - [R0 - ])/2 (6)

Hine, Thorws, and Ehrenson (13) have previously demonstrated that equa-

tion (5) is the rate equation required by the reaction of sodium methoxide with the methylene halides. Table 2. Rate Constants for the Reaction of the Methylene Halides frith Alkoxides in Alcoholic Solvents at 36°

Methylene Halide CH I CH Br CH BrC1 CH C1 2 2 2 2 2 2 2 6 a 6 a 6 a 6 a 10 k 10 k 10 k 10 k -1 -1 -1 -1 -1 -1 -1 -1 Solvent R. mole sec. t. mole sec. R. mole sec. R. mole sec. b b b c Me0H 1.86 1.23 4.02 0.123

Me0D 2.43 ± 0.09 2.28 ± 0.03 6.89 ± 0.05 0.185 ± 0.005 2.71 ± 0.19 2.43 ± 0.04 2.40 ± 0.07 9 .53 ± 0.08 2.54 ± 0.14 iso-PrOH 6.90 ± 0.12 5.49 ± 0.03 19.2 ± 0.02 0.396 ± 0.026 6.77 ± 0.12 iso-PrOD 10.2 ± 0.1 8.17 ± 0.10 31.2 ± 0.5 0.593 ± 0.029 tert-BuOH 2.30 ± 0.03 2.83 ± 0.03 11.8 ± 0.1 0.241 ± 0.006 2.67 ± 0.04 tert-BuOD 3.97 ± 0.02 4.79 ± 0.12 19.7 ± 0.6 0.458 ± 0.057 a b c Calculated from equation (5). Calculated from reference (13). Reference (13).

135

Equation (5) may be employed to calculate the rate constants

irrespective of whether the formal is produced by a nucleophilic

displacement process or an a-elimination mechanism. If the process

is S N 2 in character as shown in mechanism (7), the basic assumption is that the rate-controlling step of the reaction is the initial nucleophilic

displacement of halogen from the methylene halide by alkoxide. The

resulting alkoxyhalomethane is then rapidly transformed into the corres-

ponding formal, i.e., k 2 > kl .

RO - + CH X —0- ROCH X + X- (7) 2 2 2

RO + ROCH X (RO) CH + X- 2 2 2

If the transformation is proceeding by an a-elimination mechanism as

shown in equation (8), then the basic assumption is that the transfor-

mation of the dihalomethyl carbanion to a monohalomethylene is the rate-

controlling step of the reaction, i.e., k > k 1 2.

RO - + CHX ROH + CHX- (8) 2 2 2 -1

k 2 CHX2- -4— CHX + X- -2

RO - + ROH(D) + CHX (RO) CH (D) + 2 2 136

Hine and co-workers (4-12) have shown that the rate-controlling step in the basic hydrolysis of the haloforms is the transformation of the trihalomethyl anion to a dihalomethylene. However, equation (5) is equally valid if the rate-controlling step of the a-elimination mechanism is formation of the dihalomethyl carbanion.

As the results show, the order of reactivity of the alkoxides toward the methylene halides is isopropoxide > tert-butoxide > methoxide.

Although there appears to be no definitive work establishing the relative nucleophilicities of these alkoxides, it was assumed that their nucleophilicities would parallel their basicities, i.e., tert-butoxide > isopropoxide > methoxide. The observed order of reactivity may be explained by assuming the expected order of nucleophilicity and envoking steric effects to account for the reversal between tert-butoxide and isopropoxide.

The relative reactivities of methoxide, isopropoxide, and tert- butoxide toward the methylene halides are shown, in Table 3. The results show that there is less than a fivefold difference between the rate constants for the reactidn of isopropoxide and methoxide (tert-butoxide falling in between) with any given halide. As previously discussed, the basicities of the alkoxides differ by several orders of magnitude and the order of decreasing strength is tert-butoxide > isopropoxide > methoxide.

If the reactions were proceeding by an a-elimination mechanism, then the order of reactivity would be expected to parallel the basicities of the alkoxides. The small differences in the reactivity of the alkoxides with the methylene halides and the unexpected reactivity pattern, i.e.,isopropoxide > tert-butoxide > methoxide, strongly suggest that the reaction is largely SN2 in character. 137

Table 3. Relative Reactivities of Several Methylene Halides with Methoxide, Isopropoxide, and tert-Butoxide

Alkoxide Methylene Halide (Solvent) Me0Na iso-PrOK tert-BuOK

CH 3.78 1.24 2 I 2 (ROH) 1 CH 1 4.02 1.56 2 2 (ROD) 1 CH 4.46 2.24 2Br2 (ROH) 1 CH 2.05 2Br 2 (ROD) 1 3.49 CH BrC1 (ROH) 4.78 2.94 2 1 CH 3.86 2BrC1 (ROD) 1 4.53 CH2C12 (ROH) 1 3.22 1.96 CH 1 3.20 2.48 2C12 (ROD)

Reaction of Chloroform with Alkoxides

The Rate of Reaction

To establish the order of reactivity of the alkoxides toward a

substrate known to react by an a-elimination mechanism, the rate constants

for the reaction of the alkoxides with chloroform were determined. The

results are shown in Table 4 ( Tables 40 and 41, Appendix).

The rate constants were calculated from equation (9):

2.303 k = b(a-x) 1°-log 3a-} )t 7E757 (9)

where a and b are the initial concentrations of chloroform and alkoxide,

respectively, and x is the amount of orthoformate [or other reaction products of similar stoichiometry given in equation (11)] as calculated 138

from equation (10).

- [R0 - ] x - (10) 3

Table 4. Reaction of Chloroform with Methoxide, Isopropoxide,and tert-Butoxide in the Corresponding Alcohols at 36°

ka -1 -1 Alkoxide Z. mole sec -6 Methoxide 4.35 ± 0.11 x 10 3 Isopropoxide 3.10 ± 0.03 x 10 -1 tert-Butoxide 1.0 ±10 x 10 aCalculated from equation (9).

IUD basic assumptions are necessary to derive equation (9). The first is that the rate-controlling step of the reaction is formation of dichloromethylene and, the second is that the stoichiometry of the reaction demands three moles of alkoxide per mole of haloform. The former is consistent with the earlier work of Hine, et al., (4-12) and the latter with any one of the product-determining reaction paths shown in equation (11).

ROH ) (RO) CH + 2C1 - 3 CC1 + 2R0 - 2 ) R (olefin) + 2C1 - + CO + ROH ) ROR + 2C1 - + CO 139

If the assumptions regarding the stoichiometry of the reaction were grossly in error, calculation of the rate constants by equation (9) would produce trends in the observed rate constants reflecting the faulty choice. As the results show (Tables 40 and 41, Appendix), the rate constants for the reaction of chloroform with sodium methoxide and potassium isopropoxide show no such trends.

It was possible to determine the rate constants for the reactions of sodium methoxide and potassium isopropoxide with chloroform; however, the reaction of potassium tert-butoxide with the haloform proved too fast to study. A crude estimate of the rate constant for this reaction, however, can be obtained from the fact that at initial potassium tert- -3 -2 butoxide and chloroform concentrations of 2.51 x 10 M and 2.03 x 10 M, -3 respectively, the base concentration fell to 0.07 x 10 M iin 21 minutes. -1 1 From these data, a rate constant of 0.1 to 0.2 Q. mole sec. can be estimated; however, this may be in error by as much as an order of magnitude.

Despite the uncertainties in the actual value of the rate constant for the reaction of tert-butoxide with chloroform, it is evident that the

order of reactivity of the alkoxides toward , the haloform is tert-butoxide

> isopropoxide > methoxide. This is the order expected from the basicities of the alkoxides. These results verify our previous assumption that the rate of a reaction proceeding by an a-elimination mechanism should increase, with increasing strength of the base and strongly suggest that the reactions of the methylene halides with the same alkoxides are proceeding largely, if not entirely, by a nucleophilic displacement process. 140

Reactions of the Methylene Halides with Alkoxides

The Solvent Kinetic Isotope Effect

A surprising result of the study of the reactions of the methylene halides with alkoxides in the corresponding protio- and deuteroalcohols is the unusually large solvent kinetic-isotope effect (32b). As shown in

Table 5, a given methylene halide reacts between 1.37 and 1.91 times as fast in ROD as in ROH; the average value of is 1.68. kROD/kROH Although the isotope effect can be rationalized, it seems unlikely that either its magnitude or direction (i.e., greater or less than unity) could have been predicted in advance. In regard to direction, Hine and co-workers (4-12) have shown that for most of the haloforms (a kD/kH reaction known to proceed by an a-elimination mechanism) is about 0.57.

If the methylene halides were reacting largely by the same mechanism, then a fractional value for kROD/kROH would be expected because, in most cases, the rate of deuterium exchange of the methylene halides runs well ahead of the rate of formal formation. As a result, the equilibrium quantity of deuterium is incorporated into the halide before any significant amount of formal has been produced; the rates of formal formation, therefore are essentially those of the dideuteromethylene halides. 141

Table 5. The Solvent Kinetic Isotope Effect in the Reactions of the Methylene Halides with Alkoxides

Methylene k /k /k k /k Halide Me0D Me0H k PrOD PrOH BuOD BuOH

CH I 1.37 1.49 1.73 2 2 CH Br 1.91 1.49 1.74 2 2 CH BrC1 1.72 1.63 1.67 2 CH C1 1.51 1.50 1.90 2 2

a - Deuterium Isotope Effects. The role of deuterium substitution in the a-position to the reaction center has been reviewed by

Streitwieser (23). For reactions such as the acetolysis of tosylate esters, where bond breaking runs ahead of bond making in the transition state (i.e., the reactions that are largely S N1 in character), kH/kD

(per deuterium) ranges from 1.10 to 1.19 (24, 25). On the other hand, for reactions which are largely S N 2 in character, such as the hydrolysis of methyl-d halides (26), the values of kH/kD range from 0.98 to 1.00

(per deuterium). The reactions of the alkoxides with the methylene halides would probably best fit in the latter category and little or no effect on the rate constants for formal formation would be expected by replacing the hydrogens of the halide with deuterium. Furthermore, it will be noted that in the previously cited examples the reactions proceed faster with protium rather than deuterium in the a-position. Our results showed that the reactions of the alkoxides with the methylene halides proceeded faster in ROD than in ROH or in cases where deuterium was in 142

the a position (due to exchange). The contrast between our observations

and those reviewed by Streitwieser suggests that the observed rate

acceleration in ROD is not due to an a-deuterium isotope effect.

Hydrogen-Bonding Effect. The most likely explanation for the pronounced solvent kinetic-isotope effect lies in the relative abilities of ROH and ROD to stabilize the alkoxide ion by hydrogen bonding (5b).

Bellamy and Rogasch (27) and Dahlgram and Long (28) have reported that

deuterium bonds are comparatively weaker than hydrogen bonds. Literature

data (29) on self-association also show lower values of equilibrium constants and enthalpies for deuterium bonds. In contrast, Singh and

Rao (30) have shown that the ratio of the equilibrium constants, K H /KD , vary from donor to donor and suggest that K H/KD is a function of the basicity of the donor. Their values of the ratios of equilibrium constants varied from 0.2 to 4.3. However, if it is assumed that ROH is a better hydrogen bonding solvent than ROD, then it follows that an alkoxide is better solvated in ROH than in ROD. As a result, the alkoxide is more

stable hence less reactive in ROH as compared to ROD.

The Rates of Deuterium Exchange

The rate constants for deuterium exchange of the methylene halides in methyl alcohol-d, isopropyl alcohol-d, and tert-butyl alcohol-d catalyzed by the respective conjugate bases of the alcohols are shown in

Table 6 (Table 42 through 55, Appendix). The rate constants in isopropyl alcohol-d and tert-butyl alcohol-d were calculated from equation (12),

(Derivation 1, Case 2, Appendix): 143

Table 6. Rate Constants a for Deuterium Exchange of the Methylene Halides

SolVent Me0D iso-PrOD tert-BuOD 3 10 6k 304k 10 k Methylene -1 -1 -1 -1 -1 -1 Halide 32. mole sec. k. mole sec. Z. mole sec.

CH Cl too slow 0.287 ± 0.15 1.06 ± 0.03 2 2 CH BrC1 ± 0.37 6.82 ± 0.45 2 too slow 3.75 7.55 ± 0.62

CH Br 5.95 ± 0.37 2 2 16.2 ± 0.3 28.0 ± 1.7 6.36 ± 0.10 33.8 ± 1.0 34.5 ± 1.8

CH I 3.62 ± 0.09 11.2 ± 0.3 25.2 ± 0.8 2 2 aValues of k were determined between one and 35 percent reaction based on exchangeable protium in the methylene halide. 144

2.303 b(a-p/2) k - log t(ac-b/2) a(b-cp) (12) where a and b are the initial concentrations of the methylene halide and base (alkoxide), respectively, p is the change in the protium content of the solvent at time t, and c is a proportionality constant relating the base concentration to p. Since the rate constants for deuterium exchange in these solvents are several orders of magnitude larger than the corresponding rate constants for formal formation, the amount of deuterium in the formal can be neglected.

The rate constants for deuterium exchange of the methylene halides in methyl alcohol-d were calculated from equation (13), (Derivation 1,

Case 1, Appendix):

.2.303 b(a-mp) k lo 2 , = Tia fbm)t a(b-cp) (13) where a, b, and c have the same designations as in equation (12) and m is a proportionality constant relating the deuterium content of the formal to p. The rate constants for deuterium exchange of a given methylene halide in methyl alcohol-d are roughly comparable to the corresponding rate constants for formal formation; consequently, the amount of deuterium in the formal must be taken into consideration in establishing the rate constants for deuterium exchange.

Derivation of the Rate Equations for Deuterium Exchange

The derivation of equation (12) and (13) include several assumptions which lead to simplification of the resulting rate expressions.

The validity of these assumptions is best evaluated by considering each 145

separately and in some detail.

The Mechanism of Formal Formation. The first assumption is that the formals are formed by a nucleophilic displacement process and that the rate-controlling step of this transformation is the initial attack of the alkoxide on the methylene halide. This assumption is necessary in order to derive equation (13) but is arbitrary in equation (12). Since the objective of this study was to ascertain whether the methylene halides were reacting by an a-elimination mechanism, it may seem rather biased at this point to introduce such an assumption into the derivation of the rate equations; however, the results of the studies establishing the rates of formal formation suggest that the reactions are at least largely S N 2 in character.

Neglect of Reversibility. Equations (12) and (13) also neglect reversibility (k_l and k_ 3 ), i.e., they assume that the carbanions,

CHX2 and CDX2 are invariably deuterated rather than protonated. In the ' detailed reaction mechanism (14) the simplifying assumptions are equiv- alent to assuming that k 2 [ROD] > k_l [ROH] and k4 [ROD] > k_ 3 [ROH].

CH X + RO - CHX2 + ROH 2 2 - 1

k 2 CHX- + ROD CHDX + RO - 2 -6-- 2 -2

k 3 CHDX + RO - --÷ CDX- + ROH 2 2 -3

k4 CDX- + ROD CD + RO 2 2 X2 -4 146

If the primary kinetic isotope effect (k H/k D ), i.e. k_ /k 2 and k_ 3 /k4 , 1 is small, ca., 1.0, then these assumptions are reasonable becaupe the ratio of deuteration to protonation is controlled by the relative concentrations of ROD and ROH. These assumptions are undoubtedly best in methyl alcohol-d which is about 25 M. The methyl alcohol-d employed in the deuterium exchange studies contained 0.2 to 0.3 M methyl alcohol.

During the course of a deuterium exchange study, the protium content of the solvent increased about 0.5 M; therefore, the concentration of methyl alcohol ranged from 0.2 to 0.8 M. Assuming that k H/kD is 1.0, the ratio of deuteration to protonation is between 30:1 and 80:1 throughout the entire study. In isopropyl alcohol-d and tert-butyl alcohol-d where the molarities are 13:1 and 10.5:1, respectively, the ratio of deuteration to protonation is lower; nevertheless, the approximations seem reasonable even in these systems. Only if the kinetic-isotope effect were large, ca., three or greater, would these assumptions require more detailed consideration.

Our derivations also neglect the reverse reactions defined by k_ 2 and k_14 . Experimentally, these assumptions can be favored if the rate constants are determined during the first 5 to 35 percent of reaction, i.e., while the methylene halide still contains most of its original exchangeable protium. Under these conditions, on a purely statistical basis, the alkoxide will be more apt to find protium than deuterium. Furthermore, the ratio of rate constants, k3 /k and k /k -4 1 -2' constitute a primary kinetic-isotope effect. Both values, therefore, would be expected to be equal to or greater than one. In conjunction with the previously discussed experimental restrictions, the primary 147

kinetic-isotope effect makes the decision to neglect the reactions defined by k_ 2 and k_14 quite plausible.

The Secondary Kinetic-Isotope Effect. The final assumption employed in the derivation of equations (12) and (13) is that of neglecting the secondary kinetic-isotope effect, i.e., assuming that k l = 2k 3 .

This assumption greatly simplifies the derivation of the resulting rate equations. Since secondary kinetic-isotope effects are usually small (5c),

ca., 30 percent at the most, this assumption seems to be justified.

Results of the Study of Deuterium Exchange

As the results (Table 6) show, the rate constants for deuterium exchange (i.e., carbanion formation) of a given methylene halide increase with increasing base strength, i.e., tert-butoxide > isopropoxide > methoxide. . Furthermore, the differences in reactivity between the respective bases is about that which would be predicted from the auto- protolysis constants of the alcohols (29). It will be noted that the differences in the rate constants for deuterium exchange of a given methylene halide are about the same as the differences observed in the rates of reaction of chloroform with the three alkoxides. The fact that the rate constants for formal formation do not follow this trend suggests that the latter are formed largely by a nucleophilic displacement process.

The relative rates of deuterium exchange (i.e.,carbanion formation) of the methylene halides stand in the order CH 2Br2CH2 I2 >CH2 C1Br>CH2 C1 2 .

The observed order coincides with that previously observed by Hine , and co-workers (4, 5) in their work with the haloforms and leads to the conclusion that the ability of a-halogens to stabilize the carbanion stand in the order Br--I>C1>>F. This order seems to be consistent with the 148

expected ability of the halogens to expand their octets and thus stabilize the resulting dihalomethyl anion by resonance.

Study of the Deuterium Content of the Formal

Evidence for an a-elimination mechanism in the reactions of the methylene halides with alkoxides in deuteroalcohols can be obtained by

determining the deuterium content of the methylene group of the formal produced. If the formals are formed by an a-elimination mechanism, at least one gram-atom of deuterium will be incorporated into the methylene group of each mole of the formal produced [mechanism (8)]. On the other hand, if the formals are formed by a simple nucleophilic displacement at

carbon (SN 2 reaction mechanism), then the amount of deuterium will be governed by the rate of deuterium exchange of the methylene halide, and

its subsequent rate of formal formation.

In studies of this nature, it is necessary to establish that the formals produced do not undergo deuterium exchange under the reaction conditions. To ascertain the susceptibility of the formals to deuterium exchange, a 0.5 M solution of methylal was prepared in 1 M sodium methoxide in methyl alcohol-d. A portion of the reaction mixture was sealed in an nmr tube and the ratio of 0-H to -CH 2- absorptions of methanol and methylal, respectively, determined by nmr spectroscopy. The

sample was then heated at 80° for 24 hours after which the nmr, spectrum was again recorded. No change in the ratio of 0-H to -CH 2 - absorptions was observed indicating that no deuterium exchange had occurred; therefore, examination of the deuterium content of the methylene group of the formal produced in ROD would give a valid indication of the reaction mechanism. 149

Despite the simplicity of this method for obtaining information

about the reaction mechanism, it is subject to certain experimental

limitations, namely, that the rate of deuterium exchange of the methylene

halide being studied is not markedly faster (an order of magnitude or

more) than the rate of formal formation. If this condition is not met,

the methylene halide will be significantly deuterated before a sufficient

quantity of formal can be produced for analysis. The formal isolated

under these conditions will be almost completely deuterated and no quan-

titative information regarding the reaction mechanism can be derived.

Examination of the rates of formal formation (Table 2) and the rates of

deuterium exchange (Table 6) shows that the reaction of methylene bromide with sodium methoxide in methyl alcohol-d is ideally suited for this type of study, the ratio of rate constants, k /k being about 3:1. exchange formal' Two independent studies of the methylal produced from the reaction

of methylene bromide with sodium methoxide in methyl alcohol-d were

carried out. The first reaction was terminated after nine percent

reaction and the second after 13 percent reaction (based on the amount of methoxide consumed). The methylal produced in the two reactions was

isolated and purified by preparative gas-liquid chromatography and the

ratio of methylene to methyl protons determined by nmr spectroscopy. The nmr spectrum of each sample was integrated ten or more times and the

average value of CH /CH used in the subsequent calculations. The ratio 2 3 of the methylene to methyl proton signals was also determined on a sample of pure (unexchanged) methylal by the same procedure and at the same

instrument settings employed in the analysis of the exchanged methylal.

This value was used to correct the experimentally determined values to 150

the theoretical value of 1:3.

The ratio of methylene to methyl protons was then calculated from equations (12) and (13), (Derivation 2, Appendix). These equations were derived assuming that the methylal was formed solely by a nucleophilic displacement process (SN2 reaction mechanism). In this model, the only way that deuterium can be incorporated into the formal is by prior exchange of the methylene bromide followed by its subsequent reaction with methoxide to yield methylal. Comparison of the calculated and experimentally determined values of CH 2 /CH 3 is shown in Table 7.

Table 7. Deuterium Content of the Methylal Produced in the Reaction of Methylene Bromides with Met ho xi deb in Methyl Alcohol-d

CH /CH CH /CH Reaction Time Percent 2 3 2 3 Seconds Reaction calc. c exp.

44400 9.0 0.326 0.301 ±0.016 71040 13.1 0.322 0.299 ±0.008 a h Initial concentration, 1.305 M. Initial concentration, 1.052 M. c -6 -6 k = 6.36 x 10 and k = 2.35 x 10 . I 5

As the results show, the experimentally determined ratios of methylene to methyl protons are somewhat lower than the calculated values indicating that a small fraction of the methylal is formed by an a-elimination mechanism. Unfortunately, the method chosen for determining

11 I 151

the deuterium content of the formal did not prove as sensitive as one would have liked because the reaction was largely S N 2 in character. For example, in the run carried to 71,040 seconds, it was necessary to determine the difference between a theoretical CH 2 /CH 3 ratio of 0.333 and a calculated value of 0.322. By using a standard method for determining the CH /CH ratio, a consistent set of values could be obtained; however, 2 3 by choosing another standard method for making the measurements, a variation of as much as ten percent could be obtained. The method, therefore, was not especially sensitive to small differences. As a result, it is difficult to state unambiguously that there is clearly

some participation from an a - elimination mechanism. It is clear, however, that the reaction is largely S N 2 in character. If the reaction was proceeding either largely or entirely by an a-elimination mechanism, the values of CH /CH would have been 0.166 or less which is well beyond 2 3 the experimental uncertainty of the nmr method.

The Mass-Law Effect

As previously discussed, the mass-law effect constitutes some of the most convincing evidence for the presence of a dihalomethylene intermediate in the basic hydrolysis of the haloforms (7). In an effort to obtain similar evidence in the methylene halide series, the reaction of methylene iodide with sodium methoxide in methyl alcohol-d was studied in the presence of added sodium iodide.

Reaction of Methylene Iodide with Sodium Methoxide in Methyl

Alcohol-d in the Presence of Sodium Iodide and Sodium Perchlorate. The reaction of methylene iodide with methoxide in methyl alcohol-d was studied in the presence of 0.50 M sodium iodide. For comparative 152

purposes, the reaction was also studied in the presence of 0.50 M sodium perchlorate. The latter was necessary to cancel the ionic strength effect on the rate constants thereby establishing a valid case for comparison. The results of this study (Tables 56 through 59, Appendix) are shown in Table 8.

Table 8. Reaction of Methylene Iodide with Methoxide in the Presence of Iodide and Perchlorate

Concentration k x 10 6 Electrolyte R. mole -1 sec-1 none - 2.40 ±0.07 sodium iodide 0.50 2.37 ±0.10 sodium iodide 0.50 2.36 ±0.22 sodium perchlorate 0.50 2.49 ±0.25 sodium perchlorate 0.50 2.21 ±0.10

As the results show, the average value for the rate constants for the reaction of methylene iodide with sodium methoxide in methyl alcohol-d in the presence of sodium perchlorate is 2.35 ± 0.17; for the same reaction in the presence of sodium iodide it is 2.36 ± 0.16. This excellent agreement seems to rule out the presence of a mass-law effect.

Unfortunately, direct comparison of the average values of the rate constants from the two studies is not entirely valid because the rate constants have a tendency to decline as the reaction progresses. This trend, however, was not entirely unexpected. Hine, Thomas, and Ehrenson (7) observed declining rate constants in the reaction of methylene iodide 153

with methoxide (in methyl alcohol) and attributed this trend to photolysis and/or oxidation of the methylene halide during the extended period required to obtain the data. In agreement with the earlier work, our studies in methyl alcohol-d in the absence of added salts (Tables 13 through 17, Appendix) also showed similar trends. In contrast, however, the rate constants obtained for the reactions of methylene iodide with potassium isopropoxide and tert-:butoxide in the corresponding protio- and deuteroalcohols do not exhibit this trend. These observations suggest that the reason for the declining rate constants may be different than previously proposed.

Hine and co-workers determined the rate constant for the reaction of methylene iodide with methoxide by extrapolating the rate constants to zero time. Application of the same technique to our data gives widely divergent values for a given added salt, i.e., sodium iodide or sodium perchlorate. Furthermore, a plot of the rate constants versus percent reaction does not suggest any apparent correlation during the latter stages of the reaction. Because of the large average deviations in the rate constants coupled with the observed trends toward falling rate constants, it is difficult to state unambiguously whether there is an observable mass-law effect; however, the majority of the evidence seems to be against it.

As previously mentioned, the mass-law effect is indeed strong evidence for an intermediate dihalomethylene in the basic hydrolysis of the haloforms; however, the absence of such an effect is not equally satisfactory proof that divalent•carbon intermediates do not exist. In order for mass-law effects to be observable, the intermediate must be 154

sufficiently stable to be selective in its subsequent reactions with the various nucleophilic reagents present; otherwise, it will combine with one of the first molecules it meets, very probably a solvent molecule.

In regard to the differences in stability between mono- and dihalomethylenes, Closs and Schwartz (17) have shown that the intermediate generated by the reaction of n-butyllithium with methylene chloride was only twelve times as reactive toward 2,3-dimethylbutene-2 as toward pentene-1. On the other hafid, the intermediate generated by the reaction of the same organometallic compound with chloroform was 400 times more reactive toward 2,3-dimethylbutene•2 was toward pentene-1. If indeed divalent-carbon species are intermediates in these reactions, it seems evident that the dihalomethylenes are markedly more stable than the mono- halomethylenes. It is possible, therefore, that the reaction of methylene iodide with sodium methoxide could be proceeding through a divalent- carbon intermediate but that due to the instability of the methylene intermediate, the mass-law effect is not observed.

At this point, the logical extension of the mass-law studies to the reactions of potassium isopropoxide and tert-butoxide (in the corresponding alcohols)with methylene iodide or other methylene halides was considered. The rates of deuterium exchange of the methylene halides in 3 4 these systems are 10 to 10 times as fast as the corresponding rates of formal formation. In the reaction of methylene iodide with methoxide in methanol, the rate of deuterium exchange was only about 15 times as fast as the rate of formal formation. As a result, the chances of a dihalomethyl anion losing halogen to give a monohalomethylene are greatly 155

extended in isopropyl and tert-butyl alcohols. Unfortunately, potassium iodide, bromide, and chloride are virtually insoluble in these solvents and the mass-law study was necessarily terminated with the results obtained in methyl alcohol-d.

Product and Stoichiometry Studies

The preparation of diisopropyl formal has been reported by

Knoefel (39)1 by Jansson (40), and by Normant and Crisan (41). The preparation of di(tert-butyl) fdomal has been reported by Jansson (40) and by Leimu (42). A variety of methods have been employed including

1) direct acid-catalyzed condensation of formaldehyde with the appropriate alcohol, 2) acid-catalyzed alcoholysis of methylal or other simple formals,

3) the pyridine-catalyzed reaction of alkyl halomethyl ethers with alcohols, and 4) the reaction of alkyl halomethyl ethers with magnesium and iodine in tetrahydrofuran. Unfortunately, the preparation of these formals by the reaction of potassium isopropoxide and tert-butoxide with the methylene halides has not been reported. In order to demonstrate that the expected formals were the principal produc(ts from the reaction of isopropoxide and tert-butoxide with the methylene halides and also to demonstrate that the stoichiometry of these reactions was indeed that employed in deriving the kinetic equations, a study of the stoichiometry and the reaction products from these reactions was undertaken.

Preparation of Di(tert-Butyl) Formal. Di(tert-butyl) formal was prepared by the reaction of potassium tert-butoxide with methylene bromide in tert-butyl alcohol. The boiling point, 145° (uncorrected) is in good agreement with that reported by Leimu (42), i.e., 50P to 52°

(18 mm.), if the atmospheric boiling point is interpolated from the vapor 156

pressure-temperature nomograph developed by Lippincott and Lyman (43).

The infrared and nmr spectra of di(tert-butyl) formal are shown in

Figures 3 and 8, Appendix. The nmr spectrum in particular leaves no

doubt as to the structure of the product. The singlet at 5.20 T is due to the hydrogens of the methylene group of the formal and the singlet at

9.0 T is due to the hydrogens of the methyl groups. The integrated areas of the absorptions are in the expected 1:9 ratio.

Preparation of Diisopropyl Formal. Diisopropyl formal was prepared by the reaction of potassium isopropoxide with methylene chlorobromide in isopropyl alcohol. The boiling point, 117° (uncorrected), is in good agreement with that reported by Knoefel (39), i.e., 118°, and is identical to that reported by Normant and Crisan (41). The infrared and nmr spectra of diisopropyl formal are shown in Figures 2 and 9, Appendix.

The nmr spectrum in particular leaves no doubt as to the structure of the product. The singlet at 5.62 T is due to the hydrogens of the methylene group. The, sextet at 6.45 T is due to the methine hydrogens. The observed multiplicity is apparently due to the non-equivalence of the methyl-protons (44). The doublet at 9.18 T is due to the methyl hydrogens.

The integrated areas of the absorptions are in the expected 1:1:3 ratio.

Stoichiometry of the Reactions. The stoichiometry of the reactions of potassium tert-butoxide and isopropoxide with the methylene halides was determined, following the changes in the concentration of strong base and, halide. The results of this study are shown in Table 9 (Tables 60 and 61, Appendix). As the results show, the stoichiometry for the reaction of methylene chlorobromide with potassium isopropoxide is exactly that demanded for formal formation. The reaction of potassium 157

tert-butoxide with methylene bromide, however, yields values of

A[RO- ]/A[X- ] ranging from 0.95 to 1.04. Although the average of these values is close to unity, the variation leads to some question as to the stoichiometry of the reaction; however, it seems evident that the stoichiometry is at least very close to that employed in the derivation of the rate equations.

Table 9. Stoichiometry of the Reactions of Isopropoxide and tert-Butoxide with the Methylene Halides

Methylene Percent Alkoxjde Halide Reaction A[RO - ]/.A[X-] tent-BuOK CH Br 35 0.944 2 2 tert-BuOK CHBr 35 0.950 22 tert-BuOK CH Br 85 1.04 2 2 tert-BuOK CH Br 85 2 2 1.03 iso-PrOK CH 58 1.00 2 C1Br 158

CHAPTER IV

CONCLUSIONS

A study of the reactions of methylene iodide, bromide, chloride, and chlorobromide with sodium methoxide and potassium isopropoxide and

tert-butoxides in the corresponding protio- and deuteroalcohols has shown

that the transformations (which give the corresponding formals) proceed

largely by a nucleophilic displacement process. No conclusive evidence

that the reactions proceed either largely or in part by an a - elimination mechanism could be obtained. The following observations support the S N2 reaction mechanism: 1) The relative reactivities of the alkoxides toward a given methylene halide stand in the order isopropoxide > tert- butoxide > methoxide. The relative reactivities of the alkoxides toward chloroform (a substance known to react by an a-elimination mechanism) stand in the order tert-butoxide > isopropoxide > methoxide. The con- flicting orders suggest that the methylene halides react largely by a nucleophilic displacement process. 2) The rate constants for deuterium exchange (carbanion formation) of the methylene halides have been shown to increase with increasing base strength, i.e., tert-butoxide > isopropoxide > methoxide. This is the same order observed for the reactivities of the alkoxides toward chloroform. Since both reactions proceed through carbanion formation, the reactivities would be expected to parallel each other. 3) The deuterium content of the methylal produced by the reaction of sodium methoxide with methylene bromide in methyl 159

alcohol-d is very close to that predicted by assuming that the methylene halide first undergoes deuterium exchange and then reacts with methoxide

(by a nucleophilic displacement process) to give the formal. The deuterium content is nowhere near that predicted for an a-elimination mechanism. 4) There is no observable decrease in the rate constants for reaction of methylene iodide with sodium methoxide in the presence of added sodium iodide,i.e., no observable mass-law effect. The absence of a mass-law effect suggests that monoiodomethylene is not an intermediate in the reaction and serves as additional evidence against the a-elimination mechanism. 160

APPENDI X 161

Derivation 1. Rate Expression for the Deuterium Exchange of the Methylene Halides

Case 1: Rate of Deuterium Exchange Comparable to Rate of Formal Formation

Chemical Equations

k l CH2 X2 + RO- kCHX. + ROH (1)

k 2 CHX2 + ROD 4--- CHDX + RO - (2) k 2 -2

k 3 CHDX + RO- CDX2 + ROH 2 k _3 2 (3)

k 4

CDX + ROD "" CD2 + RO- 2 4-k X2 -4

k CH X (CHDX or CD X ) + RO 5 2 2 2 2 - ROCH 2 X + X-

k ROCH X + RO- (RO)., CH + X- 2 2

Assumptions,

1. k [ROD] » k [ROH]. k [ROD] » k [ROH] 2 -1 ' 4 -3

2. k = 2 k 1 3

3. k [ROCH 6 2 X] » k 5 [CH 2 X2 ]

4. Neglect k_ 2 and k_ 4 5. k k 1 5 162

Symbols and Definitions

p = A[ROH] = [CHDX ] + [CHD(OR) ] + 2[CD (OR) 2] + 2[CD 2 X2 2 2 2 ]

2 (OR) 2 = [CH ] + [CHD(0R) 2 ] + [CD2 (OR) 2 ]

a = [CH 2X 2], °

] + [CHDX a-x= [CH 2 X2 2 ] + fq X

b = [R0 - ] 0

b-2x = [R0 - ]

d = [CHD(OR) ] + 2[CD2 (OR) 2 ] 2

M = [CHDX2 ] + 2[CD2 X2 ]

Derivation

dx = k ([CH 2X2 ] + [CHDX ] dt 5 [CD2 X2 D[R0 - ] = k 5 (a-x)(b-2x) (7)

= k ([CH X ] + [CHDX ]/2)[R0 - ] d 1 2 2 2

= k (a-x-p/2+d/2)(b-2x) d 1 dd = k 5 [R0- MCHDX D2X2])2[C dt 2 ] + 2[CD2 X2 ])

dd = k dt 5 (b-2x)(p-d)

Solution of Equation (11)

1. Plot p versus t (Figure 10, Appendix).

2. Plot [R0- ] versus t (Figure 11, Appendix).

3. Divide the time in Figure 10 into a number of small intervals. Determine the average value of [R0 - ] during each time interval from Figure 11. Determine the average value of p from the first time interval (Figure 10), then calculate Ad'.

163

Ad' = k [R0-] ang At 1 5 avg p

Then determine Ad l .

Ad = k [R0- ] (p - -Ad ) At 1 5 avg avg -A di)

Calculate Ad for the second time interval. 2 1 Ad - [R0-] d - -A d ) At 2 = k5 avg avg 1 2 1

Continue the calculations for Ad n n-1 1 Ad = k [R0-] (p - E Adf - At n 5 avg avg m-1 m 2-Adn-1)

4. Plot AA versus t (Figure 12), .n 5. Using Figure 12, obtain the values of d at the time (t) at which p was determined.

6. Plot (a-x-p/2+d/2) versus p (Figure 13).

7. Express (a-x-p/2+d/2) as a-mp, where m is the slope obtained from the plot of (a-x-p/2+d/2) versus p as shown in Figure 13.

8. Plot [R0 - ] versus p (Figure 14).

9. Express [RO - ] as b-cp where c is the slope obtained from the plot of [RO - ] versus p and b is the intercept, [R0 - ] 0 , as shown in Figure 14.

Solution to Equation (9)

1. Substitute the values obtained from the solution of equation (11), i.e., (a-mp) and (b-cp), into equation (9).

dP = k 1 (b-cp)(a-mp) (12) dt

2. Integrating equation (12) between limits gives equation (13).

2.303 b(a-mp) (13) k l = a(b-cp) 164

Case 2: The Rate of Deuterium Exchange is Much Faster than the Rate of Formal Formation

Chemical Equations

Same as Case 1.

Assumptions

Same as Case 1 except that kJ. » k 5 .

Symbols and Definitions

Same as Case 1.

Derivation = k ([CH ] + [CHDX ]/2)[R0 - ] (17) d 1 2 X2 2 k (a-x-p/2+d/2)(b-2x) dt 1 (18)

Since k >> k p/2 » x + d/2 1 5'

d= k (19) 1 (b-2x)(a,7-P12).

Solution to Equation (19)

1. Plot [RO- ] versus p (Figure 18).

2. Express [RO- ] as b-cp where c is the slope obtained from the plot of [R0 - ] versus p and b is the intercept, [RO - ] o , as shown in : Figure 15.

k (b-cp)(a-p/2) dt 1 (20)

3. Integrating equation (20) between limits gives equation (21).

2.303 b(a-p/2) k - log 1 (ac-b12)t ' a(b-cp) (2 1) 165

Derivation 2. Expression for Calculating the Deuterium Content of Formals Formed in Deuteroalcohols

Chemical Equations

Same as Derivation 1, Case 1,

Assumptions

Same as Derivation 1, Case 1.

Symbols,

Same as Derivation 1, Case 1.

Derivation

.42i= k (a -x)(b -2x) dt 5

dd = k (ECHDX 2 [CD X ])[R0] dt 5 21 2 2

dd = k M (b-2x) dt 5

dM = (k [CH2 X2 ] + k [CHDX ] - k M)(b-2x) dt 1 3 2 5

dM = (k1 [CH 2 X2 ] + k [CHDX ]/2 - k dt 1 2 5 M)(b-2x)

Since: a-x = [CH X ] + [CHDX 2 2 2 ]/2 + M/2

dM _ k a-x-M/2)(b-2x) - k 5 M(b-2x) dt

Dividing equation (7) by (1) gives equation (8)

a-x (8) 166

Dividing equation (3) by (1) gives equation (9). dd = M/(a-x) dx ( 9 )

Integrating equation (8) gives equation (10): M = 2(a-x)[1 - (1 - x/a) k" ] (10) where k' is k /2k . 1 5

Substituting equation (10) in (9) gives equation (11). dd dx = 2[1 - (1 - x/a) k ]

Integrating equation (11) gives equation (12).

k'+1 . d x + a[(1 x/a) - 1.] (12) 2 - k' + 1

The ratio of methylene to methyl proton signals for partially- exchanged methylal may be calculated from equation (13).

CH /CH = (2x-d)/6x 2 3 (13 ) 167

Table 10

Extinction Coefficients for Methyl Alcohol in Methyl . Alcohol-d

Absorbance , Absorbance Added Me0H Cell-Path Length Me0Na c

MeQD McOD .. t Me0H M cm.x 10 3

0.171 0.724 0.737 5.60 0.0 134 0.171 0.469 0.380 5.60 0.0 135 0.192 0.455 0.340 5.78 0.269 134 0.192 0.718 0.673 5.78 0.269 135 0.182 0.668 0.655, 5.70 0.863 130 0.182 0.419 0.307 5.70 0.863 135 0.190 0.662 0.725 5.84 1.143 112 0.190 0,429 0.363 5.84 1.143 113 0.174 0.579. 0.682 5.82 1.434 102 0.174 0.388 0.356 5.82 1.434 103 168

Table 11

Extinction Coefficients for Isopropyl Alcohol in Isopropyl Alcohol-d

Added Absorbance Absorbance iso-PrOH Cell-Path L9gth iso-PrOK isp-PrOD iso-PrOD + iso-PrOH M cm. x 10

0.251 0.660 0.455 5.34 0.0 170 0.251 0.454 0.227 5.34 0.0 168 0.251 0.359 0.114 5.34 0.0 171 0.262 0.730 0.518 5.32 0.0 170 0.262 0.494 0.259 5.32 0.0 168 0.341 0.703 0.428 5.37 0.095 158 0.416 0.761 0.425 5.38 0.189 151 169

Table 12

Extinction Coefficients for tert-Butyl Alcohol in tert-Butyl Alcohol-d

Added Cell-Path Absorbance Absorbance tert-BuOH Length tert-BuOK tert-BuOD tert-BuOD + tert-BuOH M cm. x 10 3

0.091 0.256 0.403 2.83 0.0 145 0.115 0.344 0.450 3.38 0.0 150 0.091 0.466 0.880 2.83 0.0 150 0.112 0.308 0.431 3.37 0.284 132 0.112 0.458 0.798 3.37 0.284 129 0.102 0.248 0.413 3.38 0.565 104 0.102 0.395 0.816 3.38 0.565 106 170

Table 13

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6d HC10 CH I 10 Time 4 Me0Na 2 2 Percent Seconds ml. b Reaction g. mole sec. -1 c 0 0.4280 0.4397 9840 15.42 0.4391 46680 13.60 0.3873 0.4193 9.5 2.49 87540 12.43 0.3539 0.4026 17.3 2.88 125800 11.60 0.3303 0.3908 22.8 2.49 260640 9.20 0.2620 0.3567 38.8 2.40 371700 7.80 0.2221 0.3367 48.1 2.32 445860 7.01 0.1996 0.3255 53.4 2.30

Avg. 2.43 ±0.09 a b c 0.1139 M. Per 4.0 ml. sample. Calculated from dilution factors. d Calculated from equation (5). 171

Table 14

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6 d HC10 CH I 10 k Time 4 Me0Na 2 2 Percent -1 Seconds ml. b Reaction t. mole sec. c 0 0.4478 0.4446 2880 15.57 0.4433 0.4424 1.0 26700 14.20 0.4328 0.4371 3.3 161700 10.87 0.3095 0.3755 30.9 2.81 265560 9.16 0.2608 0.3511 41.8 2.60 440780 7.04 0.2005 0.3210 55.2 2.45 538460 5.08 0.1446 0.2930 67.7 3.00

Avg. 2.71 1 0.19

c d a0.1139 M. bPer 4.0 ml. sample. Extrapolated. Calculated from equation (5). 172

Table J5

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6 d HC10 CH I 10 k Time 4 Me0Na 2 2 Percent -1 -1 Seconds ml. b M M Reaction Z. tole sec. c 0 44.11 0.5596 0.4672 c 88380 36.20 0.4593 0.4171 17.9 2.53 c 150840 35.81 0.4073 0.3911 27.2 2.47 181140 33.72 0.3836 0.3792 31.5 2.49 255300 29.80 0.3390 0.3569 39.4 2.42 268020 29.24 0.3326 0.3537 40.6 2.41 317820 27.04 0.3076 0.3412 45.0 2.39 424080 23.11 0.2629 0.3189 53.0 2.35 512700 20.50 0.2332 0.3040 58.3 2.32 596040 18.40 0.2093 0.2921 62.6 2.30 663060 16.82 0.1913 0.2831 65.8 2.30

Avg. 2.40 ±0.07 a 0•04550 M. bPer 4.0 ml. sample. cM1. of 0.05075 M HC10 4 per 4.0 ml. d sample. Calculated from equation (5). 173

Table 16

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36°

6 d HC10 a CH I 10 k Time 4 Me0Na 2 2 Percent -1 -1 Seconds ml. b M M Reaction L. mole sec.

0 0.4489 c 0.4441 360 15.68 0.4465 0.4429 0.5 16.8 e e 3600 15.63 0.4450 0.4422 0.9 2.73 10800 15.37 0.4370 0.4382 2.7 2.82 e 24780 14.85 0.4228 0.4311 5.8 2.76 40260 14.39 0.4097 0.4245 8.7 2.61 87480 13.11 0.3733 0.4063 16.8 2.48 112320 12.43 0.3539 0.3966 21.2 2.52 124800 12.14 0.3457 0.3925 23.0 2.51 183180 10.91 0.3106 0.3750 30.8 2.47 216600 10.38 0.2955 0.3674 34.2 2.40

Avg. 2.53 ±0.08

a b c d 0.1139 M. Per 4.0 ml. sample. Extrapolated. Calculated from equation (5). e0mitted from average, < five percent reaction. 174

Table 17

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d at 36°

6 c HC10 CH I 10 k Time Me0Na 2 2 Percent -1 Seconds ml. b M M Reaction R. mole sec.

0 54.45 0.6677 0.5020 0 54.16 0.6641 0.5020 66360 45.80 0.5616 0.4498 15.7 2.70 93600 42.55 0.5218 0.4299 21.6 2.81 189240 34.90 0.4280 0.3830 35.7 2.68 267060 30.83 0.3781 0.3580 43.2 2.53 339660 27.84 0.3414 0.3397 48.7 2.42

431 .520 23.24 0.2850 0.3115 57.2 2.55 590580 18.12 0.2222 0.2801 66.6 2.58 787380 15.01 0.1841 0.2610 72.4 2.37 935220 13.50 0.1655 0.2518 75.2 2.22

Avg. 2.54 1 0.14 a b 0.04905 M. Per 4.0 ml. sample. cCalculated from equation (5). 175

Table 18

Reaction of Methylene Iodide with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 c HC10 CH I 10 k Time 4 iso-PrOK 2 2 Percent -1 -1 Seconds, ml. b M M Reaction Z. mole sec.

0 38.84 0.4928 0.4175 47040 30.54 0.3875 0.3649 21.4 6.56 78840 25.98 0.3296 0.3359 33.1 6.85 134700 20.48 0.2598 0.3010 47.3 6.79 152100 18.95 0.2404 0.2913 51.2 6.88 164340 17.90 0.2271 0.2847 53.9 6.96 189600 16.11 0.2044 0.2733 58.5 7.03 220860 14.33 0.1818 0.2620 63.1 7.03 232260 13.73 0.1742 0.2582 64.6 7.04 239940 13.34 0.1692 0.2557 65.7 7.05 308280 10.90 0.1383 0.2403 71.9 6.81

Avg. 6.90 '1 0.12

a b 0.05075 M. Per 4.0 ml. sample:, from equation (5). 176

Table 19

Reaction of Methylene Iodide with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 d HC10 CH I 10k Time 4 iso-PrOK 2 2 Percent -1 -1 Seconds ml. b M M Reaction Z. mole sec. c 0 0.4885 0.5272 e 1440 36.40 0.4841 0.5250 0.9 5.97 e 3540 36.00 0.4788 0.5224 2.0 5.40 e 3540 36.00 0.4788 0.5224 2.0 5.40 e 6300 35.35 0.4701 0.5180 3.8 5.83 20940 31.80 0.4229 0.4944 13.4 6.75 54780 25.28 0.3362 0.4511 31.2 7.02 84840 21.31 0.2834 0.4247 42.0 6.84 101820 19.31 0.2568 0.4114 47.4 6.85 106680 18.80 0.2500 0.4080 48.8 6.85 116340 17.96 0.2389 0.4024 51.1 6.76 173340 13.72 0.1825 0.3742 62.6 6.54

186240 . 12.85 0.1709 0.3684 65.0 6.56

Avg. 6.77 ±0.12 a b c d 0.05319 M. Per 4.0 ml. sample. Extrapolated. Calculated from e equation (5). Omitted from average, < five percent reaction. 177

Table 20

Reaction of Methylene Iodide with Potassium ,Isopropoxide in Isopropyl Alcohol-d at 36°

a 6 c HC10 CHII 10 k Time 4 iso-PrOK 2 Percent -1 Seconds. ml.b Reaction Z. mole sec. -1

0 39.08 0.4958 0.4232 39720 28.77 0.3650 0.3578 26.4 91o 9 56820 25.41 0.3224 0.3365 35.0 10.1 63420 24.33 0.3087 0.3297 37.7 10.1 69720 23.32 0.2959 0.3233 40.3 10.1 124140 16.66 0.2114 0.2810 57.4 10.2 133840 15.50 0.1966 0.2736 60.3 10.4 143820 14.70 0.1865 0.2686 62.4 10.4 154500 13.97 0.1772 0.2639 64.3 10.3 179580 12.28 0.1558 0.2532 68.6 10.2 211620 10.70 0.1357 0.2432 72.6 10.0

Avg. 10.2 1 0.1

a b c 0.05075 M. Per 4.0 ml. sample. Calculated from equation (5). 178

Table 21

Reaction of-Methylene Iodide with Potassium tert-Butoxide in tert-Butyl Alcohol at 36°

a 6 d HC10 CH I 10 k Time 4 tert-BuOK 2 2 Percent -1 -1 Seconds ml. M M Reaction L. mole sec.

0 0.4892 c 0.4985 e 1320 35.10 0.4879 0.4979 0.3 2.02 e 7080 34.67 0.4819 0.4949 1.5 2.14 e 11100 34.35 0.4775 0.4927 2.4 2.20 29340 32.86 0.4567 0.4823 6.6 2.39 111300 27.65 0.3843 0.4461 21.4 2.30 119760 27.12 0.3770 0.4424 22.9 2.32 168300 24.75 0.3440 0.4259 29.7 2.28 185760 23.90 0.3322 0.4200 32.1 2.29 197520 23.40 0.3252 0.4165 33.5 2.28 254520 21.09 0.2931 0.4005 40.0 2.27 273780 20.30 0.2822 0.3950 42.3 2.28

Avg. 2.30 ±0.03 a b c d 0.05560 M. Per 4.0 ml. sample. Extrapolated. Calculated from equation (5) 179

Table 22

Reaction of Methylene Iodide with Potassium tert-Butoxide in tert-Butyl Alcohol--d at 36°

a 6 d HC10 10 k Time 4 tert-BuOK CHI Percent2 -1 Seconds ml. b Reaction k . mole sec. c 0 0.5397 0.4979 e 4020 35.21 0.5317 0.4939 1.5 3.75 e 8460 34.61 0.5226 0.4894 3.2 3.86 13140 33.97 0.5129 0.4845 5.0 3.95 21300 32.91 0.4969 0.4765 7.9 3.98 92340 25.59 0.3864 0.4213 28.4 3.96 104400 24.50 0.3699 0.4130 31.5 4.01 155880 20.90 0.3156 0.3859 41.5 3.96 170940 20.00 0.3020 0.3791 44.0 3.95 187800 18.95 0.2861 0.3711 47.0 3.98

Avg. 3.97 1 0.02 a d 0.06040 M. bPer 4.0 ml. sample. -cExtrapolated. Calculated from e equation (5). Omitted from average < five percent reaction. 180

Table 23

Reaction of Methylene Bromide with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6 d HC10 CH Br 10 k Time 4 Me0Na 2 2 Percent ml -1 -1 Seconds b M M Reaction k. mole sec.

0 44.20 0.5608 0.5358 61980 38.10 0.4834 0.4971 13.8 2.32 91260 35.64 0.4522 0.4815 19.4 2.32 152220 34.96 c 0.3977 0.4543 29.1 2.30 168660 33.75 c 0.3839 0.4474 31.5 2.30 256620 28.40 c 0.3230 0.4169 42.4 2.29 269082 27.90 c 0.3174 0.4141 43.4 2.27 c 318660 25.53 0.2904 0.4006 48.2 2.26 c 341340 24.50 0.2787 0.3948 50.3 2.26 c 425640 21.32 0.2425 0.3767 56.8 2.23

513. 880 18.41 0 0.2094 0.3601 62.7 2.24

Avg. 2.28 •±0.03

a c 0.05075 M. bPer 4,0 ml. sample. Ml. of 0.0485q M HC104 per 4.0 ml. sample. ; dCalculated from equation (5). 181

Table 24

Reaction of Methylene Bromide with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6 d HC10 CH Br 10 k Time 4 Me0Na 2 2 Percent -1 -1 Seconds ml. M M Reaction L. mole sec.

0 0.5608 c 0.7598 e 3600 19 ,.45 0.5538 0.7563 1.2 2.30 18600 18.40 0.5239 0.7414 6.6 2.44 65760 15.52 0.4419 0.7004 21.2 2.49 100680 13.83 0.3938 0.6763 29.8 2.46 150720 11.76 0.3349 0.6469 40.3 2.45 184860 10.56 0.3007 0.6298 46.4 2.46 240720 1.)08 0.2585 0.6087 53.9 2.39 274860 8.32 0.2369 0.5979 57.8 2.36

Avg. 2.43 ±0.04 a b 0.1139 M. Per 4.0 ml. sample. 'Extrapolated., calculated from e equation (5). Omitted from average, < five percent reaction. 182

Table 25

Reaction of Methylene Bromide with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 f HC10 CH Br 10 k Time 4 iso-PrOK 2 2 Percent -1 -1 Seconds ml. M M Reaction t. mole sec. e 0 0.4972 0.4980 1680 35.40 0.4921 0.4965 1.0 6.18 g 5460 0.4813 6.02 g 34.63 0.4901 3.2 h 9120 33.81 0.4699 0.4844 5.5 6.30

18480 33.93 0.4438 0.4713 10.7 6.35 86280 23.25 0.3232 0.4110 35.0 5.55 108100 21.11 0.2934 0.3961 41.0 5.54 117960 20.30 0.2822 0.3905 43.2 5.49 127500 19.51 0.2712 0.3850 45.5 5.48 199500 15.35 c 0.2041 0.3515 59.0 5.45 208260 14.82 0.1971 C).3480 60.4, 5.45 d 269040 5.90 0.1569 0.3279 68.4 5.45

Avg. 5.49 10.03 a b c 0.05560 M. Per 4.0 ml. sample. Ml. of 0.05319 M HC10 per 4.0 ml. 4 d sample. Ml. of 0.05319 M HC104 per 2.0 ml. sample. eExtrapolated. f Calculated from equation (5).gOmltted • from average, < five percent h reaction. Omitted from average, deviation from the average > four times the average deviation.

183

Table 26

Reaction of Methylene Bromide with Potassium Isopropoxide in Isopropyl Alcohol-d at 36°

a HC10 CH Br 1060 Time 4 iso-PrOK 2 2 Percent -1 Seconds ml.b M M Reaction Z. mole sec. -1

0 39.30 0.4986 0.4493 d 2700 38.40 0.4872 0.4439 2.3 9.59 d 5160 37.60 0.4770 0.4388 4.3 . 9.66 53760 27.27 0.3460 0.3733 30.6 8.33 68580 24.98 0.3169 0.3588 36.4 8.28 85680 22.68 0.2877 0.3442 42.3 8.23 141300 17.09 0.2168 0.308'7 56.5 8.07 152760 16.15 0.2040 0.3028 58.9 8.07 171180 14.80 0.1878 0.2942 62.3 8.05 196000 13.11 0.1663 0.2835 66.6 8.11 226560 5.62 0.1426 0.2716 71.4 8.24

Avg. 8.17 ±0.10

a b c _ 0.05075 M. Per 4.0 ml. sample. Calculated from equation (5). d Omitted from average, < five percent reaction. 184

Table 27

Reaction of Methylene Bromide with Potassium tert-Butoxide in tent-Butyl Alcohol at 36°

a 6 d HC10 CH Br 10 k Time 4 tert-BuOK 2 2 Percent -1 -1 Seconds ml.b M M Reaction Z. mole sec. c 0 0.4392 0.4652 240 15.43 0.4394 e 13020 15.20 0.4328 0.4620 1.4 1.22 153960 10.57 0.3010 0.3961 31.5 2.87 260280 8.43 0.2400 0.3656 45.4 2.84 435040 6.01 0.1711 0.3311 61.0 2.82 521940 5.20 0.1481 0.3196 66.3 2.78

Avg. 2.83 ±0.03 a b d 0.1139 M. Per 4.0 ml. sample. Extrapolated. Calculated from e equation (5). Omitted from average, < five percent reaction. 185

Table 28

Reaction of Methylene Bromide with Potassium tert-Butoxide in tert-Butyl Alcohol at 36°

a HC10 CH Br Time 4 tert-BuOK 2 2 Percent ml. Seconds b M M Reaction Z. mole sec. -1 c 0 0.4692 0.4917 e 1938 35.10 0.4668 0.4905 0.5 2.69 e 4320 34.84 0.4634 0.4888 1.2 2.94 e 6900 34.72 0.4618 0.4880 1.6 2.35 e 10740 34.30 0.4562 0.4852 2.8 2.68 48180 31.05 0.4130 0.4636 12.0 2.77 86940 28.30 0.3764 0.4453 19.8 2.71 108840 26.95 0.3584 0.4363 23.6 2.68 176940 23.30 0.3099 0.4121 33.9 2.62 205620 21.90 0.2913 0.4028 37.9 2.62 28680 18.46 0.2455 0.3799 47.7 2.64 344700 16.50 0.2194 0.3668 53.2 2.64

Avg. 2.67 -1 0.04

b c a0A5319 M. Per 4.0 ml. sample. Extrapolated. dCalculated from e equation (5). Omitted from average, < five percent reaction. 186

Table 29

Reaction of Methylene Bromide with Potassium tert-Butoxide in tert-Butyl Alcohol-d at 36°

a 6 d HC10 CH Br 10 k Time 4 tert-BuOK 2 2 Percent -1 -1 Seconds ml. b Reaction . mole sec. c 0 0.5454 0.4982 3180 35.60 0.5376 0.4943 1.4 4.56 e e 7860 34.75 0.5247 0.4879 3.8 4.99 12600 33.98 0.5131 0.4821 5.9 4.94 23580 32.31 0.4879 0.4695 10.5 4.89 83640 24.97 0.3770 0.4140 30.9 4..88 101820 23.34 0.3524 0.4017 35.4 4.82 120180 22.02 0.3325 0.3918 39.0 4.70 191280 17.60 0.2658 0.3584 51.3 4.51

Avg. 4.79 1 0.12 a b c d 0.06040 M. Per 4.0 ml. sample. Extrapolated. Calculated from equation (5). Omitted from average, < five percent reaction. 187

Table 30

Reaction of Methylene Chlorobromide with Sodium Methoxide in Methyl Alcohol-d. at 36°

a b c HC10 CI ClBr 10 k Time 4 Me0Na 1 2' Percent -1 -1 Seconds ml.b M M Reaction L. mole sec.

0 45.71 0.5199 0.6019 47580 31.41 0.3573 0.4793 31.3 7.07 79560 25.33 0.2881 0.4620 44.6 6.92 88920 23.83 0.2711 0.4577 47.9 6.90 94500 22.96 0.2612 0.45$3 49.8 6.91 101340 22.00 0.2502 0.4525 51.9 6.89 108000 21.10 0.2400 0.4500 53.8 6.88 181320 13.68 0.1556 0.4289 70.1 6.82 190800 12.92 0.1470 0.4267 71.7 6.84 202320 12.17 0.1384 0.4246 73.4 6.81

Avg. 6.89 ±0.05 a b c 0.04450 M. Per 4.0 ml. sample. Calculated from equation (5). 188

Table 31

Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 d HC10 CH ClBr 10 k Time 4 iso-PrOK 2 Percent -1 Seconds Reaction Z. mole-1 sec.

0 0.4955 c 0.4749 e 1560 34.71 0.41325 0.4684 2.6 18.1 3480 33.53 0.4661 0.4602 5.9 18.8 5400 32.40 0.4504 0.4524 9.1 19.1 9540 30.10 0.4184 0.4364 15.6 19.5 12180 28.85 0.4010 0.4277 19.1 19.3 15120 27.60 0.3836 0.4190 22.6 19.0 20160 25.32 0.3519 0.4031 29.0 19.5 25380 23.41 0.3254 0.3899 34.3 19.4 30420 21.90 0.3044 0.3794 38.6 19.0 8808b 10.52 0.1462 0.3003 70.5 19.0 113160 7.92 0.1101 0.2822 77.8 19.1

Avg. 19.2 ±0.2 a b c 0.05560 M. Per 4.0 ml. sample. Extrapolated. dCalculated from equation (5). eomitted from average, < five percent reaction. 189

Table 32

Reaction of Methylene Chlorobromide with Potassium Isopropoxide in Isopropyl Alcohol-d at 36°

a 6 d HC10 CH C1Br 10 k Time 4 iso-PrOK 2 Percent -1 -1 Seconds ml.b M Reaction L. mole sec.

0 43.25 0.4920 0.4830 5280 36.88 0.4195 0.4468 14.7 32.5 10260 32.20 0.3663 0.4202 25.5 32.0 13680 29.60 0.3367 0.4054 31.6 31.5 15720 28.20 0.3208 0.3974 34.8 31.2 18720 26.25 0.2986 0.3863 39.3 31.1 21900 24.42 0.2778 0.3759 43.5 30.9 25500 22.52 0.2562 0.3651 47.9 30.8 28080 21.28 0.2421 0.3581 50.8 30.8 30420 20.20 0.2298 0.3519 53.3 30.8 c 105840 4.30 0.0601 0.2671 87.8 30.1

Avg. 31.2 ±0.05 a b c 0.04450 M. Per 4.0 ml. sample. Mi. of 0.05588 M HC10 4 per 4.0 ml. d sample. Calculated from equation (5). 190

Table 33

Reaction of Methylene Chlorobromide with . Potassium tert-Butoxide in tert-Butyl Alcohol at 36°

a HC10 C13 C1Br 106k Time 4 tert-BuOK 2 Percent -1 -1 Seconds ml.b Reaction . mole sec. e 0 0.4917 0.4815 180 35.14 0.4884 0.4799 0.7 38.9 g 2280 34.51 0.4797 0.4755 2.4 11.3 g 4140 33.80 0.4698 0.4706 4.5 11.6 g 6120 33.04 0.4592 0.4653 6.6 11.8 9000 32.04 0.4453 0.4583 9.4 11.7 12780 30.83 0.4285 0.4499 12.9 11.6 16860 29.40 0.4087 C).4400 16.9 11.9 23400 27.40 0.3809 0.4261 22.5 12.1 30480 25.66 0.3567 0.4140 27.5 11.8 c 47400 33.55 0.3052 0.3883 37.9 11.7 d 5300 10.15 0.2828 0.3771 42.5 12.0

Avg. 11.8 ±0.1 a b c 0.05560 M. Per 4.0 ml. Sample. An additional 5.0 ml. of 0.1298 M d e NaOH added. Ml. of 0.05560 M HC10 per 2.0 ml. sample. Extrapolated. 4 (Calculated from equation (5). gOmitted from average, < five percent reaction. 191

Table 34

Reaction of Methylene Chlorobromide with Potassium tert7Butoxide in tert-Butyl Alcohol-d at 36°

a b d HC10 CH2C1Br 10 k Time 4 tert-BuOK Percent -1 -1 Seconds ml.b M M Reaction t. mole sec. c 0 0.4713 0.3863 e 1560 34.60 0.4602 0.3808 2.4 19.9 3600 33.50 0.4455 0.3734 5.5 20.6 5400 32.63 0.4340 0.3677 7.9 20.3 11100 30.01 0.3991 0.3502 15.3 20.4 16920 27.70 0.3684 0.3349 21.8 20.3 23280 25.55 0.3398 0.3206 27.9 20.0 28800 23.83 0.3169 0.3091 32.8 20.0 48720 19.10 0.2540 0.2777 46.1 19.6 85440 13.67 0.1818 0.2416 61.4 18.8 103680 11.73 0.1560 0.2287 66.9 18.6 116600 10.40 0.1383 0.2198 70.7 18.8

Avg. 19.7 ±0.6

c d a0..05:313M. bPer 4.0 ml. sample. Extrapolated. CaltUlated,,from e equation (5). Omitted from average,'< five percent reaction. 192

Table 35

Reaction of Methylene Chloride with Sodium Methoxide in Methyl Alcohol-d at 36°

a 6 e HC10 CH Cl 10 k Time 4 Me0Na 2 2 Percent -1 -1 Seconds ml.b M M Reaction t. mole sec. d 0 0.5633 2.7408

3660 42.20 0.5611 2.7397 0.4 0. . 195 E 7500 42.02 0.5587 2.7385 0.8 0.200 E 12420 41.08 0.5558 2.7370 1.3 0.197 E 96180 38.28 0.5090 2.7136 9.6 0.193 c 182640 36.70 0.4656 2.6919 17.3 0.192 c 267240 33.90 0.4301 2.6742 23.6 0.187 c 353760 31.17 0.3954 2.6568 29.8 0.186 c 432660 28.80 0.3654 2.6418 35.1 0.186 c 524040 26.62 0.3377 2.6280 40.0 0.182 615120 24.72 ° 0.3136 2.6159 44.3 0.178 c 706000 22.98 0.2915 2.6049 48.3 0.17.5

Avg. 0.185 ±0.005 a b c 0.05319 M. Per 4.0 ml. sample. Ml. of 0.05075 M HC104 per 4.0 ml. sample. dExtrapolated. e Calculated from equation (5). tOmitted from average, < five percent reaction. 193

Table 36

Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 e HC10 CH 2 C1 10 k Time 4 iso-PrOK 2 Percent -1 -1 Seconds ml.b M M Reaction Z. mole sec.

d 0 0.4333 2.4594 5400 32.23 0.4286 2.4570 1.1 0.411 f 11160 31.79 0.4228 2.4541 2.4 0.447 f 20220 31.20 0.4150 2.4502 4.2 0.435 86760 27.30 0.3628 2.4241 16.3 0.419 117360 25.65 0.3409 2.4132 21.3 0.420 198240 21.13 0.2809 2.3832 35.2 0.452 c 284700 19.92 0.2527 2.3691 41.7 0.393 396360 17.29 ° 0.2193 2.3524 49.2 0.358 455160 15.02 ° 0.1905 2.3380 56.0 0.378 534800 13.12 c 0.1664 2.3259 61.6 0.376 627080 11.35 ° 0.1440 2,.,3147 66.8 0.370

Avg. 0.396 ±0.26 a 0.05319 M. bPer 4.0 ml. sample. cM1. of 0.65075 M HC10 4 per 4.0 ml. sample. dExtrapolated. e Calculated from equation (5). fOmitted from average, < five percent reaction. 194

Table 37

Reaction of Methylene Chloride with Potassium Isopropoxide in Isopropyl Alcohol-d at 36° a 6 c HC10 CH 2 C1 10 Time 4 iso-PrOK 2 Percent -1 -1 Seconds ml. b M M Reaction Q. mole sec.

0 34.63 0.4393 2.3611 85980 26.57 0.3371 2.3100 23.3 0.660 115560 24.60 0.3121 2.2975 29.0 0.636 140880 23.10 0.2931 2.2880 33.3 0.619 172080 21.43 0.2719 2.2774 38.1 0.602 202140 19.98 0.2535 2.2682 42.3 0.589 257760 17.59 0.2232 2.2530 49.2 0.571 277560 16.65 0.2112 2.2470 51.9 0.574 313260 15.41 0.1955 2.2392 55.5 0.564 345660 14.32 0.1817 2.2323 58.6 0.558 368220 13.50 0.1713 2.2271 61.0 0.560

Avg. 0.593 ±0.029

a b c . 0.05075 M. Per 4.0 ml. sample. Calculated from equation (5).

1 195

Table 38

Reaction of Methylene Chloride with Potassium tert-Butoxide in tert-Butyl Alcohol at 36°

a 6 e HC10 CH Cl 10 k Time 4 tert-BuOK 2 2 Percent b -1 -1 Seconds ml. Reaction Z. mole sec. d 0 0.4037 2.6636 6420 30.08 0.4000 2.6617 0.9 0.269 E f 10500 29.91 0.3977 2.6606 1.5 0.268 16620 29.64 0.3941 2.6588 2.4 0.272 E 107820 26.30 0.3497 2.6366 13.4 0.251 c 194280 24.63 0.3125 2.6180 22.6 0.250 278880 22.25 c 0.2823 2.6029 30.1 0.244 c 355340 20.04 0.2542 2.5888 37.0 0.241 434180 18.28c 0.2319 2.5777 42.6 0.244 c 535500 16.60 0.2106 2.5670 47.8 0.233 606700 15.10 c 0.1916 2.5575 52.5 0.236 c 696940 13.88 0.1761 2.5498 56.4 0.229

Avg. 0.241 ±0.006 a b c 0.05319 M. Per 4.0 ml. sample. Ml. of 0.05075 M HC1011 per 4.0 ml. sample. dExtrapolated. e Calculated from equation (5). tOmitted from average, < five percent reaction. 196

Table 39

Reaction of Methylene Chloride with Potassium tert-Butoxide in tert-Butyl Alcohol at 36°

a 6 c HC10 CH C1 10 k Time 4 tert-BuOK 2 2 Percent -1 -1 Seconds Ml. b M M Reaction Q. tole sec.

0 35.,39 0.4026 2.7122 98280 29.15 0.3316 2.6767 17.6 0.366 204320 19.48 0.2216 2.6217 44.9 0.549 212880 19.10 0.2173 2.6196 46.0 0.544 267540 17.35 0.1973 2.6096 51.0 0.502 290460 16.66 0.1895 2.6057 52.9 0.489 375420 14.40 ' 0.1638 2.5928 59.3 0.453 445200 12.88 0.1465 2.5842 63.6 0.430 537240 11.22 0.1276 2.5747 68.3 0.407 611460 10.12 0.1151 2.5685 71.4 0.390

Avg. 0.458 ±0.057 a 0.05319 M. bPer 4.0 ml. sample. 'Calculated from equation (5). 197

Table 40

Reaction of Chloroform with Sodium Methoxide in Methyl Alcohol at 36°

a 6 c HC10 10 k Time 4 Me0Na HCC1 Percent b 3 -1 Seconds ml. Reaction Z. mole -1 sec.

0 27.40 0.3828 0.6531 d 1560 27.12 0.3789 0.6518 3.36 9300 25.38 0.3546 0.6457 4.23 20040 23.20 0.3241 0.6335 15 4.31 24060 22.20 0.3101 0.6289 19 4.56 32640 20.73 0.2896 0.6220 24 4.48 103260 11.90 0.1662 0.5809 56 4.41 109920 11.41 0.1594 C.5786 58 4.36 166800 8.01 0.1119 C.5628 71 4.11

Avg. 4.35 ±01l a c 0.05588 M. bPer 4.0 ml. sample. Calculated from equation (9). d Omitted from average, < five percent reaction. 198

Table 41

Reaction of Chloroform with Potassium Isopropoxide in Isopropyl Alcohol at 36°

a 6 c HC10 HCC1 10 k Time iso-PrOK 3 Percent 4b Seconds ml. Reaction mole

0 7.41 0.01036 0.02295 2700 4.33 0.00604 0.02151 42 3.01 4320 3.06 0.00428 0.02092 59 3.14 5520 2.45 0.00342 0.02064 66 3.11 6480 2.04 0.00285 0.02045 73 3.10 7920 1.57 0.00219 0.02023 79 3.08 9300 1.18 0.00165 0.02005 84 3.13 10500 0.95 0.00133 0.01994 87 3.12

Avg. 3.10 ±0.03 a -3 b c 5.588 x 10 M. Per 4.0 ml. Sample. Calculated from equation (9). •

Table 42

Deuterium Exchange of-Methylene Bromide in Methyl Alcohol-d at 36°'

e 6 g HC10 10 k Time Cell Pa Me0H 4 Me0Na Percent -1 1 Seconds Absorbance cm. x 10 M ml. M_ Reaction Z. moles sec. a 0 0.220 6.89 0.244 40.76 0.5694 a 660 0.225 0.249 1980 0.229a 0.253 3960 0.231a 0.254 a 7260 0.244 0.268 1.2 5.80 a 11160 0.258 0.282 1. 5. .81 16500 0.269 0.291 2.3 5.25 17220 37.21 0.5198 b L 24480 0.286 0.309 3.1 4.96 b 33900 0.320 0.344 4.8 5.86 34680 34.00 0.4750 c 79560 0.425 0.453 10.1 6.13 805 80 27.20 0.3800 c 93720 0.455 0.483 11.5 6.26 c 99420 0.467 0.495 12.1 6.34 c 107400 0.480 0.506 12.7 6.28 112680 23.68 0.3308 d 114540 0.492 0.514 13.0 6.14 170040 0.578d 7.17 0.602 17.3 6.5.9 195120 16.80 0.2347

Avg. 5.95 ±0.37

e, 133; d aExtinction coefficients interpolated from Figure c, 131; bc, 132; c E 134. e 0.05588 M. f e P r 4.0 ml. sample. gCalculated from equation (13) where: a=1.0250, b=0.560, c=0.875, m=0.900. Table 43

Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d at 36°

d 6 e- HC10 10 Time Cell Path MeOH 4 Me0Na, Percent -1 -1 Seconds Absorbance cm. x 10 3 M ml. - M Reaction Q. moles sec. f 0 0.5620 180 0.189 6.24 0.229 3600 19.45 0.5538 15780 0.220a 6.26 0.266 2.4 6.31 18600 18.40 0.5239 b 63060 0.300 6.39 0.353 8.2 5.86g 65760 15.52 0.4419 100680 13.83 0.3938 b 103000 0.375 6.5 1 0.433 13.4 6.56 148440 0.426 c 6.56 0.485 16.8 6.31 150720 11.76 0.3349 184860 10.56 0.3007 185880 0.465 6.57 0.528 19.7 6.49 c 237600 0.495 6.52 0.567 22.2 6.29 240720 9.08 0.257 c 271920 0.531 6.58 0.602. 24.5 6.22 274860 8.32 0.237

Avg. 6.36 ±0.10

e, 133; aExtinction coefficient interpolated from Figure 5: 6, 132; b C E, 134. d0.1139 M. ePer 4.0 ml. sample. e Calculated from equation (13) where: a=0.7540, b=0.5620, c=0.888, m=0.874. (Extrapolated. gOmitted from average, > four times the average deviation. Table 44

Deuterium Exchange of Methylene Iodide in Methyl Alcohol-d at 36°

d 6 g HC10 10k Time Cell Pal Me0H 4 Me0Na Percent -1 Seconds Absorbance cm. x 10 ml.e Reaction 2. moles sec.

0 0.248 0.4560E a 240 0.222 6.77 0.248 a h 7980 0.253 6.77 0.283 4.0 2.25 9840 15.42 0.4391 a 15180 0.306 6.73 0.344 10.9 3.49 b 22740 0.351 6.80 0.388 15.9 3.57 b 33420 0.391 6.87 0.428 20.4 3.61 : b 38880 0 433 6.98 0.466 24.8 3.59 46680 13.60 0.3872 b 47220 0.472 6.92 0.513 30.1 3.84 87540 12.43 c 0.3539 98020 0.608 6.97 0.651 45.8 3.551 c 1 120840 0.705 7.03 0.748 56.8 4.45 127800 11.60 0.3303

Avg. 3.62 ±0.09

aExtinction coefficients interpolated from Figure 5 e, 132; b e, 133; CE, 134. d0.1139 M HC104 . e Per 4.0 ml. sample. From dilution factors. gCalculated from equation (13) where: a=0.4397, b=0.4560, c=0.2657, d=0.6321. hOmitted from average, > four times the average deviation. 10mitted from average,> 35 percent reaction. Table 45

Deuterium Exchange of Methylene Chloride in Isopropyl Alcohol-d at 36°

c 5 e HC10 10 k Time iso-PrOH 4 iso-PrOK Percent a b -1 -1 Seconds Absorbance ' ml.'' M Reaction ft. moles sec.

0 0.305 0.331 3.60 0.101 1260 0.309 0.335 0.2 3.12 f 2100 0.309 0.335 0.2 1.87 E 5940 0.320 0.347 0.7 2.50 E 12600 0.344 0.373 1.9 3.10 20340 0.368 0.399 3.1 3.13 27720 0.379 0.411 3.6 2.71 28380 3.44 0.0961 32760 0.403 0.437 4.8 3.07 36180 0.414 0.449 5.4 3.11 84120 0.530 0.574 11.1 2.91 86400 3.29 0.0919 94320 0.551 0.597 12.1 2.87 107700 0.572 0.620 13.1 2.76 180600 0.710 0.769 19.9 2.68 183780 0.712 0.772 20.0 2.65 194640 2.96 0.0827 198360 0.770 0.834 22.9 2.89 205620 0:750 0.812 21.9 2.64 210600 0.780 0.845 23.4 2.80 285120 2.53 0.0707

Avg. 2.87 t0.15

b 3 aExtinction coefficients interpolated from Figure 6. Cell-path length, 5.39 x 10 cm. c 5.59 x 10 -2 M. dPer 2.0 ml. sample. e Calculated from equation (12) where: a=1.1018, f b=0.101, c=0.0349. Omitted from average, < one percent reaction. Table 46

Deuterium Exchange of Methylene Chlorobromide in Isopropyl Alcohol-d at 36°

c 4 f HC10 10 k Time iso-PrOH 4 iso-Pr9K Percent a b -1 Seconds Absorbance ' M mi . d M x 10 Reaction L. mole. sec.

0 0.298 0.333 2.13 5.94 600 0.305 1200 0.299 2400 0.308 0.344 34 80 0.321 0.359 5280 0.339 0.379 4.2 3.27 6900 0.352 0.393 5.6 3.60

9240 0.364 0.407 6.8 3.63 ' 14340 0.386 0.431 9.1 4.00 17340 0.387 0.433 9.2 3.39 20520 0.403 0.450 10.8 4.63 22260 0.410 0.458 11.6 5.95g

Avg. 3.75 1 0.37

b -3 aExtinction coefficients interpolated from Figure 6. Cell-path length, 5.39 x 10 cm. c d e 0.05588 M. Per 2.0 ml. sample. The base concentration during the remainder of the -6 reaction was palculated from the rate constant for formal formation, i.e., 31.2 x 10 . g (Calculated from equation (12) where: a=0.5418, b=0.06200, c=0.4730. Omitted from average, > four times the average deviation. Table 47

Deuterium Exchange of Methylene Bromide in Isopropyl Alcohol-d at 36°

3 e HC10 10 k Time iso-PrOH iso-PrOK Percent a b -1 Seconds Absorbance , ml. M x-10 3 Reaction Z. moles sec.

0 0.329 0.344 3.43 7.82 1800 0.366 0.394 2.3 3.48 3480 0.396 0.423 4.2 3.37 4140 3.00 6.83 4980 0.420 0.452 5.8 3.29 10740 0.504 0.543 11.1 3.24 11760 2. 54 5.87 16140 0.574 0.618 1 5.5 3.31 16800 2.20 5.00 23400 0.665 0.716 21.3 3.60

Avg. 3.38 ±0.10

-3 aExtinction coefficients interpolated from Figure 6. b Ce11.-path length, 5.45 x 10 cm. c -3 d e 4.55 x 10 M. Per 2.0 ml. sample. Calculated from equation (12) where: a=0.8509, b=0.00775, c=0.0100. Table 48

Deuterium Exchange of Methylene Bromide in Isopropyl Alcohol-d at 36°

c 3 e HC10 10 k Time iso-PrOH 4 iso-PrOK Percent a b -1 Seconds Absorbance ' mi.d M x 10 3 Reaction Z. moles sec.

0 0.251 0.280 4.25 11.9

720 0.267 0.297 0.8 1.95E 1440 0.278 0.310 1.4 1.67 2640 0.299 0.333 2.5 1.66 4500 3.62 10.1 4800 0.335 0.373 4.4 1.72 6480 0.360 0.401 5.6 1.66 7920 0.375 0.418 6.3 1.58 8460 3.19 8.92 10800 0.414 0.461 8.3 1.62 12840 2.82 7.88 13200 0.442 0.492 9.8 1.62 15600 0.466 0.519 11.0 1.60 17880 0.490 0.546 12.2 1.62 18720 2.37 6.53 22080 0.524 0.584 13.9 1.60 23460 0.532 0.593 14.4 1.58 24120 1.98 5.56 26100 0.550 0.613 15.3 1.57 29400 0.573 0.638 16.5 1.58

Avg. 1.62-±0.03

aExtinction coefficients interpolated from Figure 6, bCell-path length, 5.28 x 10 -3 cm. c 5.58 x 10-3 M. dPer 2.0 ml. sample. e Calculated from equation (12) where: a=1.0903, f b=0.01185, c=0.01940. Omitted from average, < one percent reaction. Table 49

Deuterium Exchange of Methylene Iodide in Isopropyl Alcohol-d at 36°

c HC10 10 3ke Time iso-PrOH 4 iso-PrOK Percent -1 -1 Seconds Absorbance a ' b M ml.d M x 10 3 Reaction R. molessec.

0 0.269 0.297 2.64 7.37 f 600 0.312 0.344 1.4 0.59 1020 0.355 0.392 4.5 1.18 1440 0.370 0.408 5.6 1.05 1860 0.398 0.439 7.7 1.12 2220 0.422 0.466 9.4 1.17 2700 0.447 0.493 11.3 1.16 3060 2.52 7.04 3240 0.475 0.524 3840 0.504 0.556 15.5 1.1c 4320 0.525 0.579 17.0 1.13 5040 U.559 0.617 19.5 1.13 5580 2.38 6.65 5820 0.598 0.660 22.4 1.14 6420 0.622 0.686 24.2 1.13 7080 0.651 0.718 26.3 1.13 7860 0.679 0.749 28.4 1.12 8460 0.693 0.765 29.4 1.08 8940 2.21 6.17 9360 0.728 0.803 31.9 1.08 10260 0.752 0.830 33.7 1.06 10740 2.15 6.01 10980 0.772 0.852 35.2 1.04

Avg. 1.12 i0.03 a b Extinction coefficients interpolated from Figure 6. Cell-path length, 5.53 x 10 -3 cm. -3 c5.59 x 10 M. dPer 2.0 ml. sample. eCalculated from equation (12) where: a=0.7463, f b=0.0768, c= 0.00323. Omitted from average, < one percent reaction. Table 50

Deuterium Exchange of Methylene Chloride in tert-Butyl Alcohol-d at 36°

c 3 e HC10 10 k Time tert-BuOH 4 tert-BuOK Percent a b -1 -1 SecondS Absorbance ' ml.d M x 10 2 Reaction Q. moles sec.

0 0.367 0.462 8.43* 1.92 f 960 0.380 0.478 0.8 0.83 1740 0.396 0.498 1.7 1.02 2280 0.407 0.512 2.3 1.07 3120 0.427 0.537 3.5 1.19 3600 8.37 1.90 4200 0.441 0.555 4.3 1_09' 6360 0.479 0.602 6.5 1.11 7920 8.42 1.91 8280 0.505 0.635 8.0 1.06 9720 0.524 0.659 9.1 1.03 12900 0.569 0.716 11.7 1.02 15420 8.30 1.89 16020 0.614 0.772 14.3 1.02 18180 0.652 0.820 16.5 1.05 20160 0.686 0.863 18.5 1.07 21720. 0.700 0.881 19.3 1.04 21960 8.20 1.86 24540 0.737 0.927 21.5 1.04 28380 0.783 0.985 24.1 1.03

Avg. 1.06 1 0.03 a b -3 Extinction coefficients interpolated from Figure 72 Cell-path length, 5.30 x 10 cm. c -3 d e 4.55 x 10 M. Per 2.0 ml. sample. Calculated from equation (12) where: a=1.0847, f b=0.0192, c=0.00122. Omitted from average, < one percent reaction. Table 51

Deuterium Exchange of Methylene Chlorobromide in tert-Butyl Alcohol-d at 36°

d 3 f HC10 10 k Time tert-BuOH 4 tert-BuOK Percent a,b -1 -1 Seconds Absorbance ml.e M x 10 Reaction R. moles sec. c c 0 0.365 0.472 8.33 c 180 0.370 0.479 0.5 6.86g 900 0.395 0.511 3.1 8.42 1440 0.410 0.530 4.6 8.00 2340 6.17 7.83

2400 0.440 0.570 7.7 8.23 • 3120 0.456 0.590 9.3 7.80 4380 0.488 0.632 12.6 7.75 5460 0.514 0.665 15.3 7.73 6600 0.536 0.694 17.5 7.51 9240 5.55 7.04 10080 0.643 0.832 28.5 9.00 11940 0.623 0.806 26.4 6.89 12540 0.633 • 0.819 27.5 6.89 14580 0.661 0.856 30.3 6.77 17040 0.694 0.898 33.7 6.70 18780 0.707 0.915 35.0 6.43 18800 4.85 6.15

Avg. 7.55 ±0.63

b 3 aExtinction coefficients interpolated from Figure 7. Cell-path length, 5.15 x 10 cm. c d -3 e Extrapolated. 5.08 x 10 M. Per 4.0 ml. sample. (Calculated from equation (12) where: a=0.6318, b=0.00833, c=0.00478. gOmitted from average, < one percent reaction.

Table 52

Deuterium Exchange of Methylene Chlorobromide in tert-Butyl Alcohol-d at 36°

3 f HC10 d 10 k Time tert-BuOH 4 tert-BuOK Percent -1 Seconds Absorbance a ' b ml. e M x 10 3 Reaction R. moles sec. -1

0 0.377c 0.469 c 60 3.55 9.02 300 0.386 0.480 0.9 6.66g 1080 0.412 0.512 3.5 7.43 1560 0.428 0.532 5.0 7.64 2220 0.446 0.555 6.8 7.43 3180 3.04 7.72 3960 0.492 0.612 11.4 7.38 4740 2.87 7.29 5580 0.527 0.655 14.8 7.18 6740 7140 0.550 0.684 17.1 6.69 8040 2.67 6.78 8940 0.589 0.733 21.0 6.95 10380 G.601 0.747 22.2 6.44 12660 0.634 0.788 25.4 6.40 14880 2.48 6.30 15300 0.672 0.836 29.2 6.50 17100 0.685 0.852 30.5 6.22 19200 0.706 0.878 32.5 6.16 21300 0.730 0.908 34.9 6.25 24180 0.755 0.939 37.4 6.21h h 26280 0.770 0.958 38.9 6.14 h 28740 0.779 0.969 39.8 5.86

Avg. 6.82 i0.45

aExtinction coefficients interpolated from Figure 7. hCell-path length, 5.36 x 10 -3 cm. cExtrapolated. d5.08 x 10 -3 M. ePer 2.0 ml. sample. (Calculated from equation (12) where: a=0.6287, b=0.00900, c=0.00984. gOmitted from average, < one percent reaction. hOmitted from average, > 35 percent reaction. Table 53

Deuterium Exchange of Methylene Bromide in tert-Butyl Alcohol-d at 36°

c 2 e HC10 10 k Time tert-BuOH 4 tert-BuOK Percent a b -1 -1 Seconds Absorbance ' M ml.d M x 10 Reaction x. moles sec.

0 0.361 0.463 0.42 9.6 420 0.372 0.477 0.8 3.90 E 9n0_ 0.379 0.486 1.3 2.99 1680 0.403 0.517 2.9 3.81 2580 0.426 0.546 4.5 3.90 3540 0.437 0.560 5.3 3.35 5160 0.474 0.608 7.9 3.51 5940 0.37 8.4 7200 0.519 0.665 11.0 3.64 9120 0.548 0.706 13.0 3.48 10980 0.582 0.746 15.4 - 3.51 12000 0.34 7.8 12420 0.607 0.778 17.1 3.52 14100 0.632 0.810 18.8 3.49 16200 0.657 0.842 20.6 3.39 17820 0.670 0.859 21.5 3.25 18900 0.33 7.8 19800 0.697 0.894 23.4 3.26 21900 0.714 0.915 24.5 3.14 22740 0.31 7.1

Avg. 3.45 1 0.18 a b -3 Extinction coefficients interpolated from Figure 7. Cell-path length, 5.20 x 10 cm. -3 d e x 10 M. Per 2.0 ml. sample. Calculated from equation (12) where: a=0.9222, f b=0.000940, c=0.000494. Omitted from average, < one percent reaction. Table 54

Deuterium Exchange of Methylene Bromide in tert-Butyl Alcohol-d at 36°

d 2 f HC10 10 k Time tert-BuOH 4 tert-BuOK Percent a b - 1 -1 Seconds Absorbance ' M ml.e M x 10 3 Reaction Z. moles sec.

0 0.493 c 120 0.65 1.65 240 0.396 0.480 1440 0.436 0.528 3.1 2.68 1920 0.442 0.536 3.7 2.43 2700 0.453 0.549 4.9 2.29 3120 0.470 0.570 6.7 2.74 3780 0.486 0.589 8.4 2.86 4500 0.501 0.607 10.0 2.89 5100 0.63 1.60 6420 0.543 0.658 14.5 3.02 7320 0.555 0.673 15.4 2.91 8400 0.573 0.694 17.7 2.88 9720 0.600 0.727 20.5 2.95 10260 0.61 1.55 11220 0.622 0.754 22.9 2.90 12840 0.660 0.800 26.9 3.08 14160 0.675 0.818 28.5 2.99 16140 0.688 0.834 29.9 2.78 17760 0.707 0.857 31.9 2.75 18480 0.60 1.52 19380 0.703 0.876 33.6 2.69 21180 0.741 0.898 35.5 2.64g 23580 0.755 , 0.915 37.0 2.51g 25440 0.777 0.942 39.3 2.52g 26040 0.59 1.50

Avg. 2.80 *0.17

aExtinction coefficients interpolated from Figure 7. bCell-path length, 5.50 x 10 -3 cm. cExtrapolated. d5.08 x 10 -3 M. e Per 2.0 ml. sample. fCalculated from equation (12) where: a=0.5707, b=0.001643, c=0.000351. gOmitted from average, > 35 percent reaction. Table 55

Deuterium Exchange of Methylene Iodide in tert-Butyl Alcohol-d at 36°

c 2 e HC10 10 k Time tert-BuOH 4 tert-BuOK Percent a b -1 -1 Seconds Absorbance ' mi . d M x 104 Reaction J. moles sec.

0 0.401 0.514 0.40 9.10 540 0.448 0.574 6.1 2.54 900 0.487 0.624 11.1 2.88 1560 0.521 0.678 15.5 2.39 2100 0.39 8.87 2340 0.571 0.732 22.0 2.37 2820 0.610 0.782 27.0 2.51 3420 0.644 0.826 31.5 2.49 3900 0.37 8.41 4200 0.682 0.874 36.4 2.45

4860 0.712 0.913 40.3 2.42E

5520 0.745 0.955 44.5 2.45 E 6300 0.37 8.41

6600 0.772 0.990 48.0 2.29 E

7380 0.792 1.015 50.6 2.23 E 7740 0.37 7.96 f 8640 0.822 1.054 54.5 2.13 11820 0.33 7.50

Avg. 2.52 1 0.08

aExtinction coefficients interpolated from Figure 7., bCell-path length, 5.17 x 10 3 cm. c 3 d e 4.55 x 10 M. Per 2.0 ml. sample. Calculated from equation (12) where: a=0.4952, b=0.00912, c=0.000170. (Omitted from average, > 35 percent reaction. H N.) 213

Table 56

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Iodide at 36°

a 6 c HC10 CH2 I 2 10 k Time 4 Me0Na Percent -1 Seconds ml.b M M Reaction Q. mole sec. -1

0 48.30 0.5494 0.5401 83640 38.77 0.4410 0.4859 19.7 2.57 96120 37.70 0.4288 0.4798 21.9 2.54 145740 34.00 0.3867 0.4588 29.6 2.43

167040 32.48 0.369, 5 0.4502 32.7 2.42 252000 27.72 0.3153 0.4231 42.6 2.33 340680 23.77 0.2697 0.4003 50.9 2.28

424020 20.80 0.236 , 6 0.3837 56.9 2.22 491100 18.87 0.2146 0.3727 60.9 2.18 593040 16.20 0.1846 0.3577 66.4 2.16 696660 11.49 0.1307 0.3308 76.2 2.56

Avg. 2.37 -10.14 a b c 0.04550 M. Per 4.0 ml. sample. Calculated from equation (5). 214

Table 57

Reaction of Methylene Iodide with Sodium Methoxide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Iodide at 36°

a 6 c HC10 10 k Time 4 Me0Na CHI Percent2 -1 -1 Seconds ml.b M M Reaction Z. mole sec.

0 48.88 0.5994 68520 41.31 0.5066 0.4592 15.5 2.55 156540 33.14 0.4064 0.4091 32.2 2.74 246960 28.31 0.3472 0.3795 42.1 2.55 343440 24.03 0.2947 0.3532 50.8 2.48 443700 20.76 0.2546 0.3332 57.5 2.40 590160 18.00 0.2207 0.3163 63.2 2.18 774360 15.00 0.1839 0.2979 69.3 2.05 952320 13.01 0.1595 0.2857 73.4 1.92

Avg. 2.36 1 0.22

a b c 0.04905 M. Per 4.0 ml. sample. Calculated from equation (5). 215

Table 58

Reaction of Methylene Iodide with Sodium Methoxide in Methanol-d in the Presence of 0.5 M Sodium Perchlorate at 36°

a 6 c HC10 I 10 k Time 4 NeONa CH 2 Percent2 -1 Seconds ml. b Reaction Z. mole sec.

0 40.40 0.5644 0.4500 60300 34.42 0.4808 0.4132 14.8 3.07 84000 32.88 0.4593 0.4024 18.6 2.87 168000 28.21 0.3939 0.3697 30.2 2.62 239460 25.30 0.3534 0.3495 37.4 2.47 332940 22.16 0.3096 0.3276 45.1 2.36 404220 20.15 0.2815 0.3135 50.1 2.31 517080 17.50 0.2445 0.2950 56.7 2.26 601020 15.90 0.2221 0.2838 60.6 2.22 696900 14.30 0.1998 0.2727 64.6 2.19

Avg. 2.49 ±0.25 a b c 0.05588 M. Per 4.0 ml. sample. Calculated from equation (5). 216

Table 59

Reaction of Methylene Iodide with Sodium Iodide in Methyl Alcohol-d in the Presence of 0.5 M Sodium Perchlorate at 36°

6 c HC10 a CHI 10 k Time 4 Me0Na 22 Percent -1 -1 Seconds ml.b M M Reaction 9- mole sec.

0 48.35 0.5929 65340 41.80 0.5126 0.4605 13.5 2.32 154140 34.80 0.4267 0.4176 28.0 2.34 243720 29.84 0.3659 0.3872 38.3 2.27 340440 25.35 0.3109 0.3596 47.6 2.26 440340 21.83 0.2677 0.3380 54.8 2.23 585660 18.62 0.2283 0.3184 61.5 2.10

771060 14.83 0.1 . 819 0.2951 69.3 2.07 -949080 12.30 0.1508 0.2796 74.6 2.03

Avg. 2.21 ±0.10 a c 0.04905 M. bPer 4. . 0 ml. sample. Calculated from equation (5).

217

Table 60

Stoichiometry of the Reaction of Potassium tert-Butoxide with Methylene Bromide

b CH Br tert-Bu0Ka Final Vol. Aliquot HC10 AgNO 3 c A[OH] mi 2 ml. ml. ml. ml. ml. Z577f-

5.0 25.0 100.0 20.0 30.57 8.40 .944 25.0 38.00 10.55 .950

5.0 25.0 50.0 10.0 6.94 18.70 1.040 10.0 6.95 18.84 1.030

a b c 0.4600 M. 0,04905 M. 0.1010 M.

Table 61

Stoichiometry of the Reaction of Potassium Isopropoxide with Methylene Chlorobromide

b CH C1Br iso-PrOKa Final Vol. Aliquot HC10 AgNO 3 c A[OH- ] 4 mI ml. ml ml. ml. ml. A[C]

5.0 25.0 50.0 10.0 21.92 14.37 1.00

a b c 0.5052 M. 0.04905 M. 0.1010 M. 218

1,0

400

Figure 1. Attempted Determination of the Rate of Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d by Nuclear Magnetic Resonance Spectroscopy.

219

■ ■ 100 .====aimesim iwz IMIMM MiMMEMEM NNIIIMIIMIHNIIIIMIINMMOMMIUMPINIMMINISM.OMP23•51111.C12.030 Effinumffirk.I®® EH moo .11111111 1.1=arrAMWAS 80 Iri MINIMINEEEINEENENffillMEIMENEN1 IlL 1111M=INENNIINMEN ENE F 1 IMINIMIEMENMEMBENE NE IN INNIMENEENNIMIELI 111 NNESINII EINEENEEI

IMENNIENIIIIMIN i 40 Mill =IMMENSE= ENMENENINNI = INE `i20 11•11=11==== El INEMME111=== =NE 10 12 13 14 15

Figure 2. Infrared Spectrum of Disopropyl Formal (Neat).

10 0 707 0.0 I-

p 0.2 it

0.4

0.6

0.8 11.0 1.5 1 I ...a.. 3 4 5 6 7 9 10 11 12 11 1 4 15

Figure 3. Infrared Spectrum of Di(tert-Butyl) Formal (Neat).

1 220

5000 4000 3000 10,111111 la IIIIIIIIIIImmill III i 111111 11111111111111111111 11111111411 11111111111111111111

11111111.11111111111.111111111111 111111111111111111111 11111111...... 1., A IA 1111111111111 MI 11111111 I RI 13111111111111 111111 """", ''- 1I 11111111111111 11 11II 1.1111111111111111111411111111111111111 1111111m :11 BIElliENIONNMIll IIIIMMIM li M IiIIIINCEINIONNONINIM IONNEIN g IMMENEMONNIONONOMENI 0 2 NEMMEEMM a ENIMMEMOMMONIMMIEMMIN MINIM' RIO 11.1111111:11011EIMMNI NINON= qmplimi MINENIONOMOO NUMMI 11011mmEmmlil NOMMMOMMMON ■■MENE■MMOW■IIMMI■ OMEN= POO E MOM :MM NMEMMEMI MEN■MIIMM■M■ ■■MIMI■■■ M mumnimmummi mom mum • 11 sinimn umucom m MUNERENN • • INNIUMNRMIONWEENUN IIIMMI ■■■ RRI ON ENNOIIIRMIINNINNME MOINMINNE: I ■ MENNE HEONI NONNI= MMUS= • MN MIMIEVEREIMINMENUMMI MENUMMEN NI N Ell ■■■■ INI■A■■■■■■■■ =MUNN MINN M • IMMIr•••••••••• NMEEMM•m Mi MN MILIMNEW1 WOMMIIMMEIMEM NMENMMUM • MN M IMMIMEIMNMNI:MME MENEM= • • :MUIRI; IMRE ENIEN MU MMUNE=1 • NM MU IR NIMOIMMMIMMENIMMU mwmommm m mom= UMMEMOM IMUMMIMM mmommomm . mommumplimmulimmommilm 0.4 MEMMUMEMME MITIMINIAll UMMEMOMMEM MMMEMMUMMI iii MEXIMMA 111111•MMII UMMUMMIEM ! ME 11111111MMEM IMIMm MMMMMM MMOMMMEM I ON MMMMM EMMEN Emma mingummarm MIMS I ..m. "" mom ....m.. IIII . ..*I12 1 . ..121"mil Imimpli iiiiiiiii■■ ■■■h■■■ ■■■■■■ NM NEMO ■ INI :I AME0 MEMNON■ ■■■■ON MONNEIN■■ IN NMO normNOMMEN' MImomM mommaimmummM ORM immumummumm m mu NMENREN • MI BRONNVANNEHMMNIMM NNE 0.6 MEMO= R • 11•••••••••/1171111 NE= =MUNE MN NRMILINUMMMUM MINIMUM iii iii 1■M NNW MUNNIMMINIM■MN= milimwm.01 mmummiumm•PP 0 moms mommismwm MUmi mommovimmimmommum mommmum • • ism mummummummommim a MMMMM imm m mu mg mimmiramitm= N. g. 1 mum1 mimmw,.mgm MMMMM ml 0.8 mil I. AffillEfigsms...... amm! mem= 1656:=BEF----pai 1 0 limill il isilimmillmilli 11111111 II 11111111111111111111 11 MOON= EN II INOMEMENINEOMMONEENEN 0 Imams ■• ■■■■mpliumgiu ■■■■m =k"I .5 M NNW I fililtilluallmil (MICRONS)

Figure 4. Infrared Spectrum of Methyl Alcohol in Methyl Alcohol-d. Determination of the Absorbance of the 0-H Absorption by the Empirical Ratio Method. 221

140

130

120

110

100 0.4 0.8 1.2

Me0Na, M

Figure 5. The Effect of the Concentration of Sodium Methoxide on the Extinction Coefficients of Methyl Alcohol in Methyl Alcohol-d.

222

170

160

150

140 0 0.1 , 0.2

i•PrOK, M

Figure 6.... The Effect of the Concentration of Potassium Isopropoxide on the Extinction Coefficients of Isopropyl Alcohol in Isopropyl Alcohol-d. 223

150

140

130

120

110

100 0 0.1 0.2 0.3 0.4 0.5 0.6

tert-BuOK, M

Figure 7. The Effect of the Concentration of Potassium tert-Butoxide on the Extinction Coefficient of tert-Butyl Alcohol in tert-Butyl Alcohol-d.

2214

7.0 T sT T

Figure 8. Nuclear Magnetic Resonance Spectrum of Di(tert-Butyl) Formal (Neat).

10 9A 4.0 5;0 PPM M.. 6.0 1.0 9A - L 1-1.0• 0 CPS

Figure 9. Nuclear Magnetic Resonance Spectrum of Diisopropyl Formal (Neat). 225

0.4

0.3

p 0.2

0.1

0 1 2 3 -5 Time, seconds x 10

Figure 10. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d. Data Shown in Table 43. 226'

0.6

0.5

0 . 4

0.3

0.2 0 1 2 3 -5 Time, seconds x 10

Figure 11. Calculation of the Rate of Constants for Deuterium Exchange of Methylene Bromide in Methyl Alcohol-d. Data Shown in Table 43. 227

0 x

0 1 2 3 -5 Time, seconds x 10

Figure 12. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methanol-d. Data Shown in Table 43. 228

0 0 .1 0.2 0.3

Figure 13. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methanol-d. Data Shown in Table 43.

229

0.6

0.5

z ro 0. 14 a)

0.3

0.2 0 0.1 0.2 0.3

Figure 14. Calculation of the Rate Constants for Deuterium Exchange of Methylene Bromide in Methenol•d. Data Shown in Table 43. 230

1.2

1 .1

1.0

0. '1

0.6

0.5 0 0.1 0.2 0.3 0.4

Figure 15. Calculation-of- the Rate Constants for Deuterium. Exchange.. of Methylene Bromide in Isopropyl Alcohol-d. Data shown in Table 48.

'1 231

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31. J. Hine, B. C. Menon, J. H. Jensen, and J. Mulders, J. Am. Chem. Soc., 88, 3367 (1966).

32. Lars Melander, Isotope Effects on Reaction Rates, The Ronald Press Co., New York, N.Y., 1960, (a) Chapters 4 and 6; pp: (b) 125-6.

33. Organic Syntheses, Coll. Vol. I, John Wiley and Sons, 1956, p. 3357.

34. R. B. Duke, unpublished results.

35. Sadtler Standard Spectra Index (1966), Spectra Nos: (a) 5939(CH 2Br2 ), C1 ) 4636(CH2 I ) 904(CH ClBr); (b) 9685. 1011(CH 2 2 2' - 2 36. C. E. Meloan, Elementary Infrared Spectroscopy, The Macmillan Co., New York, N.Y,, 1963, p. 169.

I 11 233

37. Connecticut Instrument Co., Instruction Manual for Infrared Sampling Accessories, Wilton, Conn., p. 17.

38. H. A. Flaschka, EDTA Titrations, Pergamon Press, New York, N.Y., 1959, p. 117.

39. P. K. Knoefel, Journal of Pharmacology, 50, 88 (1934).

40. I. Jansson, Suomen Kemistilehti, 34-p, No. 78, 83 (1961).

41. H. Normant and C. Crisan, Bulletin la societe dhemique de France, 199 (1959).

42. R. Leimu, Suomen Kemistilehti, 19-B, 66 (1946).

43. S. B. Lippincott and M. M. Lyman, Industrial and Engineering Chemistry, 38, 320 (1946).

44. K. Nukada, Bulletin of the Chemical Society of Japan, 33, 1606 (1960). PART III

THE REACTION OF METHYLMAGNESIUM BROMIDE WITH BENZOPHENONE-- THE MECHANISM OF THE GRIGNARD REACTION 234

CHAPTER I

INTRODUCTION

Background

Since the first reports by Barbier (1, 2) and Grignard (3) that

alkyl halides react with magnesium to give a reagent useful for a variety

of synthetic purposes, considerable effort has been expended by numerous

workers to determine the nature of the reactive organometallic species

in solution and the reaction mechanism describing these facile trans-

formations. Although these investigations have brought out many inter-

esting and important aspects concerning the reagent, the nature of the

reactive moiety and the details of the reaction mechanism are still

unsettled.

The Reactive Grignard Species

In regard to the nature of the reactive species, it was originally

thought that the Grignard reagent could be described simply by its

empirical, formula, RMgX; however, Schlenk and Schlenk (4) showed that

addition of dioxane to a Grignard reagent prepared in the usual manner

precipitates a magnesium halide as the dioxanate leaving in solution the

corresponding dialkylmagnesium compound.

2 RMgX F R Mg + MgX2 2 (1)

To explain their.results, Schlenk and Schlenk proposed an equilibrium between the Grignard reagent and the dialkylmagnesium compound; the 235

equilibrium could, of course, :be shifted to the right by removal of the magnesium halide (by precipitation with dioxane).

From these observations, it is evident that either RMgX or R 2 Mg may be the reactive species. On the other hand, both may be reactive, the normal Grignard products being formed by two parallel reactions. No experimental evidence has been presented which distinguishes between these possibilities. It can be independently demonstrated that dialkylmagnesium compounds react with ketones and nitriles to give normal

Grignard addition products; however, because of equilibrium (1), it is impossible to prepare RMgX in the absence of the corresponding equilibrium quantity of R2 Mg. As a result, it is impossible to study the reactions of RMgX.completely independent of R 2 Mg.

Association of the Grignard Reagent

another factor which tends to obscure the reactive species in solution and to further complicate kinetic studies of Grignard reactions is that the organomagnesium compounds form associated species. Although it was originally reported (5, 6) that Grignard reagents were predominantly dimeric in solution, more recent investigations (7, 8, 9) have shown that the degree of association is a function of the concentration of the reagent, association increasing with increasing concentration of the organomagnesium compound. The nature of the associated species is unknown; however, several structures have been proposed (10-15). It was originally thought that the most highly associated species was the dimer,

R Mg X More recent investigations (16), however, have shown that 2 2 2. association proceeds at least as far as the trimer stage and perhaps 236

beyond. As a result, many of the earlier kinetic investigations which purportedly support either RMgX or as the reactive species are R2Mg2X2 now open to question because, without exception, these studies were carried out at Grignard concentrations at which both species were undoubtedly present. The earlier workers, therefore, had several species with which to contend. As a result, mathematical treatment of their data became essentially impossible and indeed led to various trends in the rate constants which were never explained adequately.

The Grignard-Ketone Complex

Another problem to contend with in a study of the kinetics of the reaction of Grignard reagents with ketones is the formation of complexes between the reacting species. The presence of Grignard-ketone complexes was first suggested by color changes observed upon addition of the organomagnesium compounds to certain ketones. Although earlier workers

(17-21) thought , that they had isolated stable, ether-insoluble complexes,

Nesmayanov (22) showed that the majority of these were actually the expected magnesium halide carbinolates to which a molecule of ketone was complexed. However, much of the controversy over the existence of stable, ether-insoluble complexes has recently been revived by Cowan and

Dolak's report (23) of a fenchone-phenylmagnesium bromide complex which they isolated and characterized.

Smith (24-26) has offered spectroscopic evidence for the existence of ether-soluble complexes formed by the reaction of methylmagnesium bromide and 2,4-dimethy1-4 --thiomethyl benzophenone. These complexes were assumed to be intermediates in the addition of the Grignard reagent to the ketone. There is evidence, therefore, that both ether-soluble and

237

ether-insoluble Grignard-ketone complexes exist and additional evidence

that they are reaction intermediates; however, the stoichiometry of these

complexes has not been determined and in only one case has the equilibrium

constant been evaluated.

In order to derive the appropriate rate equations for following

the Grignard reaction, it is essential to know both the stoichiometry of

the complex and the equilibrium constant governing its formation. In the

absence of such information, certain assumptions must be made. Since

most workers have assumed, a 1:1 complex, we employed the same value in

the derivation of our kinetic equations. In addition, experiments were

designed so that the equilibrium constant, if it were large, could be

evaluated.

Kinetics of the Grignard Reaction

Several workers (28, 29, 30, 33) have examined the kinetics of the

reactions, of Grignard reagents with ketones and nitriles. The conclusions

as to the appropriate rate laws governing these transformations are quite

varied even when identical systems were studied. For example,'Bikales and

Becker (29) report that the reaction of methylmagnesium bromide with

benzophenone follows the rate law (2):

v = k[D][K] (2)

where D is the Grignard dimer, R 2 Mg2 X2 , and K is the ketone. In contrast,

Anteunis (33, 34) reports that the same reaction follows the rate law (3):

v = k[G] 2 [K] (3)

where G is the Grignard monomer, FMgX. Unfortunately, the evidence for 238

both rate laws is not as convincing as would be desired. The rate law proposed by Bikales and Becker correlates only the first 20 to 30 percent of their data after which the rate constants decline markedly. On the other hand, the rate law proposed by Anteunis gives consistent values for the rate constants throughout the meaningful portion of the reaction but the rate constants are dependent upon the initial concentration of the reactants.

Another anomaly appears in the reaction of Grignard reagents with nitriles. Storfer and Becker (30) maintain that the reaction of ethylmagnesium bromide with benzonitrile follows the rate law (4) which is similar to that proposed for the reaction of benzophenone with methylmagnesium bromide.

v = DaPhONJ (4)

On the other hand, Swain (28) found that the reaction of'n-butylmagnesium bromide with benzonitrile followed the rate law (5).

v = [G][PhCN] (5)

Again, the results, particularly those of Becker, are not as convincing as some would like.

Reaction Mechanisms

Several different mechanistic proposals have emerged from the kinetic studies of the reactions of ketones and nitriles with Grignard reagents. These differ primarily with respect to the nature of the reactive organometallic species in solution and the order of the reaction with respect to this moiety. With only a single exception (27), all of 239

the previous investigators have concluded that the reaction is first-

order with respect to the ketalecr nitrile. At this point, however,

agreement ends. In order to present the aforementioned mechanistic con-

cepts, some of the earlier contributions will be discussed briefly.

The Meisenheimer Mechanism. Meisenheimer (17-19) seems to have

been the first to suggest that Grignard reagents react with ketones to

form complexes which then rearrange to the observed carbinolates in the

rate-controlling step of the reaction.

K + G F C (6)

C 4- P

Smith and Su (25, 26) interpreted their results from rate studies on the

reaction of 2,4-dimethy1-4'-thiomethyl benzophenone in terms of mechanism

(6). In addition, Swain (28) also reported that the reaction of n-butylmagnesium bromide with benzonitrile in ether was first-order in monomeric Grignard reagent and first-order in nitrile, second-order

over-all which follows if the Grignard-nitrile complex is present in low

concentration with respect to the reactants and the steady-state approxi-

Mation can be applied. Swain proposed that the reaction proceeded by the

following mechanism;

Bu 4 Ph-CEN + BuMgX Ph-CEN...Mg-X

Bu Bu -k I Ph-C=N...Mg-X Ph-C=N-MgX

Although Smith and Su did not propose a detailed mechanism, they did

240

envoke an intermediate Grignard-ketone complex to explain their results.

The Dimer Mechanism. Frcm their study of the reaction of methyl-

magnesium bromide with benzophenone in tetrahydrofuran, Bikales and

Becker (29) proposed that the reaction proceeds by the following path:

R R N X 1 4 1 1 C=0 + R Mg Mg 2 2 X2 ,,,.,C\ . !// R 0 \\X 1

/,' Mg R R Mg R - -0MgX + RMgX • N 1 I R 0 X Rl 1

Storfer and Becker (30) employed, the same mechanism to explain the reac-

tion of ethylmagnesium bromide with benzonitrile, as did Dessy and

Salinger (31), to account for the results from the study of the reaction

of ethylmagnesium bromide with substituted benzalanilines (Schiff bases).

The Swain Mechanism. From a study of, the effects of added magnesium

bromide on the product distribution, i.e., addition, reduction, and

enolization, in, the reaction of n-propylmagnesium bromide with diisopropyl

ketone, Swain and Boyles (32) suggested the following reaction mechanism:

241

R K 11 CO= + R MgX ...Mg-X 1 R

z Mg 11 k T1 .Mg R-C-0MgX + R MgX " 1 0 ° R

In support of the Swain mechanism, Anteunis (33, 34) found that the

reactions of methylmagnesium bromide and iodide with benzophenone and

pinacolone were first-order in ketone and second-order in monomeric

Grignard reagent, third-order over-all.

Purpose

In an effort to resolve some of the uncertainties regarding the

nature of the reactive Grignard species in solution and hopefully to

distinguish between the numerous mechanisms proposed for the Grignard

reaction, a study of the kinetics of the reaction of methylmagnesium

bromide with benzophenone was undertaken.

Approach

Examination of the conditions employed by earlier workers (28-30,

33) in their investigations of the kinetics of the Grignard reaction showed

that the reaction had always been studied at concentrations at which the

organometallic reagent was largely associated. Since the nature of the

associated Grignard reagent is uncertain, it is impossible to examine

earlier data and to establish the relative reactivities of the various

species. To eliminate dealing with a multiplicity of species, the addition 242

of methylmagnesium bromide to benzophenone was studied under conditions where the organomagnesium compound was largely monomeric. For methylmagnesium bromide, this can be accomplished by studying the reaction at -2 concentrations below 2 x 10 M.

Methylmagnesium bromide and benzophenone were chosen as reactants in order to eliminate the undesirable side-reactions,i.e., reduction and enolization, often accompanying the normal Grignard addition process, benzophenone being incapable of enolization and methylmagnesium bromide being incapable of reduction. 243

CHAPTER II

EXPERIMENTAL

Apparatus

Instrumentation

A Cary Model 14 recording spectrophotometer was used in conjunction with a matched set of quartz, 10-mm. cells (Beckman No. 75191) for making all of the ultraviolet measurements.

Constant-Temperature Bath

A constant-temperature water bath (Sargent No. S-84845) and a

Sargent Thermonitor (Sargent No. S-82055) were used for controlling the temperature ( 1 0.01°).

Timer

An electric stopwatch, (Sargent No. S-77490) was used.

Inert Atmosphere Box

Kewaunee, Model 2C1020, was employed.

Reaction Flasks

Special reaction flasks were designed and constructed for use in the kinetic studies. These were 25-ml., glass-stoppered volumetric flasks with a glass sidearm between the stopper and the graduation mark.

The sidearm was of the same diameter as the neck of the flask (ca., 10 mm.) and was about 20 mm. in length. It was tilted downward so that it formed an angle of about 60 degrees with the neck of the flask. The glass stoppers were carefully ground to fit each individual flask and both the 244

stoppers and flasks labeled to insure a proper fit. The stoppers had small glass hooks protruding from one side which allowed the stoppers to be held securely in the flasks by rubber bands extending between the hooks and the sidearms of the reaction flasks.

Chemicals

Benzophenone

Eastman's reagent was recrystallized from either pentane or isooctane prior to use.

Diethyl Ether

Baker's reagent (anhydrous) was distilled from lithium aluminum hydride prior to use.

Methyl Bromide

Eastman's reagent was distilled prior to use.

1,1-Diphenylethanol

Eastman's reagent was recrystallized from pentane prior to use.

Magnesium

Ingots of sublimed magnesium (courtesy of the Dow Chemical Co.) were milled into turnings using a Carballoy cutting tool with special care being taken to prevent contamination. The turnings were washed with ether, dried in vacuo, and stored in an atmosphere of dry nitrogen.

Manganese Carbonate

Baker Analyzed.

Oxalic Acid

Baker Analyzed.

Ammonium Chloride

Baker Analyzed. 245

Methods for Following the Grignard Reaction

Various physical methods have been employed for following the

Grignard reaction. These include 1) spectroscopic determination of the

unreacted ketone either in an ethereal solution of the Grignard reagent or after hydrolysis (26, 29, 33); 2) gas evolution methods for determin-

ing the amount of unreacted Grignard reagent (28); 3) dielectric constant measurements (36); 4) gravimetric analysis of the reaction products,

i.e., ketones derived from the reaction of nitriles with Grignard reagents

as the corresponding 2,4-dinitrophenylhydrazones (30); and 5) gas chroma-

tographic analyses of the reaction products after hydrolysis (33).

The method most often employed for studying the reaction of

Grignard reagents with ketones is that of following the disappearance of the carbonyl chromophore in ethereal solutions of the Grignard reagents.

This method, although direct, is quite limited in scope. Ethereal solutions of Grignard reagents begin absorbing strongly below 290 mu and

0.01 M solutions are usually opaque (in 10-mm. cells) below 270 mu. As a result, the high extinction absorptions of common carbonyl compounds cannot be employed. To circumvent this difficulty, two alternative approaches have been used. The first employs a longer wavelength absorption band of the carbonyl compound. Unfortunately, these absorptions have much lower extinction coefficients (ca., 200) and require that the

concentration of the carbonyl compound be in the order of 0.1 M or greater to produce adequate response from the spectrophotometer. If comparable concentrations of the Grignard reagent are employed, consideration of the association of the organomagnesium compounds in the mathematical treatment of the data is necessary. Since the latter involves gross uncertainties, 246

this approach should be avoided if possible.

The second approach employs substituted ketones which have high extinction absorptions above 290 mp. Benzophenone, for example, has its

A at about 250 mp. The substituent, therefore, must be capable of max shifting the X 50 mp or more to longer wavelengths. Smith and Su (25, max 26) found that a thiomethyl group attached to a diaryl ketone would effect the desired shift. The drawback to this approach involves the question of whether the functional group complexes with the Grignard reagent. Both ethers and carbonyl compounds have been shown to form such complexes. The uncertainties added by employing a functional group which could presumably enter into complex formation would certainly weaken any quantitative arguments that could be derived from the resulting kinetic data; consequently, this approach should also be avoided if possible.

Another spectroscopic method for following the reaction of Grignard reagents with ketones was developed by Anteunis (33). In studying the reaction of methylmagnesium bromide and iodide with benzophenone, Anteunis followed the disappearance of the ketone after first hydrolyzing the reaction mixture. This method, although it requires separate reactions for each point in the kinetic study, markedly broadens the scope of the analysis. For example, when diethyl ether is employed as solvent, the analytical range is extended down to 235 mp. This modification permits the study of most diaryl ketones at their A near 250 mp. In addition, max the high extinction coefficients of the absorptions (E = 10 4 ) allow the -4 reaction to be studied at low ketone concentrations (ca. 10 M) which are advantageous for two reasons: 1) the lower ketone concentrations 247

-4 allow correspondingly lower Grignerdrconcentrations (ca. 143 . M) to be employed which assures that the association of,the Grignard reagent is

minimized and 2) the lower concentrations also slow the rate of the , reaction and minimize the uncertainties in the measurement of time.

The experimental procedure developed during the course of this study for following the reaction of benzophenone with methylmagnesium bromide was modeIad after that of Anteunis (32) in that analysis of unreacted ketone was effected after hydrolysis; however, the apparatus

employed and , the experimental details are somewhat different as will be discussed later.

KineticStudies

Preparation of the Reaction Flasks and Addition of Benzophenone

The reaction flasks were dried by heating with.the flame of a

Bunsen burner while they were clamped in an upright position and purged with dry, oxygen-free nitrogen or argon (by means of a syringe needle extending to the bottom of each flask). After cooling to room temperature

(with the inert gas still flowing into the flasks), an aliquot of a standard solution of, benzophenone was introduced into the sidearm of each flask by means of a calibrated Hamilton syringe. The quantity employed in these studies was 10 (10 x 10 liters). After transfer of the ketone, the stopper of the reaction flask was dried (by heating in an open flame), lightly greased (with silicone stopcock,grease), and then placed in the reaction flask while still hot. Stoppering was done quickly after removal of the syringe needle used to purge the flask to prevent contamination of the flask and the ketone with either oxygen or moistures. 248

The reaction flasks were then clamped in a rack which was sized so

that it could be taken into the inert atmosphere box through the entry

port. The rack held each flask firmly in an upright position. This

facilitated the transfer of the Grignard reagent to the flasks (in the

inert atmosphere box) and eliminated the danger of tipping the flasks

over and causing premature mixing of the reactants.

The reaction flasks (containing benzophenone) and 250 ml. of dry,

freshly-distilled ether (from lithium aluminum hydride) were then placed

in the inert atmosphere box. The syringes, needles, etc., necessary for

the subsequent transfers were also placed in the box at this time.

Normally, this was done 12 to 16 hours in advance of the subsequent trans-

fers of the Grignard reagent. During this period the box atmosphere was

recirculated through the purification system to insure that it was dry

and oxygen free.

Preparation of Methylmagnesium Bromide Solutions

An aliquot of a standard solution of methylmagnesium bromide

(ca.,l-2 M) was transferred to the freshly distilled ether by means of a

calibrated syringe and thoroughly mixed. The dilute solutions of methyl- -2 4 magnesium bromide (ca.,10 to 10 M) so prepared were clear and color-

less and showed no signs of cloudiness. An aliquot (2 to 10 ml.) of the

dilute Grignard reagent was then transferred by a syringe to the lower portion of each reaction flask taking care to prevent the organomagnesium

compound from getting into the sidearm containing the benzophenone. The

ground-glass stoppers were then securely seated in the necks of the flasks

and fastened in place with, rubber bands. After the transfers were complete, the reaction flasks and the remainder of the dilute methylmagnesium 21+9

bromide solution employed in the reaction were removed from the inert- atmosphere box.

Initiating, Timing, and Quenching of the Reactions

All but two of the reaction flasks were placed in a constant- temperature bath at 25°. Just the lower bulbs of the flasks were immersed leaving the necks and. the ground glass stoppers above water in order to carry out the subsequent operations. The contents of the two remaining reaction flasks, one of which did not contain benzophenone, were hydrolyzed with ten percent ammonium. chloride. The sample containing the benzophenone, served as a measure of the initial ketone concentration at zero time and the other was employed in the reference beam of the spectrophotometer to compensate for the absorbance of the solvent.

After adequate time had elapsed for the reactants to come, to thermal equilibrium, the reactions were initiated by repeated inversions of the flasks. For the short reaction times, ca., 1000 seconds or less, only one reaction was run at a time and the electric stopclock was started at the same time the reactants were mixed. When longer reaction times were necessary, the reactions were all run simultaneously by starting the reactions at 50 second intervals. The reactions were quenched by adding 5 ml. (accurately measured) of a ten percent solution of ammonium chloride solution from a syringe. With the short reaction times, the stopclock was stopped the moment the syringe plunger was depressed. After some experience with this experimental technique, the entire operation of opening the reaction flask, quenching the reaction mixture, and stopping the stopclock could be accomplished in a second or less. The time the contents of the flasks were in contact with 250

the atmosphere was minimized. The turbulence created by forcing the ammonium chloride solution into the flask from the syringe seemed to be sufficient to completely quench the reaction; however, in order to insure complete hydrolysis, the flasks were immediately restoppered, removed from the constant-temperature bath, and shaken.

Analyses

Benzophenone. The hydrolyzed samples were diluted to volume with ether and shaken for several minutes to assure that equilibrium was established between the ether and aqueous layers of the two-phase system.

Both phases, ethereal and aqueous, were clear and colorless. The ethereal layer of each individual sample was in turn transferredto a 10-mm, quartz cell and the absorbance recorded at the minimum in the absorption spectrum

at 308 mu and the maximum at 25 .1 mu. A sample of the hydrolyzed Grignard reagent was employed in the reference beam of the spectrophotometer.

After the absorbance of each individual sample was recorded, the quartz cell was washed first with water, then with acetone, and finally with ether. Rinsing with water was necessary to prevent the clouding of the cells presumably due to deposits of water-soluble salts on the cell walls.

Methylmagnesium Bromide. The volumetric flask containing the remainder of the methylmagnesium bromide solution employed in the kinetic study was first weighed and then treated with a measured excess of standard perchloric acid. Each of the volumetric flasks employed in the study was individually tared so that the weight of the ethereal solution of the Grignard reagent could be easily ascertained. From the density of ether at room temperature (i.e., 0.710 g/ml.) and the weight of the

Grignard reagent, the volume of the methylmagnesium bromide solution was 251

determined.

The volumetric flask containing the hydrolyzed Grignard reagent was placed on a steam bath and the ether evaporated. After the ether layer had disappeared, the aqueous solution was heated for an additional

30 minutes; otherwise, the endpoints tended to fade. The volumetric flask was then removed from the steam bath and the aqueous solution titrated with standard sodium hydroxide (while still warm) to a phenolphthalein endpoint. The titrations were carried out by direct addition of the base to the solutions in the volumetric flasks. From the titration data and the volume of the ethereal solution, the concentration of the Grignard reagent was calculated.

Extinction Coefficients

The extinction coefficients of benzophenone and diphenyl methyl carbinol were determined by preparing standard solutions in diethyl ether and then recording the ultraviolet spectra. The extinction coefficients

were calculated from equation (7) ::

A E: = (7) tl where E is the extinction coefficient, 1 is the cell-path length in centimeters, c is the molar concentration of the substrate, and A is the absorbance. The absorbance was measured from the minimum in the absorption spectrum at 308 mp to the maximum at 251 mp. The extinction coefficients of benzophenone and diphenyl methyl carbinol (at 251 mil) 4 were found to be 1.84 x 10 and 342, respectively. Both values are in good agreement with those reported by Bikales and Becker (29). The 252

extinction coefficient of Biphenyl methyl carbinol was not determined at its X but at the X of benzophenone. This was done to correct max max for the amount of the alcohol formed during the course of the reaction.

Preparations

Methylmagnesium Bromide

The desired amount of sublimed magnesium (plus a 50 percent excess) was weighed into a round-bottom flask containing a glass-covered, magnetic stirring bar. The reaction flask was connected to a dry-ice condenser and arranged so that diethyl ether and methyl bromide could be distilled ditectly into it without opening the flask or apparatus to the atmosphere. The reaction flask and apparatus were then flame-dried under vacuum and subsequently filled with argon. This process was repeated three times to assure a dry, oxygen-free atmosphere. The desired amount of diethyl ether (under argon) was then distilled (from lithium aluminum hydride) into the reaction flask. When this was complete, the dry-ice trap was filled, the magnetic stirrer was started, and methyl bromide was distilled (through Drierite) directly into the reaction flask. The reaction initiated almost immediately and the rate was controlled by the rate of distillation of the halide. The amount of methyl bromide added was determined either by weighing the vial containing the halide or by observing the amount of unreacted magnesium remaining. The latter was usually employed, the reaction being terminated when about 60 to 70 percent of the magnesium had been consumed.

After addition of the methyl bromide was complete, the flask containing the Grignard reagent was removed from the apparatus, stoppered, 253

and stored in the inert-atmosphere box. Methylmagnesium bromide so prepared was clear and water-white and no color changes were evidenced after six months. No effort was made to separate the Grignard reagent from the unreacted magnesium. Samples of the reagent were withdrawn from the supernatant liquid when necessary by means of a syringe.

Inert-Atmosphere Box and Purification System

All transfers of the methylmagnesium bromide were made in an inert-atmosphere box (Keuwanee,Model No. 2C1020) equipped with an evacuable entry port. Nitrogen (oil pumped) was employed as the inert gas. The box was equipped with a recirculating system which moved the box atmosphere at a rate of 0.5-1.0 cubic feet per minute. The recirculating pump (Gelman Instruments, No. 13152) was mounted on a specially designed steel plate and then sealed into a five-gallon solvent can to prevent external leakage of the pump (Figure 1). Experience has shown that all pumps leak and that if this precaution is not taken, no reasonably sized oxygen scrubbing system can keep up with the oxygen entering the system through the pump. A piece of copper tubing connected the pump enclosure with the inert atmosphere box so that the demand of the pump leakage was met with a supply of dry, oxygen-free nitrogen.

Purification System

The purification system employed in connection with the recirculating system was essentially that proposed by Brown (35). The exit from the box was passed first through the pump and then through a dry-ice trap, then through a column of 4A molecular sieves, and finally through a column of manganese(II) oxide. The dry-ice trap removed the majority of 2514

1. IDE OX SI ) DRY-ICE II II) ( TRAP SE NE LECULAR NESE( A MO NGA A MANG M 4A

NI,TROGEN VENT THROUGH MINERAL olLEIX4ER t INERT-ATMOSPHERE BOX

PUMP AND ENCLOSURE 0

Figure 1. Inert-Atmosphere , Box and Recirculating Purification System. 255

the solvent vapors (namely ether). The ethers unavoidably get into the inert

atmosphere box each time a flask is opened or a transfer is made.

Although Brown employed a column of 13X molecular sieves for the same

purpose, we have found that the dry-ice trap is satisfactory. Some type

of scavenging system is necessary to remove the ethers and other organic

compounds from the recirculating system because of their deleterious

effects on the manganese(II) oxide. If this is not done, the Mn0 becomes

difficult to regenerate and finally must be discarded. The dry-ice trap

also removed considerable amounts of water which entered the inert

atmosphere box by diffusion through the gloves.

From the dry-ice trap, the atmosphere of the box was passed through

a column of 4A molecular sieves to remove the residual rater. In continu-

ous operation, the molecular sieve column was regenerated once a week;

however, it was evident from the amount of water produced during these

regenerations that only a fraction of the capacity of the sieves was

being utilized and that much longer periods between regenerations could

have been employed.

From the molecular sieve column, the atmosphere of the box was passed through a column packed with manganese(II) oxide (on vermiculite).

This column served to remove residual oxygen. Manganese(II) oxide is a bright-green oxide which reacts pyrophorically with oxygen to give higher

oxides of manganese which are brown in color. As a result, the manganese oxide not only serves to remove traces of oxygen in the system but also

serves as its own indicator.

Preparation of Manganese (II) Oxide

The manganese oxide was preparea by pyrolysis of manganese oxalate. 256

The oxalate was prepared by dissolving one pound (3.59 moles) of oxalic

acid dihydrate in two liters of water and heating to 70-80°. Manganese

carbonate, 400 g. (3.48 moles), was then added in small increments with

good mixing. This reaction is best carried out in a hood. The insoluble. manganese oxalate thus formed was washed twice by decantation with hot water and then vacuum filtered. The precipitate was washed until the

filtrate was neutral. The cream-colored solid was subseq4ently dried overnight in an oven at 100°. If higher temperatures are employed, the oxalate will gradually decompose to the black manganese oxide. The black oxide can be employed equally as effectively as the oxalate as far as preparing the Mn0 column is concerned but the latter is more convenient

to handle.

After drying, the oxalate was dry-slurried with a sufficient

amount of vermiculite to fill a column 90 cm. long and 10 cm. in diameter.

The mixture was then transferred to the column. The manganese oxide and

the molecular sieve columns were wired so that they could be heated to

temperatures between 300° and 350°. Hydrogen was then passed through the

column (from top to bottom) while heating to 350°. The oxalate decomposed at these temperatures as evidenced by a change from a light-brown oxalate color to the bright-green color of manganese(II) oxide. After pyrolysis, the column is ready for use.

Two manganese(II) oxide columns were employed in the gas purification system. They were arranged so that a fresh (green) column was in the recirculating system at all times. When one of the columns became spent (as evidenced by a change in color from green to black) the recirculating system was directed through the other and the spent column was 25:f regenerated. Regeneration was achieved by passing hydrogen through the column at 350°. Experience with the manganese oxide columns showed that months of continuous service could be expected from a single column if the recirculating system was properly trapped to remove organic solvents.

Pperatlahof the Inert-Atmosphere Box

The key to success in the operation of an inert-atmosphere box is to have the recirculating system running at all, times. Our system was operated continuously except for a shut don of several hours once a week to regenerate the molecular sieve column. Contamination of the box atmosphere with traces of oxygen and" moisture occurs mainly by permeation through the gloves and by taking equipment in and out of the box through the entry port. Of the two, the latter is the greatest offender. Even with the most careful manipulation of the entry port, successive entries into the box can spend a manganese oxide column in a few hours. For this reason, when crucial operations, such as the handling and transfer of the

Grignard reagents were undertaken., all of the necessary reagents and equipment were taken into the box the evening before the transfers were to be made and the actual transfers were made the following morning.

This procedure allowed about twelve hours for the recirculating system to purify the box atmosphere. The box atmosphere, therefore, was as pure as the system was capable of producing. 258

CHAPTER III

DISCUSSION AND RESULTS

Kinetic Studies

The reaction of methylmagnesium bromide with benzophenone was studied first at Grignard-to-ketone ratios ranging from 25 to 150/1. The

results of these studies (Tables 4 through 23, Appendix) are shown in

Tables 1 and 2. Under these conditions, the disappearance of the ketone

was found to be pseudo - first - order. The first - order rate constants were

calculated from equation (8):

2.303 A k = t log 17l obs (8) c where A is the corrected absorbance of'benzophenone at time, t, and A c l is the initial absorbance of benzophenone as determined by extrapolation.

The value of A was calculated from equation (9), perivation 1, Appendix: c

A-(0.0188)A 0 0.982 (9) where A is the initial Absorbance of benzophenone and A is the measured 0 absorbance at time, t. The value of A was determined by plotting A l c versus time on semi-log paper and extrapolating to zero time as shown in

Figure 2.

The value of A was always less than A in the pseudo-first-order l 0 studies. The reason for this is that when the reactants are initially 259

Table 1. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at High Grignard-to- Ketone Ratios ,at 25°

Methylmagnesium b 102k c Bromidea obs k -1 -1 -1 Benzophanona sec. P. mole sec.

22.9 0.251 i 0.014 1.09 ± 0.06 22.9 0.315 * 0.018 1.37 ± 0.08 22.9 0.302 ± 0.022 1.31 ± 0.09 23.3 0.244 ± 0.021 1.06 ± 0.09 45.7 0.558 ± 0.016 1.22 ± 0.03 45.7 0.587 ± 0.020 1.27 ± 0.05 91,5 1.65 ± 0.05 1.77 i 0.06 91.5 1.48 ± 0.05 1.59 ± 0.05 91.5 1.24 ± 0.05 1.34 ± 0.06 93.3 1.24 ± 0.03 1.35 ± 0.04 93.3 1.40 ± 0.04 1.51 ± 0.05 aConcentration determined by dilution. b Calculated from equation (8). cCalculated from equation (10), where n=1 260

Table 2. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at High Grignard-to- Ketone Ratios at 25°

Methylmagnesium 2 Bromide, 10 k kc obl -1 Benzophenone sec -1 t. mole sec.

23.2 0.263 ± 0.010 1.07 ± 0.04 23.9 0.256 ± 0.015 1.08 ± 0.06 24.5 0.308 ± 0.024 1.25 ± 0.09 44.7 0.709 ± 0.036 1.57 ± 0.08 47.8 0.781 ± 0.029 1.67 * 0.07 94.0 1.42 ± 0.10 1.49 ± 0.11

95.4 1.45 ± 0.04 1.50 ± 0.04 - 143 2.62 ± 0.09 1.73 ± 0.06 143 2.33 ± 0.06 1.60 ± 0.04 a Concentration determined by direct analysis by the Gilman method, b Calculated from equation (8). c Calculated from equation (10), where n=1. 261

0.9

0.8

0.7 ,

0.6

0.5 A c

0.4

0.3

0.2 0 100 200

Tithe, seconds

Figure 2. Calculation of Pseudo-First-Order Rate Constants for the Reaction of Methyl Magnesium Bromide with Benzo- phenone. Determination of Initial Absorbance, Al . 262

mixed, the concentration of ketone in the sidearm of the flask is consid- erably higher than an instant later when mixing is complete. The reaction, therefore, is much faster during the first second or two than later when -3 mixing is complete. Because the reaction is fast (at 9.5 x 10 M -4 methylmagnesium bromide and 1 x 10 M benzophenone, the reaction has a half-life of about 30 to 35 seconds), as much as ten percent of the ketone can be consumed before dilution is complete. Although there appeared to be no correlation between the amount of benzophenone consumed and the initial concentration of methylmagnesium bromide, the repeat- ability from sample to sample within a given kinetic run was good as evidenced by the precision of the rate constants.

Our explanation for the fact that A o and Al were not identical under pseudo-first-order conditions is, of course, open to question; however, further evidence that our explanation is indeed correct is evidenced by the fact that as the Grignard concentration is lowered to that of the ketone, Al approaches A0 . In our later studies, where the concentrations of the reactants were both comparable, i.e., both about -4 5 x 10 M, Ao was employed as the initial concentration of the ketone.

The order of the reaction with respect to the Grignard reagent under pseudo-first-order conditions was examined with the aid of equation (10):

k A l obs 2.303 k = log -A— (10 ) Go tGo 263

where G is the initial concentration of methylmagnesium bromide and n is o an integer, 1, 2, etc. It was found that the most consistent sets of values for the rate constants were obtained when n was one indicating that the reaction was first-order in the Grignard reagent. The second- order rate constants calculated from equation (10) are shown in Tables 1 and 2.

The rate constants reported in Table 1 were calculated employing

Grignard concentrations as determined by dilution of standard , solutions of methylmagnesium bromide; those reported in Table 2 were calculated from the concentrations of the organometallic compound as determined by direct analysis (Gilman method) of the dilute Grignard reagents. The reproducibility at a , given concentration of the Grignard reagent is better in the latter case showing that direct analysis of the dilute solutions is to be preferred.

Examination of the results obtained under pseudo-first-order conditions shows that the second-order rate constants tend to decrease with decreasing Grignard-to-ketone ratios. To verify this trend, we studied the reaction at Grignard-to-ketone ratios ranging from 1.4 to 6.9/1.

Under these conditions, the concentration of the methylmagnesium bromide cannot be treated as a constant. The data (Tables 24 through 28,

Appendix) from these studies were analyzed with the second-order rate equation (11):

K (G -EP]) 2.303 , o o = T77o1K ) (K -[P]) ( o log G o 264

where K and G represent the initial concentrations of benzophenone and o o methylmagnesium bromide, respectively. In these studies, the experimental quantities, A and A , were assumed to be equal to K and K (the amount o c o of ketone at time t), respectively, and P is diphenyl methyl carbinol formed in the reaction as calculated from equation (12), the stoichiometry relationship after hydrolysis.

EP] = - [K] (12)

As the results show (Table 3), the average second-order rate constants range between 0.36 and 0.72 as compared to values of 1.0 to 1.7 obtained, at Grignard-to-ketone ratios of 25 to 150/1. The trends observed in the pseudo-first-order studies are indeed real; the second- order rate constants definitely decrease with decreasing Grignard-to- ketone ratios.

In contrast, if second•order rate constants are calculated from the pseudo-first-order rate constants reported by Smith and Su (26) (for the reaction of methylmagnesium bromide with 2,4-dimethy1-4 --thiomethyl benzophenone), the values decrease with increasing concentration of the

Grignard reagent. Although the two sets of data appear to be contradictory, they nevertheless compliment one another. Smith and Su obtained their data at methylmagnesium bromide concentrations ranging from -2 4 x 10 M to 1.5 M. As shown in Figure 3 (association of methylmagnesium bromide in ether), the Grignard reagent is largely associated over much of this concentration range. In the preSent study, the maximum concen- -2 tration of methylmagnesiuM bromide, was 1.52 x 10 M. Further examination of Figure 3 shows that the Grignard reagent is predominantly monomeric

265

3.5

3.0

2.5

2.0

1.5

0 0.5 1. 0 1.5 2.0 2.5

Methylmagnesium Bromide, m

Figure 3. The Effect of the Concentration of Methylmagnesium Bromide on the Degree of Association in Diethyl Ether (Courtesy of E. C. Ashby and D. White).

266

at and below this concentration. Therefore, even as Smith and Su suggest, the declining values of the second-order rate constants observed in their studies may be due to the association phenomena, the associated species being either less reactive or perhaps totally unreactive toward the ketone.

In an effort to find a unifying mechanistic concept which would correlate the kinetic data at both high and low Grignard-to-ketone ratios, the various mechanisms that have been proposed for the Grignard reaction were examined. For each mechanism, the appropriate rate expressions were derived and the rate constants calculated from our data. The results of these studies are discussed in the following paragraphs.

The Grignard Mechanism

Grignard originally proposed that the addition of the organomagnesium reagent to a ketone was a simple second-order reaction.

k 1 G K P

Mathematical treatment of Grignard's mechanism at low Grignard-to-ketone ratios yields equation (11). Our experimental results have shown that the second-order rate constants calculated from equation (11) decrease with decreasing Grignard-to-ketone ratios. This trend seems to rule out the proposed mechanism. The results obtained at high Grignard-to-ketone ratios may also be examined for agreement with Grignard's mechanism. The observed pseudo-first-order rate constant under these conditions is given by equation (13). 267

2.303 Al _ log - k (13) k obs = t Ac 1G o

As equation (13) demands, a plot of k obs versus Go should give a straight line of slope k l passing through the origin. As shown in Figure 4, the appropriate plot does indeed yield a reasonably straight line but the best locus of points does not pass through the origin. A least-squares treatment of the data gives a slope and an intercept of 0.173 and -0.135, respectively. As a result of the inconsistencies at both the high and low

Grignard-to-ketone ratios, Grignard's mechanism seems unlikely.

The Meisenheimer Mechanism

Meisenheimer proposed that the addition of Grignard reagents to ketones proceeds by the following mechanism.

Ki G + K C

C k2 P

Smith and Su (24-26) interpreted their results from a study of the reaction of 2,4-dimethy1-4 - -thiomethyl benzophenone with methylmagnesium bromide in terms of this mechanism. Mathematical treatment of Meisenheimer's mechanism at low Grignard-to-ketone ratios applying the steady-state assumption to the concentration of the complex, C, leads to the same rate expression that Grignard's mechanism does, i.e., equation (11). As previously discussed, the rate constants calculated from equation (11) are dependent upon the Grignard-to-ketone ratio.

Mathematical treatment of Meisenheimer's mechanism at high 268

2.5

2.0

1.5

102k obs slope, 0.173 1. 0

0.5

0 intercept, -0.135

5 10 15 20 3 10 G o

Figure 4. Graphical test of the Grignard Mechanism at High Grignard-to-Ketone Ratios by Equation (13). Daa from Table 2. 269

Grignard-to-ketone ratios leads to equation (14) which is similar to equation (13) except for the definition of the observed rate constant.

2.303 A k = — = K k G (14) obs t A l 2 o c

As equation (14) demands, a plot of k versus G should give a straight obs o line of slope Klk2 passing through the origin. Graphical representation of the data according to equation (14) would be identical to that shown in Figure 4. As previously discussed, the data do yield a reasonably straight line but the best locus of points do not pass through the origin.

If the complex is present ill larger than steady-state concentrations, then the observed rate constant is given by equation (15).

K G 2.303 Al _ 1 k = —-- log A (15) obs t c 11(1G:

Rearranging equation (15) to the slope-intercept form of a straight line gives equation (16).

k2 Go1 G = o k K (16) obs 1

As demanded by equation (16), a plot of G o versus Go /kobs should give a straight line of slope k 2 and an intercept of -1/K 1. A plot of Go versus

G /k is shown in Figure 5. o obs As the results show, the best line through the points has a negative slope and a positive intercept corresponding to negative values for the rate and equilibrium constants which have no physical meaning. These 270

3 10 G o

5 6 7 8 9 10

G /k o obs

Figure 5. Graphical Test of the Meisenheimer Mechanism at High Grignard-to-Ketone Ratios by Equation (16). Data from Table 2. 271

inconsistencies suggest that Meisenheimer's mechanism is also incorrect.

The Dimer Mechanism

In recent years, there has been an overwhelming tendency to interpret the reaction of a ketone or a nitrile with a Grignard reagent as proceeding through the dimeric Grignard reagent.

K 1 2G p

K 2 D+ K -4--- C

k C ---o-3 p+ G

This model has been supported by Mosher (39, 40), Hamelin (41), Becker

(29, 30), and Dessy (12-15, 31) and their respective co-workers. Probably the strongest single piece of evidence against this mechanism (for which there is no comparably strong supporting evidence) is the recent kinetic study of Smith and Su (24-26). As previously mentioned, the second-order rate constants calculated from the pseudo-first-order rate constants reported by these workers, decline as the concentration of methylmagnesium bromide is increased. Since the fraction of the Grignard reagent in the form of the dimer increases with increasing concentration of the organomagnesium compound, the second-order rate constants should increase correspondingly. Since the observed trend in the rate constants is just the reverse of that expected from the association studies, it seems unlikely that the reaction is proceeding through the Grignard dimer. In addition, it should be emphasized that there has never been any strong kinetic evidence to support this mechanism. The second-order rate constants 272

obtained by Bikales and Becker (29) employing this concept correlated only the first 15 to 25 percent of their data for any specific run; the remaining 75 percent was left unexplained.

Still another observation which makes this reaction mechanism implausible hinges on the more recent association data of Ashby and

Walker (16). These workers have shown that association in the Grignard reagent proceeds well beyond the dimer stage and suggests that perhaps trimers are involved at high Grignard concentrations. There is, therefore, no conclusive evidence that a unique dimeric Grignard species exists and, as previously mentioned, even weaker evidence that the addition of the

Grignard reagent to a ketone or a nitrile proceeds through such an intermediate.

Nevertheless, in order to carefully examine each of the previously proposed concepts regarding the mechanism of the Grignard reaction, the necessary kinetic equations were derived for the reaction of benzophenone with methylmagnesium bromide assuming that the reaction did indeed proceed through the dimer. Under conditions where the Grignard reagent is in large excess, the resulting rate expression is given by equation (17).

A k K ED] 2.303 l 3 2 k - log — = obs (17) t Ac 1T72 77

If it is assumed that the concentration of the dimer is low with respect to the monomer under conditions where associations studies show that the Grignard reagent is largely monomeric, i.e., under the conditions of our experiments, then the steady-state approximation can be employed and equation (17) simplified to equation (18).

273

2 k K K G 3 2 1 o k - (18) obs 2 1+K K G 2 1 o

Rearranging equation (18) to the slope-intercept form of a straight line gives equation (19):

2 k G 2 3 o 1 G = o k (19) obs K3

2 where K is As demanded by equation (19), a plot of G versus 3 K1K2. o 2 G /k should give a straight line of slope of k and an intercept of o obs 3 2 2 -1/K3 . A plot of G versus Go is shown in Figure 6. o /kobs As the results show, a plot of the data in accord with equation

(19) does not suggest a straight line. The logical conclusion, therefore, from both a qualitative and quantitative viewpoint is that the addition of Grignard reagents to ketones does not proceed through the dimer mechanism.

The Swain Mechanism

From a study of the effects of added magnesium bromide on the product distribution (i.e., addition, reductioh, and enolization) in the reaction of n-propylmagnesium bromide with diisopropyl ketone, Swain and

Boyles (32) suggested that the addition of Grignard reagents to ketones proceeded by the following mechanism:

K 1 G + K C

k 2 C+ G —4 P 274

250

200

150

N oo

100

0 20 40 60 80 100 4 G /k x 10 o obs

Figure 6. Graphical Test of the Dimer Mechanism at High Grignard-to-Ketone Ratios by Equation (19). Data from Table 2. 275

Although this mechanism was largely discounted when the early association

studies reported (5, 6) that the dimeric Grignard reagent was the principal organometallic species in solution, it nevertheless offered qualitative

explanations for many apparent anomalies observed in Grignard chemistry.

In light of the more recent association studies, which show that associa-

tion is a function of the concentration of the organomagnesium reagent, mechanisms proceeding through a monomeric Grignard species must now be

considered.

In support of the Swain mechanism, Anteunis (33, 34) found that

the reaction of pinacolone and benzophenone with methylmagnesium bromide was first-order in ketone and second-order in the Grignard reagent, third- order over-all. Anteunis applied the steady-state treatment to the con-

centration of the complex in the derivation of his rate equation and in accordance with Swain, assumed that active Grignard reagent was regenerated in the rate-controlling step of the reaction. As previously mentioned, Anteunis' results were internally consistent, i.e., good precision was obtained and there was no apparent drift in the rate

constants for any individual run. The rate constants, however, showed a dependence upon the initial concentrations of methylmagnesium bromide and ketone, the rate constants decreasing with increasing initial concen-

tration of the two reactants. These trends led Becker (29) to dispute

Anteunis' conclusions as to the order of the reaction with respect to the Grignard reagent. Becker maintained that the reaction was first-order

in Grignard dimer and first-order in ketone, second-order over-all.

Unfortunately, Becker's evidence ffor the dimer mechanism was no better than Anteunis' for the Swain mechanism and no clear distinction could be 276

made between the two paths.

Our results at high Grignard-to-ketone ratios were examined for correspondence with the Swain mechanism If the steady-state approximation is applied to the concentration of the complex, C, then the resulting rate expression is given by equation (20).

2.303 to _ 2 k = Al obs A— - k G (20) c 2 o

As previously discussed (Tables 1 and 2), the studies under pseudo-first- order conditions appeared to be first-order in the Grignard reagent not second-order as demanded by equation (20). If the Swain mechanism is correct, the complex must be present in larger than steady-state concentrations.

Assuming that the Grignard-ketone complex is present in larger than steady-state concentrations, the appropriate rate expression is given by equation (21).

2 A k K G 2.303 , l 2 1 o k = obs log -:— = (21) Ac 1 o

Rearranging equation (21) into the slope-intercept form of a straight line gives equation (22).

2 k G 2 o G = o k (22) obs

should give As demanded by equation (22), a plot of Go versus Go 2/kobs a straight line of slope of k 2 and an intercept of -1/K1 . A plot of the 277

data in accordance with equation (22) is shown in Figure 7. The best straight line through the points has a positive slope and a negative intercept as required for a meaningful interpretation. Least-squares -I -1 treatment of the data gives values for k and K of 1.68 2.,mole sec, 2 1 and 1007, respectively. It is evident, therefore, that our data are entirely consistent with the Swain mechanism.

The Swain mechanism was examined further by determining the rate constants at low Grignard-to-ketone ratios using the equilibrium constant ascertained from the pseudo-first-order studies. Under these conditions, the concentration of the Grignard reagent cannot be treated as a constant.

In addition, the steady-state approximation cannot be applied to the concentration of the complex due to the magnitude of the equilibrium constant. The stoichiometry of the reaction with respect to the Grignard reagent, which has been obscured at the high Grignard-to-ketone ratios, must also be considered. The latter has never been defined experimentally.

The stoichiometry problem revolves around the question of whether a mole of Grignard reagent is regenerated in the rate-determining step of the reaction as proposed by Swain or whether the carbinolate complexes with an additional mole of Grignard reagent from which the latter is slowly regenerated.

X X N \ z R, Mg C Mg 11 11 1 R-C=0...Mg -- R-C---0-Mg-X -4- R-COMgX+R 1 MgX rl 278

(.9 0

25 50 75 100

G /k 4 o obs x 10

Figure 7. Graphical Test of the Swain Mechanism at High Grignard-to-Ketone Ratios by Equation (22). Data from Table 2.

279

Alternatively, but necessitating the same stoichiometry, the reaction

may not yield an alkoxymagnesium halide directly but l'iather an alkoxy-

magnesium alkyl which then undergoes rearrangement to the alkoxymagnesium

halide.

R X 1\ A R 4 ' Mg AX R MgX R N\ // , Ili 2 ÷ il C=0...Mg + R-C-0-MgR ÷ R-C-0MgX + R MgX 1 I 1 R R R R/ 1

The additional mole of Grignard reagent (which may or may not be complexed with the carbinolate) is regenerated, therefore, in a secondary and

presumably slower process. Ashby and Arnott (43) have recently shown

that the equilibrium between alkoxymagnesium alkyls and magnesium halides

does indeed lie largely to the right; however, the rate at which this

transformation occurs, which is of prime importance if the reaction

proceeds by the modified Swain mechanism, has not been established.

R--0,MgR + MgX 2 R-C-OMgX + RMgX R

Regardless of how complicated the actual transformations are, they

could be treated in the general mechanism (23):

K 1 G +K-7—+ C (23)

k 2 C+ G —0- P 1

K 3 P P + G 1 280

where P is the reaction product with which a mole of Grignard reagent is 1 complexed and P is expected carbinolate. Unfortunately, mathematical

treatment of this mechanism (when K is of the order of 1000) is quite 1 complex. Furthermore, since there are two unknowns, k 2 and K 3 (assuming

K to be determined by the pseudo--first-order studies), it is evident 1 that neither could be calculated directly. After due consideration, it

was decided that the precision of our data and the uncertainties in the

value of K %-did not warrant the detailed study necessary to treat the

data in this manner. As a result, only the two limiting cases were con-

sidered: 1) the original proposal of Swain where the Grignard reagent

is regenerated in the rate-determining step of the reaction and 2) a

modified Swain mechanism where the second mole of Grignard reagent remains

complexed with the carbinolate. The mathematical derivation of the appropriate rate expressions for both cases are shown in Derivations 2 and 3,

Appendix; however, to illustrate the approach employed, derivation of

the equations for the modified Swain mechanism will be outlined here.

The differential rate expression and the solution to the quadratic

equation for calculating the concentration of the complex are given by

equations (24) and (25), respectively.

d[P]/dt = k EC [G] 2 j (24)

1+K (G +K -3[P])1 274K (G K +2K )[P]+2[P] (25) 1 o o 1 o +K o 1 o oo o [C] = 2K 1

Substitution of equation (25) into equation (24) does not give an

expression which is readily integrated. It was evident, therefore, that 281 equation (25) would have to be simplified before it could be integrated.

Simplification was effected by observing that C could be expressed as a linear function of the product, P. It will be noted that the concentration of C can be calculated from equation (25) at the various concentrations of P by employing the value of K 1, e.g., 1007, as determined from the study under pseudo-first-order conditions. A plot of C versus P can then be constructed (Figure 8, Appendix) and C expressed as:

[ C] = C, - m[P] (26) where -m is the slope determined from the aforementioned plot and C o in the intercept at zero time. Substitution of equation (26) into the differential rate expression (24) gives (27).

d[P]/dt = k (C - m[P])[G] (2 7) 2 o

The concentration of the Grignard reagent, G, can be expressed by equation (28).

[G] = Go - [C] - 2[P] (28)

Substituting equation (26) into equation (28) and collecting terms, equation (29) is obtained.

[G] = Go - Co + (m-2)[P] (29)

Letting Go-00 equal Ao and substituting equation (29) into equation (27), gives equation (30).

d[P]/dt = k (C - m[P])(A + (m-2)[P]) (30) 2 o o 282

Integrating equation (30) gives equation (31).

C (A +(m-2)5N 2.303 o o k log 2 t(C(m-2) + mA) (31) o o Ao (C o -

Analyses of the data (Tables 24 through 28, Appendix) obtained at low

Grignard-to-ketone ratios in terms of both the Swain and modified Swain mechanisms are shown in Table 3.

As the results show, the rate constants calculated from the equations derived from the Swain and the modified Swain reaction mechanisms give average values of 1.35 and 1.69 k. moles -1 sec. -1 , respectively.

The average value of the rate constants calculated from the pseudo-first- -1 -1 order studies was 1.68 L. mole sec. It is evident, therefore, that the termolecular reaction mechanism originally proposed by Swain, with the added stipulations that the equilibrium constant for complex formation is large and that perhaps a modification in the stoichiometry of the reaction is necessary, best describes the transformation. The data presented in this study represent the widest range of Grignard-to-ketone ratios that have been reported. The fact that the rate constants are indeed constant over, the entire range of Grignard-to-ketone ratios strongly supports the proposed mechanism. Furthermore, the rate equations derived from the proposed mechanism correlate the rate data over the entire range wherein meaningful kinetic data may be obtained, namely, between 15 and 80 percent reaction. Some of the other reaction mechanisms that have been proposed have failed at this point.

The agreement between the average, value of the rate constants obtained at high Grignard-to-ketone ratios and that obtained at low 283

Table 3. Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at Low Grignard-to- Ketone Ratios at 25°

Methylmagnesium a b c Bromide k k k 1 -1 -1 -1 -1 Benzophenone . Z. mole -1 sec. t. mole sec. Z. mole sec.

6.86 0.58 ± 0.07 1.54 ± 0.09 1.66 ± 0.17 6.86 0.41 ± 0.03 1.10 ± 0.08 1.17 ± 0.07 2.73 0.35 ± 0.03 1.26 ± 0.10 1.56 ± 0.14 2.59 0.72 ± 0.08 1.68 ± 0.24 1.88 ± 0.22 1.37 0.36 ± 0.18 1.10 ± 0.24 2.17 ± 0.56

Avg. 1.35 Avg. 1.69 a Second-order rate constants, calculated from equation (11). b Rate constants calculated from rate equations derived from the Swain mechanism, i.e., equation (7) Derivation 2, Appendix. cRate constants calculated from rate equations derived from the modified Swain mechanism, i.e., equation (7), Derivation 3, Appendix. 284

Grignard-to-ketone ratios employing the modified Swain reaction mechanism

suggests that the latter best describes the stoichiometry of the reaction; however, due to the rather large average deviation, these data certainly

do not firmly establish this point. Furthermore, examination of the data

(Tables 24 through 28, Appendix) shows that there is a trend toward falling

rate constants when either of the two stoichiometries is employed. The

exact stoichiometry, therefore, may be more complex than our mathematical

treatment affords. The trends, however, are more pronounced when the ,

Swain stoichiometry is employed. The precision in the rate constants is

also better when the rate equation derived from the modified Swain mechanism is employed. For these reasons, we believe that the latter best describes the stoichiometry of the reaction; however, it is recognized

that more work is necessary to establish this point.

The modified Swain stoichiometry is in qualitative accord with the

observations of Storfer and Becker. (30) regarding the reaction of benzonitrile with ethylmagnesium bromide. These workers reported that the

aromatic nitrile is converted to the corresponding ketimine only when two

alkyl groups are present for each nitrile group. This is not to be

interpreted that a 1/1 ratio of Grignard reagent to nitrile is totally unreactive, it is only that the reaction is so slow in the absence of a

2/1 excess, that quantitative conversions cannot be effected in any reasonable length of time. Storfer and Becker interpreted their findings in terms of the dimer mechanism; however, they can be explained better in

terms of the modified Swain mechanism. The imine formed would be expected to be markedly more nucleophilic (basic) than the corresponding alkoxymagnesium halides formed in the reaction of Grignard reagents with 285

ketones. They are capable, therefore, of forming more stable complexes with the organomagnesium compound. As a result, the 2/1 stoichiometry is accentuated. These observations also lead us to believe that the reaction is best described in terms of the modified Swain stoichiometry.

1 1;

286

CHAPTER IV

CONCLUSIONS

The kinetics of the reaction of methylmagnesium bromide with

benzophenone in diethyl ether have been studied by a new experimental

technique which promises to be applicable to a wide range of organo-

metallic reactions. The reaction was studied at Grignard-to-ketone

ratios ranging from 1.4 to 152/1 which represent the broadest range

reported to date. Analysis of the kinetic data has shown that the

reaction is best described in terms of the Swain mechanism with the

added stipulatipn that complex formation is governed by an equilibrium

constant, K1 , of about 1000.

K 1 K C

k C + G P + G

There is also the distinct possibility that the stoichiometry of the

reaction as originally proposed by Swain may be incorrect, i.e., the

active Grignard reagent may ,not be regenerated in the rate-determining

step of the reaction, but may remain complexed to the carbinolate. In

this form it presumably has a lower reactivity toward the ketone;

however, more work is necessary to verify this point. 287

The kinetic equations derived from the modified Swain mechanism give consistent values for the rate constants which are independent of the Grignard-to-ketone ratios. In addition, the rate equations correlate the data throughout the entire region wherein meaningful kinetic data can be derived, i.e., between 15 and 75 percent reaction. The rate equations derived from other reaction mechanisms that have been proposed fail at one or both of these points. These observations lead us to believe that the mechanism of the Grignard reaction is best described in terms of the modified Swain mechanism. 288

A:PPEN DI X 289

Derivation 1. Correction of the ,Benzophenone,Absorbance for the Absorbance of Diphenyl Methyl Carbinol.

Definitions

A = measured absorbance at time, t.

A = initial absorbance of benzophenone. o A = absorbance of reaction mixture due to benzophenone. c K = benzophenone.

K = initial benzophenone concentration. o D = diphenyl methyl carbinol. 4 eK = extinction coefficient for benzophenone (1.84 x 10 ).

C D = extinction coefficient for diphenyl methyl carbinol (342)

Assumptions

The derivation is based on stoichiometry of the reaction mixture after hydrolysis.

K0 = EDI + [K]

Note:: The cell-path length, which is constant, is not carried along in the derivation.

Derivation

A = eK[K] + e D [D] (1)

A = e [K] + - [K]) (2) K eD(Ko

= [K](e K - E D ) + Ko e D (3)

Since A =K e o o K E D A = [K](e — A ( 4 ) K - e D ) + e o K 290

Solving equation (4) for [K] gives equation (5).

A (c/e ) A [K] - K o (5) el< e D

A Since [K] = E K

Ae - eb A [A - (e /e ) A ] K o D K o A - (6) c EK - E D T1-1 e ie K

. . Substituting the values of e lf and e in .equation (6) gives K equation (7).

A - 0.0188 A A - c 0.981 (7) 291

Derivation 2. Derivation of the Rate Equation for the Swain Mechanism at Low Grignard-to-Ketone Ratios

Mechanism

K 1 G + K C

C + G —4 P + G

Conditions

Definitions

G and K equal the initial Grignard and. ketone concentrations, o .o re spectively.

Assumptions

1. K is large and the steady-state assumption cannot be applied to C. 1 2. The rate-controlling step of the reaction is definitely k 2 .

3. The stoichiometry after hydrolysis is defined by:

K = [K] + [P] o

Derivation

d[P]/dt = k 2 [C][G]

„,2 2 „2, „ +K(G +K -2[P])41+K (G +K -4K (G K -[G +K ][P]+LPi ) 2) o o 1 o o 1 o o o o [C] = 2K 1

1. Calculate the concentrations of the complex, C, at the various concentrations of P obtained from equation (2).

2. Plot [C] versus [P] (Figure 8). 292

Derivation 3. Derivation of the Rate Equation for the Modified Swain Mechanism at Low Grignard- to-Ketone Ratios

Mechanism

Same as Derivation 2.

Conditions

Same as Derivation 2.

Definitions

Same as Derivation 2.

Assumptions

Same as Derivation 2.

Derivation

d[P]/dt = k2 ECHG]' (1)

1+K (G +K -3[P])-i[1+K (G +K -3[P])J2-4K 2 (G K -(G +2K )[P]+2[P] 2 ) EC] = 1 o o 1 o o 1 o o o o 2) 2K ( 1

1. Calculate the concentrations of the complex, C, at the various concentrations of P obtained from equation (2).

2. Plot [C] versus [P] (Figure 8).

3. Express the concentration of the complex, C, as C - m[P] where m is the slope obtained from Figure 8 and Co is the intercept at zero time. [C] = co-MIN (3)

293

3. Express the concentration of the complex, C, as C - m[P] where m is the slope obtained from Figure 8 and C is the o intercept at zero time. [C] = C -m[P] o (3)

4. Express the concentration of the Grignard reagent as:

[G] = G-[C]-[P] ( 4 ) 0

5. Substituting equation (3) in equation (4) gives equation (5).

[G] = Go-00+(m-1)[P] (5)

6. Letting G -00 equal Ao and substituting equations(3) and (5) into equat ion (1) gives equation (6). d[P]/dt = k (C 2 o-m[P])(Ae+(m-1)[1:]) ( 6 )

7. Integrating equation (6) gives equation (7).

C (A 2.303 +(m-1)[P]) o o k22 t[Co(m-i)i-mAo] log A (C -m[P]) ( 7 ) o o 294

4. Express the concentration of the Grignard reagent as: [G] = Go -[C]-2[P] (4)

5. Substituting equation (3) in equation (4) gives equation (5). [C] = G•Co+(m-2)[P] (5)

6. Letting G -Co equal Ao and substituting equations(3) and (5) into equaion (1) gives equation (6).

d[P]/dt = k2 (C0 -m[P])(A0+(m-2)[P]) (6)

7. Integrating equation (6) gives equation (7).

C (A +(m-2)[P]) 2.303 o o log ) k2 =t[C(m-2)÷mA] A (C (7 0 0 o 0 295

Table 4

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

10 2k c a Time k -1 -1 Seconds Absorbance sec. Z. mole sec.

0 0.964 0 0.953 51.3 0.827 0.284 1.23 102.0 0.756 0.231 1.00 151.6 0.655 0.255 1.11 201.1 0.573 0.261 1.13 251.6 0.511 0.255 1.11 303.3 0.477 0.236 1.02 407.0 0.375 0.237 1.03

Avg. 0.251±0.014 1.09 ±0.06

Initial Concentrations: -4 Benzophenone 1.009 x 10 M b Methylmagnesium Bromide 2.31 x 10 3 m

Initial Absorbance (extrapolated) ...... 0.955 aCalculated from equation (10), where n=1. bBy dilution. cCalculated from equation (8).

1 296

Table 5

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k ohs a Time k -1 -1 Seconds Absorbance sec. I Q. mole sec.

0 0.902 52.0 0.733 0.333 1.44 91.4 0.658 0.312 1.35 114.4 0.632 0.285 1.23 150.6 0.535 0.331 1.43 250.0 0.378 0.343 1.49 302.2 0.364 0.298 1.30 360.1 0,303 0.304 1.32

Avg. 0.315 ±0.018 1.37 ±0.08

Initial Concentration: -4 Benzophenone ..... 1.009 x 10 M b -3 Methylmagnesium Bromide ...... 2.31 x 10 -3 M

Initial Absorbance (extrapolated) ..... 0.866 a Calculated from equation (10), where n=1. b By dilution. c Calculated from equation (8). 297

Table 6

Reaction of Methyimagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k a Time k 1 -1 Seconds Absorbance sec. Q. mole sec.

0 0.936 61.6 0.793 0.260 1.13 101.2 0.703 0.281 1.22 126.0 0.625 0.323 1.40 161.1 0.551 0.332 1.44 201.4 0.493 0.324 1.40 252.1 0.430 0.316 1.37 301.7 0.387 0.299 1.29 373.6 0.337 0.280 1.21

Avg. 0,302 1'0.022 1.31 ±0.09

Initial Concentrations: Benzophenone 1.009 x 10 -4 M

Methylmagnesium Bromide 2.31 x 10 -3M

Initial Absorbance (extrapolated) 0.927 aCalculated from equation (10), where n=1. bBy dilution. cCalculated from equation (8). 298

Table 7

Reaction of Methylmagnesium Bromide with genzophenone in Ether at 25°

2 c 10 kobs a Time k -1 Seconds Absorbance sec. -1 Z. mole sec.

0 0.870 51 0.642 0.203 0.88 100 0.567 0.233 1.01 150 0.496 0.248 1.07 200 0.452 0.233 1.01 300 0.330 0.264 1.15 400 0.221 0.305 1.32 500 0.230 0.235 1.02 500 0.233 0.234 1.01

Avg. 0.244 -1 0.021 1.06 ±0.09

Initial Concentrations: Benzophenone 9.90 x 10 _ 3M Methylmagnesium Bromide 2.31 x 10 M

Initial Absorbance (extrapolated) 0.708 aCalculated from equation (10), where n=1. bBy dilution. c Calculated from equation (8). 299

Table 8

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 d 10 k Time obs a -k1 -1 Seconds Absorbance sec. -1 Z. mole sec.

0 0.935 c 51.0 0.624 0.237 c 0.52 76.8 0.454 0.590 1.28 90.1 0.435 0.551 1.19 105.3 0.405 0.541 1.17 126.5 0.354 0.523 1.22 152.7 0.310 0.558 1.21 202.2 0.237 0.563 1.22

Avg. 0.554 1 0.016 1.22 1 0.03

Initial Concentrations: 4 Benzophenone 1.009 x 10_3M b Methylmagnesium Bromide 4.62 x 10 -- M

Initial Absorbance (extrapolated) 0.697 aCalculated from equation (10), where n=1. bBy dilution. c Omitted from average. dCalculated from equation (8) 300

Table 9

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k a Time obs k 1 -1 Seconds Absorbance sec. -1 2. mole sec.

0 0.922 32.2 0.725 0.589 1.28 51.8 0.652 0.617 1.34 73.8 0.600 0.551 1.19 100.8 0.493 0.607 1.31 125.6 0.450 0.562 1.22 149.3 0.374 0.603 1.31 199.7 0.299 0.569 1.23 251.7 0.212 0.599 1.30

Avg. 0.587 1 0.020 1.27 1 0.05

Initial Concentrations: Benzophenone 1.009 x 10 _ 3M b Methylmagnesium Bromide 4.62 x 10 M

Initial Absorbance (extrapolated) 0.891 aCalculated from equation (10), where n=1. bBy Dilution. Calculated from equation (8). 301

Table 10

Reaction of Methylmagnesium Bromide with Benzophenone in : Ether at 25°

2c 10 k Time ohs k 1a - Seconds Absorbance sec. -1 Q. mole sec.

0 0.953 12.9 0.661 1.57 1.70 16.9 0.602 1.78 1.92 25.5 0.533 1.67 1.81 32.1 0.480 1.66 1.80 40.3 0.436 1.56 1.69 50.5 0.362 1.64 1.78 60.5 0.309 1.65 1.78

Avg. 1.65 ±0.05 1.77 ±0.06

Initial Concentrationt: - Benzophenone 1.009 x 10 ^ 3M b Methylmagnesium Bromide 9.24 - x 10 M

Initial Absorbance (extrapolated) 0.802

a Calculated from equation (10), where n=1. bBy dilution. c Calculated from equation (8). 302

Table 11

Reaction ofMethylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k Time obs ka -1 -1 Seconds Absorbance sec.-1 Z. mole sec.

0 0.964 16.6 0.659 1.43 1.55 22.8 0.590 1.54 1.66 25.5 0.580 1.45 1.57 36.1 0.487 1.52 1.65 41.6 0.451 1.51 1.63 50.6 0.425 1.37 1.48 61.9 0.346 1.47 1.59

Avg. 1:47 *0.05 1.59 ±0.05

Initial Concentrations: -4 Benzophenone 1.009 x 10 M b 3 Methylmagnesium Bromide 9.24 x 10 7' M

Initial Absorbance (extrapolated) 0.828 aCalculated from equation (10), where n=1. bBy dilution. cCalculated from equation (8). 303

Table 12

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

10kc2 a Time obs k -1 -1 Seconds Absorbance sec. Z. mole sec.

0 0.968 16.7 0.701 1.14 1.23 25.6 0.607 1.32 1.43 33.8 0.556 1.27 1.37 42.2 0.504 1.26 1.37 52.1 0.464 1.18 1.28 62.5 0.379 1.32 1.43 72.4 0.367 1.19 1.29 89.3 0.297 1.22 1.32

Avg. 1.24 1 0.05 1.34 1 0.06

Initial Concentrations: 4 Benzophenone 1.009 x 10 M b * Methylmagnesium Bromide 9.24 x 10 M

Initial Absorbance (extrapolated) 0.842 aCalculated from equation (10), where n=1. bBy dilution. cCalculated from equation (8). 304

Table 13

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2h 10 k a Time obs k 1 -1 Seconds Absorbance sec. -1 L. mole sec.

0 0.854 23.7 0.487 1.24 1.34 38.0 0.410 1.24 1.34 44.9 0.395 1.11+ 1.23 52.3 0.344 1.25 1.36 59.8 0.318 1.4 1.34 70.1 0.267 1.31 1.42 90.5 0.216 1.28 1.39 110.3 0.175 1.25 1.35

Avg. 1.24 ±0.03 1.35 ±0.04

Initial Concentrations; -5 Benzophenone 9.90 x 10 M c -3 Methylmagnesium Bromide 9.24 x 10 -3 M

Initial Absorbance (extrapolated) 0.643

..aCalculated from equation (10), where n=1. b Calculated from equation (8). cBy dilution 305

Table 14

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 b 10 k a Time obs k -1 -1 Seconds Absorbance sec. -1 Z. mole sec.

0 0.940 13.4 0.635 1.23 1.33 18.2 0.581 1.41 1.53 22.1 0.537 1.53 1.66 27.7 0.513 1.39 1.51 33.4 0.474 1.40 1.52 37.3 0.447 1.41 1.53 48.4 0.390 1.39 1.51 60.5 0.331 1.40 1.51

Avg. 1.40 1 0.04 1.51 1 0.05

Initial Concentrations: -5 Benzophenone 9.90 x 10 M c -3 Methylmagnesium Bromide 9.24 x 10 M

Initial Absorbance (extrapolated) 0.747

a Calculated from equation (10), where n=1. b Calculated from equation (8). cBy dilution. 306

Table 15

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 b 10 kobs Time ka -1 -1 -1 Seconds Absorbance sec. k. mole sec.

0 0.882 41.2 0.741 0.238 0.97 d 61.9 0.733 0.177 d 0.72 92.2 0.642 0.264 1.08 128.1 0.571 0.285 1.16 163.1 0.527 0.275 1.12 203.8 0.482 0.265 1.08 253.5 0.425 0.263 1.07 302.4 0.384 0.257 1.05 403.7 0.300 0.254 1.04

Avg. 0.263 -1 0.010 1.07 10.04

Initial Concentrations: -4 Benzophenone 1.027 x 10 M c 3 Methylmagnesium Bromide 2.45 x 10 M

Initial Absorbance (extrapolated) 0.813 aCalculated from equation (10), where n=1. b Calculated from equation (8). c . By analysis. d Omitted from average. 307

Table 16

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 b 10 }co bs a Time k -1 Seconds Absorbance sec. Q. mole see.sec.

0 0.888 61.9 0.677 0.251 1.05 94.4 0.606 0.283 1.19 126.5 0.558 0.278 1.17 160.3 0.527 0.257 1.08 198.9 0.475 0.260 1.09 247.7 0.419 0.263 1.10 301.5 0.409 0.224 0.94 402.3 0.319 0.232 0.98

Avg. 0.256 ±0.015 1.08 ±0.06

Initial Concentrations: -4 Benzophenone 1.027 x 10 M -3 Methylmagnesium Bromide c 2.380 x 10 M

Initial Absorbance (extrapolated) 0.784 a Calculated from equation (10), Where n=1. b Calculated from equation (8). c By analysis.

308

Table 17

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k Time obs ka -1 -1 Seconds Absorbance sec. -1 mole Z. sec.

0 0.870 41.1 0.750 0.359 1.45 61.2 0.732 0.283 1.15 91.1 0.645 0.332 1.34 122.6 0.596 0.312 1.27 161.9 0.546 0.291 1.18 302.5 0.353 0.306 1.24 403.2 0.300 0.272 1.10

Avg. 0.308 ±0.024 1.25 ±0.09

Initial Concentrations: -4 Benzophenone 1.010 x 10 M b -3 Methylmagnesium Bromide 2.47 x 10 M

Initial Absorbance (extrapolated) ..... 0.866 a Calculated from equation (10), where n=1. b By analysis. c Calculated from equation (8) 309

Table 18

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k a Time obs k 1 -1 Seconds Absorbance sec. -1 9— mole sec.

0 0.896 30.5 0.650 0.684 1.51 51.4 0.572 0.664 1.47 76.0 0.458 0.754 1.67 88.2 0.428 0.728 1.61 101.5 0.384 0.745 1.65 125.9 0.350 0.677 1.50 161.9 0.248 0.753 1.67 201.9 0.219 0.671 1.49

Avg. 0.709 ±0.035 1.57 ±0.08

Initial Concentrations: -4 Bentophenone 1.010 x 10 M b -3 Methyltnagnesium Bromide 4.52 x 10 M

Initial Absorbance (extrapolated) 0.795 aCalculated from equation (10), where n=1. b By analysis. c Calculated from equation (8). 310

Table 19

Readtion of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 c 10 k a Time obs k 1 -1 Seconds Absorbance sec. -1 Q. mole sec.

0 0.866 0 0.850 36.1 0.588 0.804 1.67 50.2 0.541 0.746 1.54 75.9 0.413 0.859 1.78 101.6 0.362 0.777 1.61 130.6 0.301 0.755 1.56 151.7 0.248 0.783 1.62 243.0 0.141 0.746 1.54

Avg. 0.781 ±0.029 1.62 ±0.07

Initial Concentrations: -4 Benzophenone 1.010 x 10 M b -3 Methylmagnesium Bromide 4.83 x 10 -3 M

Ihitial Absorbance (extrapolated) 0.778

a Calculated from equation (10), where n=1. b By analysis. c Calculated from equation (8). 311

Table 20

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

102 k Time obs ka 1 -1 Seconds Absorbance sec. -1 R. mole sec.

0 0.946 12.2 0.679 1.49 1.57 20.2 0.630 1.28 1.34 30.7 0.540 1.36 1.43 35.7 0.494 1.43 1.50 41.4 0.470 1.36 1.42 51.1 0.411 1.37 1.44 60.5 0.375 1.32 1.38 75.9 0.235 1.72 1.80

Avg. 1.42 ±0.10 1.49 ±0.11

Initial Concentrations: Benzophenone 1.027 x 10 _ 3M b Methylmagnesium Bromide 9.54 x 10 M

Initial Absorbance (extrapolated) ..... 0.807 aCalculated from equation (10), where n=1. b By analysis. cCalculated from equation (8). 312

Table 21

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 a b 10 k Time obs -1 -1 Seconds Absorbance sec. -1 Z. mole sec.

0 0.870 17.6 0.598 1.47 1.52 26.1 0.539 1.41 1.46 31.4 0.497 1.43 1.48 36.3 0.456 1.48 1.54 41.9 0.439 1.38 1.43 50.7 0.368 1.51 1.56 61.6 0.327 1.44 1.50

Avg. 1.45 ±0.04 1.50 ±0.04

Initial Concentrations: -4 m Benzophenone 1.010 x in n Methylmagnesium Bromide c 9.64 x 10-3 m

Initial Absorbance (extrapolated) 0.768 a Calculated from equation (8). b Calculated from equation (10), where n=1. °By analysis. 313

Table 22

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2 a b 10 k k Time obs -1 -1 Seconds Absorbance sec. -1 k. mole sec.

0 13.1 0.539 2.60 1.71 20.4 0.438 2.72 1.79 26.4 0.400 2.46 1.62 30.4 0.300 3.14 2.06 35.9 0.303 2.63 1.73 40.8 0.252 2.80 1.84 45.3 0.951 2.53 1.66

Avg. 2.62 10..09 1.73 1 0.06

Initial Concentrations: 4 Benzophenone c 1.065 x 10 M -2 Methylmagnesium Bromide- 1.52 x 10 M

Initial Absorbance (extrapolated) 0,745 a Calculated from equation (8). b Calculated from equation (10), where n=1. c By analysis. 314

Table 23

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

2a b 10 k k Time obs -1 -1 Seconds Absorbance sec. -1 Z. mole sec.

0 0.992 10.3 0.620 2.a2 1.60 14.5 0.573 2.20 1.51 20.4 0.487 2.38 1.64 25.3 0.423 2.50 1.73 30.1 0.404 2.27 1.56 35.8 0.349 2.34 1.61 40.3 0.320 2.31 1.59 45.1 0.290 2.29 1.58

Avg. 2.33 1 0.03 1.60 1 0.04

Initial Concentrations: -4 Benzophenone ...... 1.065 x 10 M Methylmagnesium Bromide c 1.45 x.10 M

Initial Absorbance (extrapolated) 0.776 a Calculated from equation (8). b Calculated from equation (10), where n=1. cBy analysis. Table 24

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

a b c Time k k k -1 -1 -1 -1 -1 -1 Seconds Absorbance Z. mole sec. Z. mole sec. mole sec.

0 0.928 0 0.926 600 0.752 0.72 1.93 1.99 900 0.694 0.61 1.65 1.71 1200 0.609 0.63 1.69 1.78 1860 0.483 0.61 1.63 1.75 2400 0.413 0.58 1.55 1.69 3120 0.360 0.52 1.40 1.54 5220 0.275 0.40 1.06 1.20

0.58 ± 0.07 1.56 ± 0.19 1.66 ± 0.17

Initial Concentrations Benzophenone 1.009 x 10-4 M -4 Methylmagnesium Bromide 6.924 x 10 M a Calculated from equation (11). ?Calculated flrom equation (7), Derivation 2, Appendix, where: a=0.6529, b=0.03950, c=0.3975, d=0.6025. cCalculated from equation (7), Derivation 3, Appendix, where: a=0.6531, 1)=0.03930, c=0.4085, d=1.5915.

Cr.) cn Table 25

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

a b c k k k Time -1 -1 -1 -1 -1 -1 Seconds Absorbance i. mole sec. i. mole sec. i. mole sec.

0 0.944 660 0.774 0.46 1.22 1.29 960 0.703 0.46 1.25 1.33 1260 0.677 0.40 1.08 1.15

1800 0.621 . 0.36 0.96 1.02 2460 0.497 0.40 1.09 1.18 3000 0.437 0.40 1.08 . 1.18 3720 0.500 0.27d 0.71c 0.78d 4500 0.329 0.37 1.01 1.12

Avg. 0.41 ± 0.03 1.10 ± 0.08 1.18 ± 0.07

Initial Concentrations -4 Benzophenone 1.009 x 10 M -4 Methylmagnesium Bromide 6.924 x 10 M aCalculate0 from ewation (1.1). b Calculated from equation (7), Derivation 2, Appendix, where: a=0.6526, b=0.03980, c=0.4114, d=0.5886. c Calculated from equation (7), Derivation 3, Appendix, where: a=0.6543, b=0.0381, c=0.3900, d=1.610. dOmitted from average, > three times the average deviation. Table 26

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

a b k k kc Time -1 -1 -1 -1 -1 -1 Seconds Absorbance R. mole sec. R. mole sec. R. mole sec.

0 0.946 d 1000 0.818 0.27 0.93 1.00 d 1500 0.729 0.34 1.17 1.30 2000 0.644 0.38 1.32 1.52 2500 0.570 0.41 1.42 1.70 3000 0.593- 0.31 1.09 1.38 3500 0.468 0.39 1.37 1.73 4500 0.443 0.35 1.24 1.62 5500 0.387 0.35 1.23 1.69

Avg. 0.35 ± 0.03 1.26 ± 0.10 1.56 ± 0.14

Initial Concentrations Benzophenone 2.054 x 10-4 -4 M Methylmagnesium Bromide 5.604 x 10 M a Calculated from equation (11). bCalculated from equation (7), Derivation 2, Appendix, where: a=0.4937, b=0.06670, c= 0.3514, d=0.6486. c Calculated from equation (7), Derivation 3, Appendix, where: a=0.5003, b=0.06370, c=0.3692, d=1.6308. d Omitted from average, > three times average deviation. Table -27

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

a b c k k k Time -1 -1 -1 -1 -1 -1 Seconds Absorbance Z. mole sec. Z. mole sec. Z. mole sec.

0 0.986 0 0.979 99.3 0.893 0.77 1.81 1.88 199.5 0.782 0.94 2.20 2.37 400 0.670 0.82 1.91 2.15 600 0.600 0.72 1.67 1.95 800 0.5 9 5 0.71 1.65 2.01 1000 0.510 0.59 1.39 1.71 1550 0.430 0.50 1.18 1.54

Avg. 0.72 ± 0.08 1.68 ± 0.24 1.94 ± 0.20

Initial Concentrations -4 Benzophenone 5.32 x 10 M -4 Methylmagnesium Bromide 13.78 . R 10 M aCalculated from equation (11). b Calculated from equation (7), Derivation 2, Appendix, where: a=1.1115, b=0.2670, c=0.5885, d=0.4115. c Calculated from equation (7), Derivation 3, Appendix, where: a=1.1155, b=0.2630, c=0.6086, d=1.3994. Table 28

Reaction of Methylmagnesium Bromide with Benzophenone in Ether at 25°

a b c Time k k k -1 -1 -1 -1 -1 Seconds Absorbance Z. mole sec. Z. mole sec. Z. mole sec.

0 0.901 720 0.692 0.60 2.37 3.09 1200 0.606 0.57 2.23 3.36 1800 0.618 0.36 1.41 2.08 2700 0.542 0.34 1.35 2.33 3660 0.536 0.26 1.04 1.82 5460 0.517 0.19 0.75 1.38 11820 0.436 0.12 0.50 1.27 20040 0.366 0.093 0.41 2.05

Avg. 0.36.± 0.18 1.25 ± 0.58 2.17 ± 0.56

Initial ,Concentrations Benzophenone 5.050 x 10-4 M -4 Methylmagnesium Bromide 6.924 x 10 M a Calculated from equation (11). b Calculated from equation (7), Derivation 2, Appendix, where: a=0.5149, b=0.1775, c=0.4500, d=0.5500. c Calculated from equation (7), Derivation 3, Appendix, where: a=0.5364, b=0.1560, c=0.4815, d=1.5185.

0.) ED 320

0.5

0.4

0.3

0.2

0.1

0 0.1 0.2 0.3 0.4 0.5 0.6 0.7

Figure P. Calculation of the Rate Constants for the Reaction of Methylmagnesium Bromide with Benzophenone at Low Grignard- to-Ketone Ratios with the Modified Swain Stoichiometry. Determination of the Relationship between C and P, Derivation 3, Step 3. Data from Table 24 , Appendix. 321

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VITA

Roy B. Duke, Jr., was born in Houston, Texas, on September 20,

1932, and attended primary and secondary schools there. His undergraduate work was done at Lamar Institute of Technology, Beaumont, Texas, and the University of Houston, Houston, Texas. In 1956, he received a

Bachelor of Science degree in chemistry from the latter. During most of his undergraduate work, he worked full, time for Texas Electric Steel

Casting Company of Houston as an analytical chemist. After graduation, he joined The American Oil Company in Texas City, Texas, as a chemist in the Research and Development Section. At night he continued his education at the University of Houston, obtaining a Master of Science degree in 1960.

In 1960, h ,joined the Research and Development Department of

Texas Eastman, a Division of Eastman Kodak, in Longview, Texas, where he served in the capacity of research chemist.

In 1962, he returned to graduate school at the Georgia Institute of Technology as an instructor in the Chemistry Department. His graduate research was begun in December, 1963, under the direction of Dr. Jack

Hine. After Dr. Hine's departure from the Georgia Institute of Technology in July, 1965, Roy continued his graduate research under the direction of

Dr. E. C. Ashby.

In June of 1966, he joined the Organic Chemistry Department of the Marathon Oil Company, Littleton, Colorado.

Roy is married to the former Theadus Ann Schrader and is the father of three sons: Larry A., Paul T., and Dwight F. Duke.