MECHANISM OF OXIDATION

OF FLUORAL AND RELATED COMPOUNDS

MICHAEL MIROSLAV MOCEK

B.A., University of British Columbia, 1957

MSc, University of British Columbia, i960

A THESIS SUBMITTED IN PARTIAL FULFILMENT OF

THE REQUIREMENTS FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

in the Department

of

We accept this thesis as conforming to the required standard

THE UNIVERSITY OF BRITISH COLUMBIA

•October, 1962 In presenting this thesis in partial fulfilment of the requirements for an advanced degree at the University of

British Columbia, I agree that the Library shall make it freely available for reference and study. I further agree that permission for extensive copying of this thesis for scholarly purposes may be granted by the Head of my Department or by his representatives.

It is understood that copying or publication of this thesis for financial gain shall not be allowed without my written permission.

Department of

The University of British Columbia, Vancouver 8, Canada.

Date The University of British Columbia

FACULTY OF GRADUATE STUDIES

PROGRAMME OF THE

FINAL ORAL EXAMINATION

FOR THE DEGREE OF

DOCTOR OF PHILOSOPHY

of

MICHAEL M. MOCEK

B.A., University of British Columbia, 1957 M.Sc, University of British Columbia, 1960

THURSDAY, OCTOBER 11, 1962, AT 3:00 P.M. IN ROOM 261, CHEMISTRY BUILDING

COMMITTEE IN CHARGE

Chairman: F.H. SOWARD

PUBLICATION F.W. DALBY D.E. McGREER D.G.L. JAMES C. REID Deuterium Isotope Effects in Organic C.A. MCDOWELL A.I. SCOTT Cations. R. Stewart, A.L. Gatzke, M. R. STEWART Mocek and K. Yates. Chem. and Ind., 331 (1959). External Examiner: W.A. WATERS, F.R.S. Balliol College Oxford, England MECHANISM OF PERMANGANATE OXIDATION OF FLUORAL AND RELATED COMPOUNDS Formic acid has been oxidised by permanganate in the region 26 to 1TL H2SO4 and the reaction in the region ABSTRACT from 25 to about 40% H2SO4 is interpreted as the oxidation of the unionised formic acid by permanganic The mechanism of permanganate oxidation of fluoral acid. An isotope effect of about 7 is observed in the hydrate and its analogue, substituted with deuterium very strongly acidic region (7 27» H2SO4), which de• at position-1, has been extensively investigated creases upon going to 257o H2SO4. throughout the pH region 0.7 to 14 and a successful The structure of fluoral hydrate, synthesised in study was made of the behaviour of permanganate in the this laboratory, was unequivocally established by region from pH 0.7 to 46.37o H2SO4. n.m.r. to be the dimer of fluoral hydrate and spin- Compared with 2,2,2-trifluoroethanol and other spin splitting of the OH proton in a large number of substituted ethanols which ionise in the pH region, secondary and primary as well as natural fluoral hydrate is unique since it has two acidic products was clearly detected. , both of which can ionise. The pKa for the first ionisation process has been determined to be 10.1 From the kinetic data as well as other supporting evidence, at least four different reaction paths can be distinguished for the permanganate oxidation of fluoral GRADUATE STUDIES hydrate: (i) reaction of the dianion of fluoral hydrate (ii) reaction of the mono-anion, (iii) reaction of the unionised fluoral hydrate with permanganate anion, and Field of Study: Physical-Organic Chemistry (iv) a tentatively proposed reaction of the unionised fluoral hydrate with permanganic acid. Advanced Inorganic Chemistry H.C. Clark The deuterium isotope effects are k^/k^ = 5 for the Surface Chemistry L.G. Harrison Crystal Structures K.B. Harvey, L.W. Reeves oxidation of the di-anion, kH/kn = 10 for the mono- anion reaction, kg/kn = 4 for the neutral molecule High Polymers B.A. Dunell

reaction and k^/kn = 6 for the permanganic acid Molecular Rearrangements A. Rosenthal oxidation of unionised fluoral hydrate in strong acid. Synthetic Methods G.G.S. Dutton Activation parameters and a positive effect show that the transition state formation involves a reaction of similarly charged species; proton tunneling Related Studies: is considered unlikely on the basis of a temperature study. The rate determining step is most consistent Theory and Application of with a hydride transfer to the permanganate ion. Differential Equations C.A. Swanson In connexion with the above study, 2,2,2-tri- Functions of a Complex fluoroethanol, 1,1,1,3,3,3-hexafluoropropanol-2 and the Variable W.H. Simons corresponding 2-d analogue have been examined in some• Structure of Metals E. Teghtsoonian what lesser detail. An isotope effect kH/kn = 20 has been observed for the oxidation of 1,1,1,3,3,3-hexa- fluoropropanol-2.

The pKa of permanganic acid has been determined in sulfuric acid by a spectrophotometric method, pK =-5.1 ii

ABSTRACT

The mechanism of permanganate oxidation of fluoral hydrate and its analogue, substituted with deuterium at position-1, has been extensively investigated throughout the pH region 0.7 to Ik and a successful study was made of the behaviour of permanganate in

the sulfuric acid region from pH 0.7 to 46.3$ H2S04.

Compared with 2,2,2-trifluoroethanol and other substituted ethanols which ionise in the pH region, fluoral hydrate is unique since it has two acidic hydrogens, both of which can ionise. The pKg, for the first ionisation process has. been determined to be 10.1.

From the kinetic data as well as other supporting evidence, at least four different reaction paths can be distinguished for the permanganate oxidation of fluoral hydrate: (i) Reaction of the di- anion of fluoral hydrate, (ii) Reaction of"the mono-anion, (iii)

Reaction of the unionised fluoral hydrate with permanganate anion, and (iv) Tentatively proposed reaction of the unionised fluoral hydra• te with permanganic acid.

The deuterium isotope effects range from kn/kD = 5 for the oxidation of the di-anion, kii/lq) = 1° for the mono-ahion reaction, kn/kD = h for the neutral molecule reaction and kHAD = 6 for the permanganic acid oxidation of unionised fluoral hydrate in strong acid.

Activation parameters and a positive salt effect show that the transition state formation involves a reaction of similarly iii charged species; proton tunneling is considered unlikely on the basis of a temperature study. The rate determining step is most consistent with a hydride ion transfer to the permanganate ion.

In connexion with the above study, 2,2,2-trifluoroethanol, l,l,l;3>3,3-aexafluoropropanol-2 and the corresponding 2-d analogue have been examined in somewhat lesser detail. An isotope effect kg/kp = 20 has been observed for the oxidation of 1,1,1;3,3,3- hexafluoropropanol-2.

The pK_ of permanganic acid has been determined in sulfuric acid by a spectrophotometry method, using two different methods of calculation; pKa = -5.1.

Formic acid has been oxidised by permanganate in the region

25 to 72 $ H2S04 and the reaction in the region from 25 to about kO %

H2SO4 is interpreted as the oxidation of the unionised formic acid by permanganic acid. An isotope effect of about 7 is observed in the very strongly acidic region (72 $ H2SO4), which decreases upon going

to 25 $> H2S04.

Certain 4,4'-disubstituted benzhydrols were oxidised by permanganate in 0.1 M sodium , however solubility difficul• ties prevented any extensive studies.

The structure of fluoral hydrate, synthesised in this labo• ratory, was unequivocally established by n.m.r. to be the dimer of fluoral hydrate.

Spin-spin splitting of the OH proton in large number of secondary and primary alcohols as well as natural products was clearly iv detected.

In the case of trifluoroethanol, two distinct species were observed; a species bonded to the dimethylsulfoxide and the monomeric species of trifluoroethanol which was detected in carbon tetrachloride and tetrachloroethylene. Solvent influence upon the hydroxyl proton shift is discussed.

Conformational analysis of diacetone-galactose was attemp• ted, and the results indicate that the molecule exists in the chair form in tetrachloroethylene.

The effect of substituents upon the magnitude of the hydroxyl proton spin-spin coupling is discussed and the experimental results to that effect reported. V

ACKNOWLEDGEMENTS

Sincere thanks are extended to the research director,

Prof. R. Stewart, for his invaluable help and suggestions during the course of this investigation.

Many valuable discussions with Dr. D.E. McGreer and his encouragement during the course of the n.m.r. projet are most gratefully acknowledged.

The author is greatly indebted to Mrs. E.B. Brion for obtaining all of the n.m.r. spectra. vi

TABLE OF CONTENTS

Page

INTRODUCTION 1

OBJECT OF RESEARCH 17

EXPERIMENTAL 19

1. Reagents Used 19

2. Kinetic Methods 23

a) Titration Methods '. 23

b) Permanganate Oxidations in Deuterium Oxide 27

c) Spectrophotometry Method in the pH Region 27

d) Titration Method for Oxidations in 25 to k6 <$>

Sulfuric Acid 29

e) Spectrophotometry Method in Concentrated Acid 31

3. The pK_ Determination of Fluoral Hydrate 32 k. Determination of Heat of Ionisation of Fluoral Hydrate . 31*-

5. Determination of pK of Permanganic Acid 35

6. Preparation of Benzhydrols 37

a) 4,4'-Dinitroberizhydrol 3^

b) Physical Constants and IR Data for Benzhydrols 38

c) Fluoral Hydrate 39

d) l,l,l;3,3,3-Hexafluoro-2-propanol-2-d hi

RESULTS ^3

1. Permanganate Oxidations of Fluoral Hydrate ^3

a) The Nature of the Substrate ^3 vii

Page

b) Alkaline Hydrolysis of Fluoral Hydrate kj>

c) Stochiometry of the Reaction kk

d) Products Analysis k-5

e) Order of the Reaction V?

2. Permanganate Oxidations of FLuoral Hydrate in the

Region pH 6 to 11 ^5

a) Determination of pK^ of Fluoral Hydrate ^7

b) Enthalpy of Ionisation of Fluoral Hydrate k$

c) Variation of pH of Buffer with Temperature 52

d) Treatment of Kinetic Results 52

e) Deuterium Isotope Effect 57 f) Activation Parameters for the Oxidation of Fluoral Hydrate in the Region pH 6 to 12 58 5. Oxidation of Fluoral Hydrate in Strongly Alkaline

Solutions 59

a) Rate Dependence on Hydroxyl Ion Concentration 59

b) Deuterium Isotope Effect ~. 65

c) Influence of Ionic Strength 65

d) Oxidation by 66 e) Activation Parameters for Oxidation of Fluoral Hydrate in Strongly Alkaline Region 67

f) Oxidation of Fluoral Hydrate at 780 71 k. Oxidation Kinetics in the Region pH 11 to 12,. 5 73

5. Oxidation of FLuoral Hydrate in Weakly Acidic Region, pH 6 to 0.7 76 6. Oxidation of l,l,lj3,3,3-Hexafluoro-2-propanol 77 viii

Page

a) Oxidation of l,l,l;3,3>3-Hexafluoro-2-propanol in

Deuterium Oxide 77

7. Permanganate Oxidation of 2,2,2-Trifluoroethanol 79

8. Permanganate Oxidation of Chloral Hydrate 80

9. Permanganate Oxidation of Substituted Benzhydrols .... 8l 10. Substituents Effect on the Rate of Permanganate

Oxidation of Compounds of the Type CF3CHRiR2 85 11. Permanganate Oxidations in Concentrated Sulfuric Acid 87

a) Determination of pKa of Permanganic Acid 87

b) Permanganate Oxidation of Fluoral Hydrate in 25 to k6 $ Sulfuric Acid 9k c) Activation Parameters for Oxidation of FLuoral Hydrate in 26 to kS $ Sulfuric Acid 95

12. Summary of Permanganate Oxidations of Fluoral Hydrate 99

13. Behaviour of Mn1^ in Concentrated Sulfuric Acid 102

Ik. Permanganate Oxidation of Formic Acid in 20 to 72 $ Sulfuric Acid 109 a) Permanganate Oxidation of Formic Acid in 20 to k2 $ Sulfuric Acid 110 b) Activation Parameters for Permanganate Oxidation of Formic Acid in 20 to k2 $ Sulfuric Acid 115

c) Permanganate Oxidation of Formic Acid in 50 to

72

DISCUSSION AND CONCLUSIONS 120

1. Mechanism of Permanganate Oxidation of Fluoral Hydrate in Weakly and Strongly Alkaline Solutions 120 2. Oxidation Mechanism of Fluoral Hydrate in the Region pH 1 to 6 129 3. Oxidation Mechanism of Fluoral Hydrate in the Acidic Region, up to k6 $ Sulfuric Acid 131 ix

Page h. Mechanism of Oxidation of Formic Acid in the Region 20 to 42 $ Sulfuric Acid 134 5. Mechanism of Oxidation of Formic Acid in the Region 50 to 72 % Sulfuric Acid 136

6. Summary 137

APPENDIX 139

N.M.R. Study of Hydroxyl Proton Spin-Spin Splitting 139

INTRODUCTION 139

EXPERIMENTAL 144

1. Technique of n.m.r. Measurements ihk

RESULTS 1^7

1. Structure of Fluoral Hydrate 14-7

2. Structure of Chloral Hydrate 168

3. Investigation of OH Spin-Spin Coupling in Alcohols 172 k. Solvent Influence on Hydroxyl Resonance 178

5. Study of Hydroxyl Proton Splitting in 2,2,2-Trifluoro• ethanol 179

6. Application of OH Spin-Spin Splitting to Natural Products and Derivatives 183

7. Conformational Analysis 6f Diacetone-Galactose 189

8. Effect of Substituents upon the Magnitude of the OH Coupling 195

DISCUSSION AND CONCLUSION 199

BIBLIOGRAPHY 201 X

TABLES

Page

I The pK*s of CF3CH(OH)2 and CP3CD(OH)2 k9

II Variation of pKg^ of Fluoral Hydrate with Temperature 5°

III pH Variation of the Buffer with Tempe• rature 52

IV Oxidation of Fluoral Hydrate in the Region pH 6 to 12; .Variation of the Rate Constant with pH 3k

V Oxidation of Fluoral Hydrate, Protio and Deuterio Analogues, in the Region pH 6 tp 12; Deuterium Isotope Effect 58

VI Thermodynamic Data for the Oxidation of Fluoral Hydrate in the Region pH 6 to 12 60

VII Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions; Rate Dependence on Hydroxyl Ion Concentra• tion 6k

VIII Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions; Influence of Ionic Strength 65

IX Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions; Oxidation by Manganate ." 67

X Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions; Activation Parameters for the Protio Compound 68

XI Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions; Activation Parameters for the Deuterio Compound 69

XII Oxidation of Fluoral Hydrate at 780 72

XIII Oxidation of Fluoral Hydrate in Weakly Acidic Region,

pH 6.12 to 0.7 76

XIV Oxidation of l,l,l;3,3,3-Hexafluoro-2-propanol 78

XV Permanganate Oxidation of 2,2,2-Trifluoroethanol .... 80 t xi Page

XVI Permanganate Oxidation of Chloral Hydrate 8l

XVII Permanganate Oxidation of -Identically Substi• tuted. Benzhydrols 6%

XVIII Effect of Substituents on the Rate of Permanganate

Oxidation of Compounds of the Type CF3CHRiR2 86

XIX Spectroscopic Data for Ionisation of Permanganic Acid 92

XX Summary of pKa's of Permanganic Acid, Evaluated

with Aid of Ho, H_ and HR (JQ) Functions 93

XXI Oxidation of Fluoral Hydrate in 26 to kS $> Sulfuric Acid 95

XXII Activation Parameters for Oxidation of Fluoral. Hydrate in 26 to k6 Sulfuric Acid 97

XXIII Permanganate Oxidation of Formic Acid in 0.1 to 1.0 M (59,60,6l) and 20 to h2 Sulfuric Acid '.. 110

XXIV Permanganate Oxidation of Formic Acid in 50 to 72 $ Sulfuric Acid 117 XXV Spin-Spin Coupling Constants of the Hydroxyl Protons

in Secondary Alcohols of the Type RjCH0HR2 197

XXVI Spin-Spin Coupling Constants of the Hydroxyl Protons in Primary Alcohols of the Type RCHsOH 198 xii

FIGURES

Page

1. Oxidation of Fluoral Hydrate; Typical Rate Plots kS

2. Oxidation of Fluoral Hydrate; Relation "between Rate Constant and pH hQ

3. Fluoral Hydrate; Variation of pKa with Temperature ... 51

k. Oxidation of Fluoral Hydrate; Relation between Rate Constant and Fluoral Hydrate Anion Concentration 55

5. Oxidation of Fluoral Hydrate; Linear Relation between pH and log kz/ (k2max - ka) • ^6 6. Oxidation of Fluoral Hydrate Mono-Anion; Variation of Rate Constant with Temperature 6l

7. Oxidation of Fluoral Hydrate in Alkaline Solutions; Relation between Rate Constant and [OH"'] 63

8. Oxidation of FLuoral Hydrate Di-Anion; Variation of

Rate Constant with Temperature 70

9. pKa of HMn04; Plot of Optical Density versus H_ 88

10. pKa of Permanganic Acid (Using G. ^ s25) 89

11. pKa of Permanganic Acid (Usong 6 ^ 45g) 90 12. pKa of Permanganic Acid (Modified Davis - Geissman Method) 91

13. Oxidation of Fluoral Hydrate in 25 to k6 <$> H2SO4; Relation between Rate Constant and H_, HQ, HR 96

Ik. Oxidation of Fluoral Hydrate in Ul.8 $> H2SO4; Variation of Rate Constant with Temperature 98

15. Oxidation of Fluoral Hydrate; Relation between Rate Constant and pH, H_ 100

16. Oxidation of Fluoral Hydrate; Relation between loga• rithm of Rate Constant and pH, H_ 101 xiii

Page

IV 17. Spectrum of Mn in 62$ HaS04 104

IV 18. Spectrum of Mn in 97$ H2S04 105

111 19. Spectrum of Mn in 62$ HaS04 107

20. Oxidation of Formic Acid in Perchloric and Sulfuric

Acids; Plot of log k2 versus H- 112 21. Oxidation of Formic Acid; Relation Between Rate . Constant and h_ 114

22. Oxidation of Formic Acid in 38$ H£>S04; Variation of Rate Constant with Temperature , 116

23. Oxidation of Formic Acid in 50 to 72$ H2S04;

Relation of log k2 with HR 119

24. H1 Spectrum of Fluoral Hydrate 150

25. H1 Spectrum of Fluoral Hydrate After Standing 151

26. H1 Spectrum of Fluoral Hydrate (+ 10 ul. Hj.0) 152

27. F1^ Spectrum of Fluoral Hydrate 153

1 28. H Spectrum of Fluoral Hydrate (+ 39 ul. H20) 155

29. H1 Spectrum of Fluoral Hydrate 156

30. F1^ Spectrum of Fluoral Hydrate 157

31. H1 Spectrum of Fluoral Hydrate (+ 30 pi. Hs-O) 159

32. F1^ Spectrum of Fluoral Hydrate (+ 30 pi. HJJO) 160

1 33. H Spectrum of Fluoral Hydrate (+ kO jul. H20) l6l

1 34. F ^ Spectrum of Fluoral Hydrate (+ kO jul. H20) l62

35. H1 Spectrum of Fluoral Hydrate (+ 160 jul. Hj.0) 163

1 36. F ^ Spectrum of Fluoral Hydrate (+ 160 ul. H20) 164

37. H1 Spectrum of Fluoral Hydrate After Standing 165 xiv

Page

1 1 38. H Spectrum of CF3CD(OH)2 169

1 39. H Spectrum of Chloral Hydrate in CDC13 170 kO. E1 Spectrum of Chloral Hydrate in Acetone 171 kl. H1 Spectrum of Chloral Hydrate in Acetone (+ 10 ul.

HaO) 173 k2. H1 Spectrun of Isopropyl Ijk

4-3. H1 Spectrum of CF3CHOHCP3 in Acetone 175 kk. H1 Spectrum of CF3CHOHCF3 in Acetone - DMS 176

1 1+5. H Spectrum of PhCH0HCF3 in DMS 177 k6. E1 Spectrum of Trifluoroethanol in DMS l8l

1 kj. H Spectrum of Trifluoroethanol in CC14 182 kS. H1 Spectrum of Diacetone-Glucosej Ring Protons iQk

4-9. H1 Spectrum of Diacetone-Glucose (Ring Protons and OH) 186

50. H1 Spectrum of Citronellol (Neat) 187

51. H1 Spectrum of Citronellol in DMS 188

52. H1 Spectrum of Diacetone-Galactose (Ring Protons) ... 190

53. H1 Spectrum of Diacetone-Galactose (Ring Protons) ... 191 1

\ INTRODUCTION

Potassium permanganate was first used by Margueritte (l) in lb%6 in the oxidation of ferrous ion. Even today - 116 years later - there is still, some doubt about the actual mechanism of some of the permanganate oxidations.

Permanganate can undergo reduction to lower valency states in three well known ways. In the weakly acidic solutions, the overall five electron change may be represented by the equation

+ 2+ Mn04~ + 8H + 5e Mn + 4H20

In the weakly acidic, neutral or alkaline solutions the permanganate can undergo a three electron change

+ Mn04~ + 4H + 3e ===== Mn02 + H20 and in the alkaline solutions (pH 12 to Ik) permanganate reduction may be arrested at the manganate stage without any further noticeable reduction of the manganate

Mn04~ + e <==== Mn04~

Generally, inorganic substrates appear to reduce acid permanganate to its lowest Mn2+ state, whereas organic substrates con•

taining oxidizable hydrogen seem to stop the reduction at the Mn02 stage.

Study of the stability of permanganate solutions has re• ceived considerable attention in the past and has been recently summarized by Belcher (2). Permanganate is stable for months in neutral solutions when kept in the dark (2,3) but decomposition is 2 noticeable in the acid (4,5) or in the alkaline solutions (2,6,7).

Light, manganous salts and dioxide also accelerate permanganate decomposition. Morse, Hopkins and Walker (8) postulate the reaction

2HMn04 + Mn02 (x + 2)Mn02 + 30 + HJZO which indicates the autocatalytic nature of the process. It should be noted that the species under consideration is permanganic acid and not the permanganate ion.

Tne purple permanganate ion absorbs strongly in the visible part of the spectrum and following.its disappearance by spectral means constitutes one of the important methods of study of the permanganate oxidations. The absorption spectra and electronic structure of the permanganate ion have been extensively studied by various workers (3,9,10,11) and its tetrahedral configuration has been confirmed by Wolfsberg and Helmholtz (12).

Of interest are the reactions of the lower oxidation states of manganese which might be helpful in elucidating the mechanism of permanganate reduction, undoubtedly, the transfer of more than two electrons simultaneously would be considered a most unlikely process on energetic grounds and thus it might be anticipated that permanganate should pass through some intermediate stage before being reduced to its final state.

The behaviour, stability and oxidizing power of manganate were investigated extensively by Waters and co-workers (11,13) and others (3,14,15). Manganate" is stable only in alkali and solutions 3 less than IM in hydroxyl ion slowly disproportionate into permanganate and Mn02 (3,l6,17K However, the exact pH value or the region at which this occurs has never been precisely defined or investigated.

3 The blue ion (Mn04 ) was first prepared by

Lux (l8) and its oxidizing properties and stability have been investigated by Pode and Waters (ll). Below 8M NaOH hypomanganate disproportionates into manganate and . The oxidizing

x power of the manganese of the type Mn04 decreases in the

2 3 order Mn04~ > Mn04 > Mn04 ~. This is explained on the basis that the oxidation process involves acceptance of electrons by the anion, whereby the movement of the electrons is contrary to the charge on the anion (l6). Wiberg and Stewart (19) found, however, that both manganate and permanganate oxidize substituted benzaldehydes at the same rate. Since the formation of hydroxyl radicals is postulated in the oxidation process, it may well be that induced oxidation is occurring in this case. This may account for the same rate of oxidation by both manganese species.

The cation Mn3+ is Stable in concentrated sulfuric acid (20) however, at lower acidities fast takes place:

2Mn3+ =5==== Mn4+ + Mn2+

It is, however, possible to stabilize the solutions of Mn3+ by complexing agents such as pyrophosphates which make them more suitable as oxidizing agents in the pH region. Waters and his collaborators have investigated extensively the Mrr1-11 Gxidatidns of aliphatic , , 1,2-diols, o< -hydroxy acids-and other organic k

compounds (21,22,23).

Manganese dioxide has long been used by the organic chemist for oxidations. These reactions, usually carried out in the presence of sulfuric acid or in aprotic solvents, are of heterogeneous nature.

No homogeneous oxidations involving Mn*V are as yet known (l6). In view of the results obtained in this investigation, this statement will be mentioned further in the text. The composition of precipitated manganese dioxide varies depending upon the conditions used (24). It was pointed out from analogy with other transition metal

that Mn(OH)4 would be too weak an acid to exist in the ionized form

in any solution and therefore it dehydrates to the more thermody•

namically stable Mn02 (l6).

In the weakly acidic or neutral solutions manganous react with permanganate to form manganese dioxide. This reaction was first utilized by Guyard (25) for the determination of manganese.

2+ + 2 Mn04~ + 3Mn 2H2O *?=^ 5Mn02 + kE

The reaction as written is probably not a truly reversible equilibrium

(l6). The mechanism of the Guyard reaction has been investigated quite extensively in the past and has been comprehensively reviewed by

Ladbury and Cullis (26); however, up to the present the mechanism is far from established. The main difficulty appears to be the fact that

one deals with a heterogeneous reaction involving solid Mn02 and also

that other manganese species of intermediate valency are formed in

the reaction. These might be responsible for the actual course of the reaction. 5

Another important reaction of manganese oxyanions is that

between permanganate and hypomanganate:

3 2 Mn04~ + Mn04 ~ =3=^ 2Mn04 "

This reaction is fast (3) and is the reason for the inability to

observe any transition state involving Mn^ in the permanganate oxida•

tions in the alkaline solutions (11).

2 All three of the manganese oxy-anions, Mn04~, Mn04 ", and

3 Mn04 ~ are of course conjugate bases of their corresponding acids,

and they would be expected to appear in the unionized form somewhere

in the pH region. If the protonation of these anions occurs in the

region other than that defined by the pH, the the pr^ of the acid must follow the H_ function, in accordance with.the postulates first

put forward by Hammett (27). The permanganic acid is usually con•

sidered to be a strong acid, however, surprisingly very little is

actually known about its strength (e.g. pKa). It has been known for

a long time that a green solution is obtained when permanganate is

dissolved in concentrated sulfuric acid (28). The green colour has

been attributed by Franke (29) to the formation of the species

(Mn03)2S04 and by Ephraim (30) to the formation of Mn207. However,

the evidence confirming either was lacking. Symons and co-workers

(31) have recently measured the pKa of permanganic acid by observing

the changes in spectra of the Mn04" anion with increasing acidity.

They have ascribed the green colour to permanganic acid and substantiated

their argument by showing that Beer's law was obeyed when the con•

centration of permanganate was varied. Perchloric acid was used as 6 the protonation medium and spectral changes were recorded only in the region where the protonation was slight, since permanganic acid pre• sumably decomposes at higher concentrations. They have also established for several ions that changes in solvent and ionic strength did not detectably alter the spectra. The spectral data were evaluated by a method previously described by Hammett and Deyrup (32), and the value obtained for the pKa of permanganic acid is given as -2.25. It should be pointed out that the H- function was first evaluated in the acidic region by Boyd (33) in 196l for a series of cyanoethylene type compounds in perchloric and sulfuric acids and by Phillips (3*0 for hydrochloric acid. The B_ function has been known for some time in the alkaline region (35,36,37,38) and a self-consistent H_ scale has been recently reported by Stewart and O'Donnell (39) and Edward (40).

More recently, Royer (kl) has made cryoscopic studies on the green solution of permanganate in 99*9$ sulfuric acid. Even though the precision of the data is not striking and variation of the cryoscopic V factor with concentration is observed. Royer postulates that at least in dilute solutions the number of particles responsible for the freezing point depression is 6. From this, as well as from extrapolated conductivity measurements, the existence of the species

+ Mn03 (permanganyl ion) is postulated. According to Royer, this species arises from the reaction:

+ + + KMn04 + 3H2S04 s ^ Mn03 + K + 3HS04" + E^)

The spectrum of this species, taken in 98$ sulfuric acid is almost coincident with the spectrum reported by Symons.(31) for permanganic 7 acid. An attempt was also made by Royer to predict theoretically the

+ expected spectrum for Mn03 .

Attempts by Symons (31) to measure the basicities of manganate and hypomanganate anions failed, due to the decomposition of both anions into permanganate and manganese dioxide before any protonation could be detected. Waters (16) however, advanced the interesting

suggestion that the ions HMn04~ and HMn04~ might be involved in the decomposition of the Mn^1 and MnV species respectively. Symons (31)

3 2_ estimates that the pK^ for the system Mn04 ~/HMn04 is less than

= ih.h and for Mn04 /HMn04" less than 12.

When permanganate is dissolved in 65$ oleum, is evolved rapidly and the resulting solution is blue. This colour is ascribed by Symons (h2) to a quadrivalent species of manganese, possibly the uncharged sulphate. When MnC-2 is dissolved in oleum, no evolution of oxygen or sulfur dioxide is observed and the solution is identical with those described above. Apparently the same species was reported originally by Lankshear (4-3) whose observations were confirmed by

Royer (hi) who attempted cryoscopic studies on the blue solution, but without any success.

Isotopic Exchange Reactions.

The exchange between Mn04" / Mn04~ in the alkaline solutions was extensively studied by Wahl and co-workers (hh,h^,h6). The values of the heat and entropy of activation for this process are reasonable for an electron transfer occurring by an outer-sphere mechanism (hf).

On the basis of the free energy of-activation they concluded that the 8 electron tunnelling hypothesis of Marcus, Zwolinski and Eyring (k8)

was not a reasonable model for the Mn04~ / Mn04~ exchange. Interesting influences of various cations were found; thus the sequence kCsOH yk KOH > kUaOH » kLiOH was ascribed to the incorporation of the cation into the activated complex to form a bridged structure

[Mn04-K-Mn04] ''" , analogous to the anion bridge structure presented by

Taube and collaborators (4-9).

Happe and Martin (50) have investigated the exchange between

2+ Mn / Mn04~ in 1 to 2M nitric acid. There the well known Guyard reaction takes place and hydrous oxide is obtained containing Mrr^.

Rapid, 30$ exchange occurred between Mn2+ and freshly precipitated

iv 2+ Mn . The exchange between Mn / Mn04" was also detected, however, due to other reactions this exchange may not be real. Adamson, however,

2+ observed a measurable exchange between Mn / Mn04~ in 3M perchloric acid (51). In this case, the appearance of the precipitate of Mn*^ was delayed for about 10 hours.

The stable isotope of oxygen 018 was used with advantage by many workers in the investigations of the mechanism of permanganate oxidations. Oxygen - 18 may be incorporated into permanganate by equilibrating permanganate and 018 enriched water for several hours at 100C° (9,52). The corresponding exchange in neutral or weakly alkaline solutions is slow at room temperature and this enables the tracing of the route of oxygen during the course of the reaction.

It was suggested that the mechanism of the 0 exchange involves the protonation of permanganate followed by the replacement 9 of the hydroxyl group (52). In 0.2M the 018 exchange

is increased (53) and in strongly alkaline solutions permanganate

undergoes a complete exchange of its with RaO18 (5^). Manganate

also undergoes exchange with the solvent and for this an type mechanism was suggested, involving protonated manganate:

OH" + HOMn03" ======HOMn03" + OH"

However, no exchange occurred between manganate and the solvent when barium salts were added and manganate precipitated as the insoluble barium manganate. It was concluded that permanganate does not exchange

its oxygens with the solvent in the strongly alkaline solutions, and

the isotopic enrichment of permanganate was explained on the basis of

the rapid reaction between manganate and permanganate described above

Oxidation of Formic Acid.

The oxidation of formic acid by permanganate has received

considerable attention in the pastj however, the progress on the mechanism was made only relatively recently.

The first systematic investigation was undertaken by Holluta,

covering the pH region up to IM sodium hydroxide (55)- The bimolecular

nature of the reaction was clearly recognized and the accelerating

effect of the fluoride ion was noted. In the strongly alkaline

region, manganate reacted with formic acid at a much slower rate than

permanganate.

Tompkins and collaborators (56,57) confirmed the second order

of the reaction and in addition noted a positive primary salt effect 10 which indicated a transition state composed of two similarly charged species. From this evidence, the conclusion was reached that the formate ion was the reactive species, and the following mechanism was postulated:

Mn04" + HCOO" i •> Mn03~ + C02 + OH" slow

2Mn03~ + Hs-0 > Mh02 + BMn04~ + OH" fast

HMn04" + HCOO" > HjjO + C02 + 20H" fast

Addition of manganous ions caused a decrease in the rate and this indicated clearly that Mn*1* or Mn*V species played little role in the oxidation. The energy of activation and negative entropy of activation were in accord for a reaction between two similarly charged species.

Wiberg and Stewart (58), using the deuterium labeling technique found a substantial isotope effect kn/kD = 7«4 for the reaction of formate ion with permanganate. This shows clearly that the C-H bond is broken in the rate determining step. To account for this observa• tion either a hydride ion or electron transfer to permanganate is postulated. However, when the experiments were carried out with 0X8 labeled permanganate, a definite (i8 to 30$) oxygen transfer to the product occurred. Several possible mechanisms were considered which would explain the 01S incorporation into the product: 11

-9

"O-C-H + Mn04" O-C-H (+J HC03 + Mn03"

0

Mn03

The objection to this proposal was the fact that the reaction between permanganate and undissociated formic acid is much slower, although it would be expected that the above intermediate would form easier with the latter species. Another possibility was discussed which would lead to both labeled and unlabeled carbonate, depending on which bonds were broken:

0 W * ls la 0 -Mn03 0 Mn03" "0-C "0-C- / H

. II 18 C02 + H0 -Mn03 O-C-OJ-SH + Mn03"

Halpern and Taylor (59,60,6l) have extended the oxidation of formic acid into the acidic region where the formic acid exists pre• dominantly in the undissociated form. The contribution of the un• ionized formic acid was detected and found to be unaffected by the

change of solvents from H2O to D20. Furthermore, no deuterium isotope 12 effect was observed for the reaction of HCOOH and DCOOH. In agreement with the findings of Wiberg and Stewart (58) and Aebi, Buser and

Luethi (62) a considerable isotope effect of 7 was observed for the formate ion reaction. The entropy of activation for the undissociated formic acid - permanganate reaction appears unusually large and it may reflect the increased charge separation in the activated complex.

Oxidation of Benzaldehyde.

The benzaldehyde oxidation was examined by Tompkins (63) who

0" I concluded that the ion CeH5-C-H is responsible for the reaction with I OH permanganate.

A more detailed study was undertaken by Wiberg and Stewart

(6k) using the deuterium and 018 labeling. A substantial isotope effect was found in the weakly acidic region; this effect decreased with increasing pH, as did the 018 transfer from permanganate to substrate. The reaction was found to be general acid catalysed. A

Hammet plot for the reaction in the weakly acidic region showed a good linear relation with a small negative value. A mechanism which fits the observed data was as follows:

. l

OH I

RCHOH + Mn04" EC — OMn03 I H

OH I k

RC — OMn03 I H :B ===> RC02H + HB + Mn03

fast

3Mn03" + H2O ======> 2Mu02 + Mn04 + 20H

In the alkaline solutions the rate was found to depend on the square root of the hydroxyl ion concentration, and most, if not all of the oxygen introduced into the was derived from the solvent.

The deuterium isotope effect was small and the ^> value found for the Hammett plot was large and positive,- In order to explain the dependence of the rate on the square root of the hydroxyl ion concentra• tion, a free radical chain mechanism involving OH radicals was suggested.

Oxidation of Benzhydrol.

This substrate which yields an oxidation product stable to further oxidation was examined thoroughly by Stewart (17). The use of deuterium substitution and 018 exchange enabled him to show that the C-H bond is broken in the rate determining step and that none of the oxygen is introduced into the substrate by permanganate. The rate showed a dependence on the hydroxyl ion concentration and a positive

salt effect. On the basis of the above observations, a most likely mechanism was postulated, involving a hydride ion transfer from the Ik benzhydrylate anion to permanganate: slow (C6H5)2CH0- + Mn04 » (C6H5)2C = Q. + HMn04

fast + HMn04~ + Mn04 > 2Mn04" + H

Oxidation of Fherryl-trifluoromethyl carbinols.

In the order to test the validity of the assumption that the alcoholate anion is the species which reacts with permanganate, a series of substituted phenyl-trifluoromethyl carbinols was subjected to further study (64,65). These alcohols ionise in the pH region'.and their pKa's can be easily determined (66). If the assumption that the alcoholate anion is the reactive species is true, then it should be possible to observe a leveling off of the rate in the region where the alcohol is almost fully ionised, and moreover the rate should follow the pKa of the alcohol. This has in fact been Confirmed.

Further observations of a positive salt effect, large negative entropy of activation and a large deuterium isotope effect, JIHAD = 16, point to a mechanism which involves a hydride ion transfer. The effect of substituents is however only slight, and some other mechanisms involv• ing a termolecular step were considered. The large deuterium isotope effect is possibly due to the loss of the stretching as well as bend• ing modes in the transition state; however a possibility of the quantum mechanical proton tunnelling is not excluded.

Oxidation of Cyanide Ion.

The reaction of the cyanide ion with permanganate was found to yield cyanate~in the region pH 12 to lk.6, however below pH 12 15 various other products were obtained (65,67). Two reaction paths may• be distinguished; one, appearing at low reactant concentrations, for which the mechanism may be represented as follows:

k p- Mn04" + CN" (NC0Mn03) ====> Mn03" + NCO. and the other, detectable at higher reactant concentrations, proceeding probably by a formation of the species H(CN)2~ which reacts with permanganate by a hydride ion transfer.

Oxidation of Molecular Hydrogen.

The kinetics of the homogeneous reaction between permanganate and molecular hydrogen were investigated in the acidic and basic regions by Webster and Halpern (68). The rate varies only slightly with the acidity and the absence of the salt effect is consistent with a reaction between an ion and uncharged molecule. The proposed mechanism v favours the formation of Mn species by a hydride ion transfer on the basis of the thermodynamic data:

+ Mn04~ + H2 '•• > Mn03" + 2H

Oxidation of Borohydride ion.

This reaction was studied in the basic solutions by Freund and Nuenke (69). Even though the reaction appears to be somewhat complex, the following observations were made: (a) there was an absence of a strong hydroxyl ion dependence, (b) the reduction of borohydride with manganate alone was found to be slow, (c) when permanganate was reacted with borohydride, it was impossible to stop 16 the reaction at the manganate stage even with the addition of barium ions, (d) all of the hydrogen evolved originated from borohydride. On the basis of these observations it was suggested that the mechanism might involve hypomanganate as an intermediate.

Oxidation of Olefins.

In 1922, Boeseken (70) has postulated the stereochemical scheme for the oxidation of olefins to cis-glycols.

V 0 0 \- 0. . 0 \-0E

\,/ 2 H20 + Mn Mn ====#. + H2Mn04~

Cv 0^ ^0 C 0 0 C-OH / \ / \ / \

No direct evidence for Boeseken1s intermediate was available until

Wiberg and Saegebarth (71) applied the 018 labeling technique. Using

018 labeled permanganate they have oxidized potassium oleate to the diol and found that up to 1.5 atoms of oxygen can be transferred from permanganate to the olefin molecule. This shows very clearly that the

Boeseken1s intermediate must have two C-O-Mn bonds which can hydrolyse between the manganese and oxygen atmos.

It is clearly apparent that almost all of the more modern mechanistic approaches to the permanganate oxidations involve the MnV species which may arise either by a hydride transfer from the substrate to permanganate or by an oxygen transfer from permanganate to the substrate (cf. however ref. 7)' 17

OBJECT OF RESEARCH

This project was undertaken with the primary view in mind of establishing the nature of the reactive species in permanganate oxidations. A simple system was sought in which the products would be stable to further oxidation. The k,k -di-substituted benzhydrols fitted these requirements well. Their study would also represent an extension of the previous work on benzhydrol (17) and perhaps suggest a reason for the small substituent effect which was observed for the phenyl-tri-fluoromethyl carbinols (64,65). As a logical continuation it was decided to study the system RCHOHCF3 where the electronic properties of the R group could be varied more greatly then is possible with the benzene ring substituents. For this reason trifluoroacetal- dehyde, trifluoroethanol, and hexafluoro-2-propanol were selected as the most likely candidates. The advantage of using these compounds lies in the fact that they all ionise in the pH region and thus it is possible'to study the reaction of the neutral molecule as well as the ionised species.

Quantitative investigations of permanganate oxidations in the strongly acidic solutions have remained essentially unknown up to the present. The reason for the apparent reluctance of the investigators to use permanganate as an oxidant in this region is not immediately obvious even though it may be attributed partly to the decomposition of permanganate,(permanganic acid) which would complicate any quantita• tive studies. 18

Fluoral hydrate and formic acid were chosen as suitable compounds, since their behaviour in strong acid may be reasonably well predicted, and is well known, in the case of formic acid. Two objectives would be accomplished by this work: (i) the mechanism of permanganate oxidation of formic acid is well known in the pH region.

Investigation of this reaction at high acidity would essentially complete the knowledge of behaviour of formic acid towards permanganate,

(ii) Extension of permanganate oxidations into the strongly acidic region would provide basis for further investigations and also indicate whether there exist any similarity between the mechanisms of chromate and permanganate oxidations. 19

EXPERIMENTAL

1. Reagents Used

Benzhydrols.

Benzhydrol and other k,h -disubstituted benzhydrols and their deuterium analogues were prepared by the reduction of the correspond• ing ketones with HAIR* or LiAlB-4. The purity was checked by the constancy and sharpness of the melting point and the isotopic purity was estimated by infra-red and in some cases by nuclear magnetic reasonance measurements. The stock solutions of benzhydrol and benzhydrol o(-d were prepared in the conventional manner by weighing out a quantity of the compound and dissolving it in the warm water

(17)• However the solubility of the substituted benzhydrols decreased markedly compared to benzhydrol itself. Thus the stock solutions of these were prepared by dissolving a suitable quantity in warm water and allowing to stand overnight. The excess of the undissolved com• pound was filtered off and the concentration was determined spectro- photometrically by comparison with a sample of known concentration.

Fluoral Hydrate.

The stock solution was prepared by weighing out a suitable quantity and dissolving in water. It was later discovered that fluoral hydrate exists In the form of a condensation dimer and thus a correction was applied to some of the kinetic data. The purity of this compound was checked by the melting point, infra-red, and n.m.r. spectroscopy, 20 carbon hydrogen analysis and molecular weight determination.

I,l,lj3,3,3-Hexafluoro-2-propanol.

A sample of this compound was generously donated to this laboratory by Dr. S. Andreades of Dupont de Nemours Company, Wilmington,

Delaware. The purity was checked on "Aerograph" vapour phase chromato• graph and was found to give a single sharp peak. The proton magnetic reasonance gave the expected spectrum (see Appendix), and the infra-red

C-H absorbtions were noted at 2960 and 1377 cm1. The stock solution was made up by dissolving a weighed quantity of the alcohol in water.

The deuterio analogue was synthesised by a method described elsewhere in the text and it was purified by the vapour phase chromatography.

The infra-red spectrum showed C-D stretching and bending frequencies and the complete absence of C-H absorptions. The n.m.r. spectrum showed only one peak due to the 0-H absorption.

2,2,2-Trifluoroethanol.

This was purchased from Columbia Chemicals Company and was purified by the vapour phase chromatography.

Formic Acid.

Stock solutions were prepared by dissolving a Baker and Adamson sample. . df sodium formate in the sulfuric acids of appropriate concentra• tion just before the kinetic runs. An appropriate correction was made to account for the decrease in the sulfuric acid concentration when necessary. 21

Chloral Hydrate.

A commercial sample was recrystallized from ether until the melting point was constant. The purity was checked by n.m.r. which gave the predicted spectrum (see Appendix).

Potassium Permanganate.

A reagent grade sample was used. The stock solutions were made up by dissolving a quantity of permanganate in boiling distilled water and filtering the hot solution through fine sintered glass crucible and filling to the mark with water (72). Permanganate solutions were standardized with sodium oxalate (2,72) and further by the iodometric method and stored in the dark.

Potassium Manganate.

Potassium manganate was prepared by the method of Pode and

Waters (ll). The stock solutions were made up in 1.0M sodium hydroxide and the concentration determined by an iodometric titration and spectrophotometrically. The solutions were kept in the freezing compartment of the refrigerator where, it was found, the solutions were stable for several months.

Solutions of MnIv.

The stock solutions of MnIV were prepared by combining a known volume of permanganate with concentrated sulfuric acid and allow• ing the decomposition to take place. The,concentration of the MnIV species was determined by the iodometric titration of a weighed aliquot and spectrophotometrically. The stock solutions were stable 22 up to a period of several days.

Sulfuric Acid.

Fisher's "Analar" grade sample was used, containing not more than 0.00001$ sulfur dioxide. Solutions of varying amounts of sulfuric acid were prepared "by dissolving the appropriate amount of the acid in the doubly distilled water. The concentration of sulfuric acids were determined by titrating a weighed quantity with standard sodium hydroxide. Same sulfuric acid solutions were used for the spectrophotometric and kinetic experiments.

Water.

Ordinary distilled water was redistilled from alkaline potassium permanganate under conditions which prevented absorbtion of carbon dioxide. This doubly distilled water was used for the prepara• tion of all stock solutions and in all kinetic experiments.

The deuterium oxide was prepared by distillation of a commercial sample from potassium permanganate to which a small amount of calcium oxide and anhydrous potassium carbonate were added. Mixing with atmospheric water and carbon dioxide was prevented during the distillation.

Potassium , dipotassium hydrogen phosphate and sodium thiosulfate were Baker and Adamson "Analar" grade. 23

2. Kinetic Methods

Essentially two kinetic approaches were used. The one in•

volved the conventional aliquot withdrawal, quenching with the

acidified potassium iodide solution and backtitration of the liberated

iodine with standard thiosulphate. The other was a spectrophotometry method.

(a) Titration Method

All solutions used in a particular kinetic run were first

thermostated at a predetermined temperature of the run. This was

usually 25.O t 0.03°. In a typical run 25.0 ml. of 0.6 M potassium

sulfate, 5.0 ml. of 1.0 M sodium hydroxide, 5.0 ml. of 0.0108 M

fluoral hydrate and 10.11 ml. water were combined and mixed in a 125

ml. Erlenmeyer flask, which was provided with a ground glass stopper.

About 30 seconds later the reaction was started by addition of 4.89 ml. of 0.0216 M potassium permanganate from a fast delivery pipette.

The delivery time was 3.5 seconds. Care was taken to employ the same

mechanical procedure every time during the delivery of permanganate.

The reason for the immediate addition of permanganate, after all

other reagents were mixed, was to avoid any possible hydrolysis of

the substrate. The time between the mixing of the reagents and

addition of permanganate was the same in all runs. After suitable

times, 4.92 ml. aliquots were withdrawn with another fast delivery

pipette and passed into the quenching solution containing acidified

potassium iodide. The delivery time of this pipette was 4.5 seconds 2k and this had to he taken into account in the reactions with a half life shorter than 15O seconds. It was found that the acidified potassium iodide oxidized upon standing and as a consequence the titer of the thiosulfate was greater and the reaction appeared slower. There• fore the following procedure was adopted. The sulfuric acid was pipetted into an Erlenmeyer flask and a few crystals of potassium iodide and about 20 mgm of sodium bicarbonate were added to it just before the aliquot was to be quenched. The concentration of the acid was kept at a minimum in order to avoid any oxidation of potassium iodide. In the case of reactions with half lives in the order of 100 seconds, the aliquots were withdrawn at suitable intervals, quenched, stoppered and stored in the dark and then titrated in a random order within about 10 minutes after the last aliquot was withdrawn. The titration time was maintained at 60 seconds; it took about kO seconds to complete the titration and a further 20 seconds were allowed for the burette to drain before the reading was taken. It was found that such a standardization of the procedure gave excellent reproducibility.

A 25$ solution of Thyodyne indicator was used for the end-point determination; half ml. of this solution was added just before the completion of the titration. The addition of the same amount of indicator assured the same end-point. VII

In the pH range from 1 to about 10.5 the Mn is reduced to

MnIV by all the organic reagents studied herein; e.g. a three equivalent change occurs. The oxidation of a secondary alcohol or aldehyde involves a two equivalent change and this sets the ratio of the substrate 25 to permanganate as 3 : 2. In the more strongly alkaline solutions, above pH 12, the permanganate is reduced to manganate which accumulats and reacts with the substrate at a negligible rate compared with the rate of oxidation by permanganate. Since here a one equivalent change occurs, the ratio of the substrate to permanganate now becomes 1:2.

In the region between pH 10.5 and 12 some intermediate ratio applies.

However this is not easily determinable since any manganate formed during the reaction does not immediately disproportionate into permanganate and manganese dioxide. Manganate is stable above pH 12 and completely unstable below pH 10.5; however in the intermediate range the disproportionation of manganate appears to be time as well as pH dependent, and therefore precise information about the oxidation rate cannot be obtained. Stochiometric ratios of the substrate to permanganate of 3 : 2 and 1 : 2 below pH 11 and above pH 12 respectively were used throughout this investigation.

Integrated rate expressions for a second order reaction were derived by Stewart (17,19) for the permanganate oxidations in the alkaline solutions where thiosulphate is used in an iodometric titration and the stochiometric concentration of the reactants is present throughout the oxidation:

k2 = 1 x Vp - Vt

[Alcohol] o t Vt - 4/5 V0 where Vo = initial volume of thiosulfate at "t = 0

Vt = volume of thiosulfate used at time t 4/5 Vo = final volume at t^ = Vco , (usually calculated) 26

t = time seconds

[Alcohol] o = initial concentration of substrate at t = 0

k2 = second order rate constant

A similar expression is found for the ratio 3 : 2

k2 = 1 x Vo - Vt

[Alcohol] o * t Vt - 2/3 V0

The plot of (Vo-Vt)/(Vt-V5 V0) and (V0-Vt)/(Vt-2/5 V0) versus time gave straight lines.

In the investigation of the permanganate oxidations in the pH range 0.7 to 6.2 it was necessary to increase greatly the concentra• tions of the reactants due to the slowness of the reaction. A further complication resulted from the precipitation of manganese dioxide which caused erroneous titres. For this reason a different kinetic pro• cedure was adopted here. The total volume of the kinetic solutions used was 10.0 ml. out of which nine one-mililiter aliquots were pipetted into separate containers which were stoppered, then immersed into the thermostated bath and allowed to react. Light was excluded by wrapping the containers into a tin foil. After a suitable time one of these containers was open and the entire contents, including pre• cipitated manganese dioxide, were quenched in the usual manner and titrated with thiosulfate. Consistent results were obtained by this method.

Similar procedure was adopted for the high temperature runs.

It was^ anticipated that the reaction would be fast and for this reason pH of 6.12 was chosen where the half life of the reaction is about 3 27 days at 25°. The individual aliquots were sealed under nitrogen and stored in dry ice. The time involved in the preparation of the reaction mixture and sealing of the aliquots did not exceed 30 minutes.

After the completion of the reaction at 780, the aliquots were quenched in dry ice and then processed in the usual manner.

(b) Permanganate Oxidations in Deuterium Oxide

Hexafluoro-2-propanol was oxidized in deuterium oxide as a solvent. The following procedure was applied in this experiment.

Sodium hydroxide (0.204 g, 98.5$), potassium permanganate (0.029k g) and potassium sulfate (2.18 g) were weighed out and dissolved in about

40 ml. of deuterium oxide. The reaction was started by addition of

2.94 ml. of 0.0322 M hexafluoro-2-propanol (solution in water) and the volume was quickly made up to 50 ml. with deuterium oxide. The reaction was then followed in the usual manner. Exactly the same procedure was repeated in another run in which distilled water was substituted for the deuterium oxide. Both runs were made at 250 and at ionic strength 1.0.

(c) Spectrophotometric Method in the pH Region

This method was used mainly in the oxidation of substituted benzhydrols which were not soluble enough to allow the use of the titration method.

The reaction mixture was made up essentially the same way as in the titration method, by combining the previously thermostated reaetants. After the addition of permanganate-a small portion of the 28 solution was quickly transferred into a cell of one cm path length and placed in the Beckman DU spectrophotometer, which was provided with thermospacers in order to keep the temperature constant. A blank cell containing all the reagents with the exception of permanganate was used to compensate for any absorbtion caused by the reagents. The optical density was followed with time at 522 and 426 mu , The absorbtion at the former wavelength is due mainly to permanganate and at the latter due essentially to manganate. At 522 mu the extinction coefficient of permanganate is 2570 and that of manganate 587. Similarly at 426 m/u the extinction coefficient of permanganate is 77 and that of manganate 1370 (9)• Using these data, the expression for the con• centration of permanganate may be derived:

[MQ04-]t = Ds22 - 0.282 D426 25413 where D stands for the optical density. The second order rate constant is given by:

k2 = 2 1 - 1

_ t [[Mn04~]t [Mn04 ] o.

The plot of 1/ [ Mn04"] t versus time gives straight line, and the rate constant may be calculated from the slope of the line. Good second order rate plots were obtained up to about 50$ reaction however there the situation was complicated due to the precipitation of the product () out of the solution. This was observed visually on the bulk of the solution and was noted by a sudden break in the plot of the optical density against time as well as in the plot of

1/ [Mn04-]t . 29

It is assumed that manganate is stable and is not reduced further to some MnIV species which might interfere with the measure• ments of the light absorbtion. It must also be assumed that no other species is formed during the reaction which might absorb at the wave• lengths measured. Since a reasonable agreement was obtained for the titration and spectral methods, at least in the case of unsubstituted benzhydrol (17), there is an indication that the spectral method is reliable.

Another complication arose when the solubility of some of the benzhydrols was so low that stock solutions of only 1 x 10~4 M and less could be prepared. It was found that the oxidation of the blank

(benzhydrol absent) made a considerable contribution to the rate if the oxidation time was more than five minutes. Since the rate of oxidation of blank was neither clean first or second order, the best which could be done under the circumstances, was to add the optical density of the blank to the optical density of the sample at 522 mu and subtract it at k26 mu.

(d) Titration Method for Oxidations in 25 to k6jo Sulfuric Acid

In the sulfuric acid solutions permanganate is reduced to

MnIV species which does not precipitate but stays in solution. This makes it rather convenient since it is possible to withdraw aliquots and treat them in the usual manner. The problem appears very simple.

However several difficulties inherent in the system are immediately apparent. It is not possible to pipette the acid solutions 30 quantitatively due to the great viscosity of the solution. This pro• blem is very easily overcome, since the volumes (Vo, V-t, V^ ) need not be known in absolute terms but only in relative terms. This follows from the expression (Vo-Vt)/(Vt-2/5 Vo). The only necessary condition is that the volumes withdrawn are the same every time.

Since permanganate decomposes in the strong acid solutions, it was necessary to run a blank, containing only permanganate, con• currently with every kinetic reaction. The correction which was applied to all oxidation rates was made simply by adding the volume differences (due to the blank) onto the volumes (Vt) of the aliquots.

In this way good second order kinetics were obtained up to 50$ reaction. It is known that the decomposition of permanganate in the weakly acidic solutions is accelerated by the presence of precipitated IV Mn species. What might be the effect on the decomposition of IV permanganate in the strongly acidic solutions by Mn which stays in solution could not be estimated in the present investigation. It is only assumed that the decomposition of permanganate is small. This is in part confirmed by the constancy of the isotope effect which was observed in the strongly acidic solutions. More work to shed light on this point is now being carried out in these laboratories. Another troublesome feature which complicated the matter further was the heat of mixing of solutions. In order to minimize any errors arising from sudden heating up of solution, the following procedure was rigorously followed. Water was introduced first into the reaction vessel followed by the acid. The solution was subsequently cooled in the ice for a 31 few seconds and then permanganate was added and the mixture cooled again momentarily. The substrate was added last and the Volume of the solution was adjusted quickly to mark with sulfuric acid of appropriate concentration. Since the substrate did not mix immediately with the remainder of the solution, the reaction could be conveniently

started by mixing the reactants and starting the timer at the same

time. Since the temperature of the reaction mixture was slightly below the temperature of the run before mixing, the heat evolved during the mixing brought the temperature to the required value. The

solution was then thermostated during the entire run. It was found

that with the above procedure the initial heat fluctuation never

exceeded one degree.

(e) Spectrophotometric Method in Concentrated Acid

There is a good indication that Mn*^, when present in the

spectroscopic concentrations, does not precipitate but forms a homogeneous

solution in concentrated sulfuric acid solutions. This allows the use

of a spectral method in this region. The disappearance of the Mn04~ was followed at a wavelength of 525 mu. Since the absorption at this wavelength does not drop completely to zero due to the absorbtion by

Mn-1-V, a small correction was made to the optical density readings.

The absorbtion due to Mn*v at 525 mu was determined from fast reactions

which were followed for more than twenty half lives until there was no

observable change. This was done at 72.3, 69.2, 62.5 and 57-5$ sulfuric

acid and the absorbtion at to© was always found to be the same. In 32 the actual procedure 0.111 ml. of 0.0333 M permanganate solution was pipetted into a 10.0 ml. volumetric flask, this was momentarily dipped into a dry-ice acetone mixture and the sulfuric acid of appropriate concentration was filled to the mark and the solution was allowed to equilibrate in a thermostated bath. Similarly 0.1 ml. of 0.5 M formic acid was introduced into another 10.0 ml. volumetric flask to which the sulfuric acid was added in the manner above and the solution was thermostated. The time for the entire procedure never exceeded four minutes. Then 2.25 ml. of the above permanganate solution was withdrawn with a glass syringe and delivered into a cell provided with a glass stopper. The formic acid solution (0.25 ml.) was likewise withdrawn with a glass syringe and introduced into the cell contain• ing permanganate and the timer started. Optical density readings were taken at suitable intervals. The rate was obtained by plotting

m 1/ [Mn ]t against time. The faster reactions could be followed only past their half lives and the rates could only be estimated. For the slower reactions, good second order rate plots were obtained up to 60$ reaction.

3. The pKa Determination of Fluoral Hydrate

For reasons which will be discussed later it was necessary to determine the pKa of fluoral hydrate, and its deuterio analogue.

The method was similar to that used previously by Stewart and Walker

(73).

Fluoral hydrate (59.k mgm) was weighed in a closed container 33 and washed into a 50 ml. volumetric flask with doubly distilled water which was previously boiled out to expel carbon dioxide. A kO ml. aliquot was pipetted into a 125 ml. four necked flask and a slow stream of nitrogen was passed through. A thermometer was inserted into the flask and the whole assembly thermostated to 25 t 0.2°. The pH of the solution was measured with the Beckman E-2 blue glass electrode which was standardized before and after the experiment against pH 7 and 10 Beckman buffers. The measurements were made on the Beckman model G pH meter. A measured quantity of 0.010 N sodium hydroxide (carbonate free) was added periodically to the solution of fluoral hydrate which was stirred with a teflon covered magnetic bar, and the pH measured. The stirrer was shut off during the measurements.

The burette containing sodium hydroxide was equipped with a drying tube filled with solid sodium hydroxide to prevent contamination by carbon dioxide. In all measurements, the pH readings were accurate to - 0.01 pH units. The sodium hydroxide was standardized with 0.01 N sulfuric acid, the normality of which was known accurately.

Calculation of the pKa,.

The pH was plotted against the volume of the base added and since no typical titration curve was obtained, the equivalence point had to be calculated from the weight of the fluoral hydrate used. The pKa values were obtained from the following equation -.lilh):

+ + H+ = U^ciAl -[H ] + [or] )/([A-] + [H ] - [OH-] ) which may be converted to the following solutions above pH 7, 3*

pKa = pH + log M + [0H~1 [A-\ - [or] where [HA]O = initial concentration of fluoral hydrate and [HA] and

[A~] stand for the concentration of fluoral hydrate and its anion

respectively, present at any time, calculated from fHA]o and the amount

of base added. The pH values corresponding to 40, 50 and 60$

titration were used to calculate the pKg,.

Sample Calculation of PKQ,.

At the half equivalence point, the solution consists of

40.00 ml. of 0.0111M fluoral hydrate and 22.20 ml. of 0.010 N sodium

hydroxide. The total volume of the solution is thus 62.2 ml. (u« 0.02)

[HA] = 40 x 0.0111 = 0.00357M = [A-] 2 x 62.2

pH = 10.18 and [0H~] = O.OOOI513M

Then pKa = pH + log [HA] + [0H"J [HA] - t0H-J = 10.18 + log 0.00372 0.00342

= 10.216

= 10.22

The same method was used for the determination of the pKa of deuterio

and protio analogues at ionic strength of 1.0.

4. Determination of Heat of Ionisation of Fluoral Hydrate

Only an approximate method was used here. The procedure was

essentially the same as described above. To 40 ml. of 0.0109 M fluoral 35 hydrate were added 20.05 ml. of 0.010 M sodium hydroxide. This corresponds to about 46$ titration. The pH of the solution was measured at various temperatures and the ionisation constant was computed in the usual manner. The pKa's at various temperatures were plotted against the corresponding l/T values and the heat of ionisation was calculated from Van't Hoff's equation (Ref. 74 p. 829 f.f.). d InK = AH dT RT2

/d InK = AE f 1 dT J ~R~/ T§"

log K = - AE 1 + C 2.5 R T

or log (Ka)g = - A H [ 1 - 1 "

(Ka)i 2.3 R [ T2 Ti

Plot of pKa versus l/T gave straight line whose slope is:

slope = AH 2.3 R

5. Determination of pKa of Permanganic Acid

Spectral Measurements.

The procedure, was essentially the same as described else• where (75). A 0.10 ml. aliquot of O.O321 M permanganate solution

(in water) was pipetted into a 10 ml. volumetric flask. This was momentarily dipped into dry ice and then quickly filled up to the mark with the appropriate concentration of sulfuric acid. The volumetric flask was vigorously swirled during addition of sulfuric acid so that any heat of solution was quickly dissipated. The 36 solution was immediately transferred into a glass stoppered cell and the optical density was measured as a function of time at two wavelengths using the Beckman DU spectrophotometer. Extrapolation to zero time was made for all the measurements. Since the introduction of the 0.1 ml. aliquot resulted in dilution of the sulfuric acid, a suitable correction was made. The cell compartment was kept at 25° by means of thermospacers. A blank containing sulfuric acid of same concentration as the test solution was run alongside to compensate for any absorption by the solvent.

Calculations of pKa,.

The pKa of an acid HA is defined (76) as

PKJJA = log CHA + log fHA

+ cA-cH+ fA- %

Since the measurements are carried out in strong sulfuric acid, where the tendency of the solution to transfer a proton to the charged base is expressed by the H- function, the pKg, relation becomes:

PK-HA = H_ + log CgA where C signifies concentration. The experimental data were first treated by a method previously described by Symons (31)- The ratio

CHA/CA" is given by:

WCA- = (AcA- - A€ ) / (Afe - AeM) where _ represents the extinction coefficient of the permanganate at wavelength X , is "the extinction coefficient of its conjugate 37 acid at the same wavelength and X^ is the extinction coefficient of

solution containing both species.

A modified Davis-Geissman procedure was also used to check

the validity of the results (75). Denoting = CD \m - OD \ A_

in a solution in which permanganate exist completely as the anion and

^\HA = OD^ . - OD^ in a solution where the permanganate ion is

effectively completely protonated. In the acid concentration which contains both permanganate and permanganic acid, the quantity A may

be defined as A = 0D\„. - 0D\ A_. The ratio is given by MA A

CJA = Am - A CA- A- ^A- and a plot of log of this ratio versus H- should give a straight line whose intercept equals the pKa. The advantage of this method is that

the need for an arbitrary choice of the inflection point of an ionisation

curve is eliminated, since the best straight line may be found by the method of least squares. The wavelength for the permanganate ion

absorption was taken at 525 nyu and that for the permanganic acid at

U58 mu. The results obtained therein will be discussed further in

the text.

6. Preparation of Benzhydrols

4,V-Dimethoxy, 4,4'-dimethyl, 4,4'-dichloro and 4,4'-bis -

(dimethylamino) benzhydrols and their deuterio analogues were

synthesised by the reduction of the corresponding ketones with LiAlH* 38 or LiAlD4 respectively, by a standard procedure. The yields of the crude products were of the order of 70 to 80$.

(a) 4,4'-Dinitrobenzhydrol

Nitration of diphenylmethane yielded 4,4'-dinitrodiphenylmethane

(77)5 this compound was oxidized with chromium trioxide to

4,4'-dinitrodiphenylketone (78).

4,4'-dinitrodiphenylketone (2.7 g) was dissolved in 30 ml. of anhydrous tetrahydrofuran and cooled to -78°. A solution of LiAlILi

(0.4 g) in tetrahydrofuran (25 ml.) was added to it over a period of

15 minutes with vigorous stirring. Formation of a red colour appeared at the start of the reaction, this changed to brown after one half hour. Stirring was continued for two hours at -78°, and the mixture was then allowed to warm up to 0° and poured over ice-sulfuric acid mixture. Sulfuric acid was neutralized with sodium bicarbonate and the mixture extracted several times with 50 ml. portions of ether.

The ethereal extracts were washed with 1 N sulfuric acid, sodium bicarbonate and water, dried over magnesium sulfate, evaporated to dryness and the solid dissolved in tetrahydrofuran - benzene - ligroin

(30-60°) - 1:1:1 mixture and recrystallized in the refrigerator.

Yield : 25$.

The deuterio analogue was prepared by the procedure described above with the exception that HAID4 was used in place of LiAlRj..

Yield : 25$.

(b) Physical Constants and IR Data for Benzhydrols

The infra-red spectra were obtained in order—to establish the 39 presence of deuterium in the isotopically substituted molecules. The

samples were made in potassium bromide discs with the exceptions of

4,4'-dimethoxybenzhydrol which was dissolved in carbon tetrachloride

and 4,4'-dinitrobenzhydrol run as a solid between two sodium chloride

plates. The C-H and C-D stretching frequencies of dinitrobenzhydrols

were very weak (79)•

C-H or C-D C-H or C-D Stretching Bending m.p. Frequency Frequency m.p. °C Benzhydrols cm-1 cm"1 °C lit. and ref.

4,4'-OMe 1364 69-70 69-70 (80)

4,4' -die- o< -d 2140 1004 69-69.5

4,4'-Me 1357 69-70.5 69-70 (81)

4,4'-Me- c/ -d 2150 1002 69-70

4,4'-Cl 2880 1350 94-95 91.5 (82)

4,4'-Cl-c( -d 2140 1000 94.5- 95.5 -IO3 4,4'-NMe2 1400? IOI.5 102-103 (83)

4,4'-NMe2- 0( -d 2120 995 IOI.5-IO3

(84) 4,4'-N02 2900 167.5-168.5 164-165.5

4,4'-N02- 0( -d 2300? 168-169

(c) Fluoral Hydrate

The synthesis of this compound was carried out essentially

by the procedure developed by Braid, Iserson and Lawlor (85).

Trifluoracetic acid was reduced with stochiometric amount of

LIAIH4 and the crude aldehydrol dehydrated in a mixture of concentrated

sulfuric acid-phosphorus pentoxide. The fluoral hydrate was obtained ko by adding less than the stochiometric amount of water to the tri- fluoroacetaldehyde which was kept at -78°. Yield: 43$, m.p. 93-94°

(sealed tube), lit.: m.p. 69° (85).

Analysis*: Required for CF3CH(OH)2: C 20.68, H 2.59

obtained: C 22.51, H 192

Molecular Weight: 223, Theoretical 116.

Infra-red: C-H stretching at 2980 cm-1.

Since the molecular weight appears to be about twice the theoretical, it appears that the compound is a condensation dimer con• taining two molecules of fluoral hydrate less one molecule of water.

Required for the dimer: C 22.40, H 1.88, mol. wt.: 214. As will be evident from the n.m.r. spectra (Appendix), the compound was identified as the dimer.

The deuterio analogue was synthesised by the above method

using LiAlD4. Yield: 77$, m.p. 92-93° (sealed tube).

Analysis: Required for CF3CD(0H(2: C 22.50, H 2.80

found: C 21.86, H 1.95

Infra-red: C-D stretching at 2220 cm_i.

The proton n.m.r. spectrum shows only one peak due to the hydroxyl protons (Appendix). The physical constants and analysis indicate that the deuterio compound is also a condensation polymer, probably the dimer.

* Analysis by A. Bernhardt, Muelheim, West Germany. hi

(d) 1,1,1;5»313-Hexafluoro-2-propanol-g-d

An attempt was made to oxidise the hexafluoro-2-propanol with acidic permanganate and then reduce the hexafluoroacetone with LIAID4.

It was assumed that the following reaction will take place.

3 CF3CHOHCF3 + 2 HMnC-4 > 3 CF3COCF3 + 2 Mn02 + k H20

Potassium permanganate (2.38 g) was dissolved in 12 ml. of 96$ sulfuric acid and the mixture heated with stirring. When the tempera• ture reached about 60° the mixture decomposed violently giving off black soot.

In a further attempt potassium permanganate (3 g) was dissolved in 15 ml. 54$ sulfuric acid, and the mixture heated to 4-5°.

Hexafluoro-2-propanol (3.8 g) was added slowly and the mixture heated at 75° for one hour. Brown colour was evident after 15 minutes of heating. Phosphorus pentoxide (5 g) was added and the mixture heated for further one hour, however no hexafluroroacetone was isolated. It appears that the permanganate decomposed before it had a chance to oxidise.

In order to obtain at least some product from the limited quantity of starting material available, sodium dichromate (5 g), dissolved previously in 5 ml. of water and 10 ml. concentrated sulfuric acid, was added to the reaction mixture and heated at 90° for two hours (86). Only 0.1 ml. of the hexafluoroacetone was collected in a cooled trap. The reaction mixture was cooled and 30$ hydrogen perioxide was added until a clear green solution resulted containing only Ma** and Cr***. The reaction mixture was extracted k2 three times with 50 ml. portions of ether, the ethereal extracts dried over magnesium sulphate and the ether distilled off. The fraction boiling over 60° was kept. A mixture of 1.5 g phosphorus pentoxide and 5 ml. of concentrated sulfuric acid was heated to 100° and to it added dropwise the fraction boiling over 60°. About one ml. of the hexaf luoroacetone was isolated and stored in a trap at -78°.

The LiAlD4 (0.42 g) was dissolved in 10 ml. of ether, cooled to -78° and the perfluoracetone was slowly bubbled into it with stirring. The mixture was stirred for one half hour and then brought to room tempera• ture during further half hour. The complex was decomposed by addition of dilute sulfuric acid. The ether layer was separated and the re• maining solution extracted several times with ether. The combined ethereal extracts were dried and the ether distilled off, The residue was chromatographed first on a 10 foot apiezon J column and finally

on a 10 foot didecylphthalate column. The 1:1 mixture of CF3CH0HCF3 and the product obtained thereby gave a single sharp peak with the

expected intensity. Yield, about 0.3 g CF3CD0HCF3. The infra-red spectrum shows the C-D stretching at 2170 cm-1 and the absence of C-H vibration. The IR spectra for the protio and deuterio compounds were almost identical with the exception of the C-H and C-D vibrations and the fingerprint region, proving the identity of the synthesised compound. The proton n.m.r. spectrum showed only one peak due to the hydroxyl protons. RESULTS

1. Permanganate Oxidation of Fluoral Hydrate

(a) The Nature of the Substrate

It is well established that trichloroacetaldehyde exists in water solutions in the hydrated form. Similarly trifluoroacetaldehyde takes up water very easily to form the hydrate. There was some doubt at the outset of this investigation as to the nature of the species being oxidised, however it was established clearly by n.m.r. measure• ments that the species which exists in water solutions is in fact fluoral hydrate in the monomeric form (for details see Appendix).

(b) Alkaline Hydrolysis of Fluoral Hydrate

Before any extensive oxidation experiments could be carried out, it was necessary to establish the extent of the alkaline hydrolysis of fluoral hydrate. Since chloral hydrate hydrolyses very easily in the alkaline solutions to chloroform and formic acid (87,88) it was anticipated that a similar reaction might occur with fluoral hydrate, yielding fluoroform and formic acid. Since the rate of oxidation of formic acid is established and is known to be constant in the region pH 6 to 9, it is then possible to estimate the extent of hydrolysis of fluoral hydrate by oxidising any formic acid which might be formed.

To check this, the following experiment was performed. A solution of fluoral hydrate (of identical concentration as that used in the kinetic experiments) in 0.15 M sodium hydroxide at ionic strength

1.0 was allowed to stand at 25° for 1720 seconds. The solution was kk quenched with 5 nil of 1 M sulfuric acid and phosphate buffer was added to adjust the pH to 6.5. A stochiometric amount of permanganate was added and the reaction was followed in the usual manner. Since the rate of oxidation of fluoral hydrate at pH 6.6 was already estab• lished by an independent experiment, and found to be slow, it was possible to estimate the extent of the hydrolysis. It was found that at most only 2$ of the fluoral hydrate were hydrolysed. Considering that the half life of the fluoral hydrate oxidation reaction at 0.1 M sodium hydroxide (ionic strength 1.0) is about 50 seconds, then 1720 seconds represent essentially completion of the reaction. Thus to all intents and purposes the extent of the alkaline hydrolysis of fluoral hydrate may be considered negligible for at least 75$ reaction for both, the protio and deuterio compounds,

(c) Stochiometry of the Reaction

In the alkaline solutions the overall reduction of permanganate is accompanied by one electron change. Since the fluoral hydrate is a two electron reductant, it follows from this that the stochiometry is given by:

= 2 Mn04" + CF3CH(0H)2 + 20H" 2Mn04 + CF3C02H + 2H20

This equation is applicable only if the manganate formed reacts with the substrate at much slower rate than permanganate. This was actually the case in the present investigation.

Below approximately pH 12 the manganate disproportionates into permanganate and manganese dioxide and the stochiometry becomes:

+ 2Mn04" + 3CF3CH(0H)2 + 2H 2Mn02 + 3CF3C02H + UHgO 45

In the strong acid permanganate is reduced to MnIV and thus

the overall stochlometry is the same as for the weakly acidic region.

(d) Products Analysis

It is now well established that permanganate and dichromate oxidations of fluoral hydrate yield trifluoroacetic acid which is stable

to further oxidation (85,89,90,91).

(e) Order of the Reaction

All of the kinetic results could be fitted onto the second order rate plots for at least 75$ reaction. The increase in the con•

centration of both reactants did not have any effect on the rate

constant. From this as well as other evidence accumulated about the permanganate reaction (16,17,19,58,59,60,61,64,65,66,67) it is concluded

that the reaction is of second order and in particular first order in

each of the reactants.

2. Permanganate Oxidation of Fluoral Hydrate in the Region pH 6 to 11

Constant ionic strength conditions (u=1.0) were used for all

the runs throughout the pH region. At constant pH and constant ionic

strength the rate of oxidation of fluoral hydrate is given by the

expression

- d fMnQ4"] = k2 [Mn04 ] [ Fluoral Hydrate] .

dt .

Good second order rate plots have been obtained for all the runs in

this region. Typical rate plots for the protio and deuterio compounds

are shown in Figure 1. time sec * x 47

A marked dependence of the rate on pH is observed, with the rate rising with increasing pH. After reaching a certain value of the pH the rate appears to level off. Purely by inspection it was recognized that the relation between the rate constant and £0H~] or

[H+] is not linear. However a plot of the rate constant versus pH shows a behaviour which is characteristic of an ionidation curve /

(Figure 2). To explain this behaviour two obvious possibilities are immediately apparent. Either the permanganate or the fluoral hydrate are undergoing some ionisation process. Since it is known that the permanganate species does not undergo any change in this region, it follows that it must be the ionisation of fluoral hydrate. To confirm this idea, following experiments were performed,

(a) Determination of pKp. of Fluoral Hydrate*

The method of measurements have been described earlier. On the basis of pKa determinations of other trifluoromethyl substituted alcohols (66) it was anticipated that the first ionisation of the fluoral hydrate will occur in the pH region below pH 12. This assumption proved to be correct. The pKa1s were first determined at low ionic strength (y. » 0.01 and 0.05). Since all the kinetic experi• ments were carried out at ja.=1.0, it was also necessary to determine the pKa at this ionic strength. The ionic strength was adjusted in this case with potassium sulphate. The results are presented in

Table I. * It should be pointed out that only the first ionisation constant of . fluoral hydrate could be determined directly. The second ionisation constant could not be determined experimentally.by present methods. Oxidation of Fluoral Hydrate 49

Table I

r- 0.01 0.0111 6.17 10.21

0.01 0.0108 6.03 10.22

-0.05 0.0544 6.46 10.19

1.00 0.0109 9-78 10.01

1.00 0.0109 9-33 10.03

The error in the pKg determinations was estimated as - 0.02 pK units.

It appears that the pKg. of the deuterio compound is slightly higher

than that for the protio compound. Even though this is within the

experimental error, the higher value for the deuterio compound would be predicted on the basis of Halevi's considerations (92). The

enhanced ionization implied by a higher value of the ionisation

constant at high ionic strength is again in accord with Debye-Hueckel prediction of the effect of ionic strength. Since the pH measurements have doubtful value at the ionic strength ju=1.0, it must be assumed

in this connexion that the quantity measured on the pH meter is

related to the pH. That this assumption is reasonable will become

apparent from the kinetic results,

(b) Enthalpy of Ionisation of Fluoral Hydrate

Since it was anticipated that ionisation equilibrium of

fluoral hydrate occurs before the oxidation reaction, any determina- 50 tion of energy and entropy of activation for the oxidation process would automatically involve the enthalpy of ionisation. For this reason it was necessary to determine the variation with temperature of the pKa. The results are summarised in Table II and the plot of pKa's versus l/T is shown in Figure 3.

Table II

Variation of pKa of Fluoral Hydrate with Temperature

Temperature °C pK^

15.7 10.10

25.2 9-97

34.3 9.82

43.6 9.68

52.5 9.57

The value of enthalpy of ionisation which was obtained from the slope in Figure 3 was found to be: A Hi = 6.2 - 0.3 kcal mole"1. This value must be regarded as approximate only since the method used in its determination is not considered to yield accurate results. The somewhat high heat of ionisation of fluoral hydrate, compared to other aliphatic acids (93,94), is in accord with findings of other investiga• tors (95,96,97,98). The enthalpy of ionisation was not determined for the deuterio compound but it was assumed that its pKa changes with the temperature in the same manner as for the protio analogue.

52

(c) Variation of pH of Buffer with Temperature

Phosphate buffer was used to keep the pH constant during the oxidation reactions. The heat and entropy of activation for the permanganate oxidation of fluoral hydrate were also determined in the buffer and this necessitated the determination of the variation of pH

of buffer with temperature. Since the oxidation reactions were fast

it was found impractical to measure the pH of the buffer during the

runs. For this reason the kinetic solution consisting of buffer

(u=1.0) and fluoral hydrate was made up and water was substituted for

permanganate and the temperature variation of the pH of this solution was measured. The results are summarised in Table III.

Table III

pH Variation of the Phosphate Buffer with Temperature at JU=1.0

Temperature °C. pH of Buffer

15 10.53 20 10.28 25 10.20 30 10.12 35 10.06

kO 10.00

(d) Treatment of Kinetic Results

The kinetic data corresponding to Figure 2 are shown in

Table IV. It has been pointed out already that the plot of the rate

constant against pH appears to follow a pattern characteristic of an 53 ionisation behaviour. Whether this is indeed the case, may be checked in several ways. The pl^ of fluoral hydrate is known at u=1.0 and thus the concentration of the fluoral hydrate anion may be easily calculated at any pH from the expression:

pKg, = pH + log [A"] / [HA] where [A-] is the concentration of the anion and [ HA] represents the concentration of the unionised fluoral hydrate. The plot of the rate constant versus the corresponding fluoral hydrate anion concentra• tion should give a straight line. That a good agreement has been obtained may be easily verified from Figure k. The maximum value of the rate constant which is attained at the full ionisation of fluoral

hydrate may be read from the graph (k2 max = l..?9). This quantity is not accessible experimentally for reasons discussed later. The plot of the rate constant versus the anion concentration becomes rather insensitive at small concentrations of the anion. Better results

should be obtained by relating the log k2 to the pH, but here the region of the high anion concentration can not be correlated. Recognising

the fact that the quantity of log k2/(k2 max - k2) is identical with log o( /(l - (X ), its plot versus pH should give a straight line whose intercept with zero must give the pKa of fluoral hydrate. Moreover, if the relation is real, a slope of unity must be obtained, as it necessarily follows from the equation

pKa = pH + log ^ .

1 - ex-

Agreement is indeed excellent as can be seen from Figure 5. The value Table IV

Oxidation of Fluoral Hydrate in the Region pH 6 to 12

Variation of the Rate Constant with pH

1 PH ks 1.mole" .sec

6.12 O.OOOi69a

6.60 0.000490

7.62 0.00352b

7.62 0.00339^

8.6l 0.0391

9.34 0.202

9.40 0.199

9-72 0.420

9.82 0.512

10.02 0.662

10.20 O.738

10.20 O.756

10.31 0.891

10.90 I.O98

11.29 I.O38

11.70 0.987

Data for the prOtio compound.

All runs at 25.00 ± 0.03° and u = 1.0 [Fluoral hydrate] 6 = 0,0015 M

[Mn04"] o = 0.001 M

(a) [Fluoral hydrate] 0 =_ 0.0612M, [Mn04" ]0 = 0.0408 M

(b) [Fluoral hydrate] 0 = 0.015M, [Mn04" ]0 = 0.01 M

FIG. 5 Oxidation of Fluoral Hydrate Linear relation between pH

OJ k2 1 and log £—: 777- CM X ffl E CM U) / o + 1 - / o

o-

-1 - / -2 - / slope = 1*0

-3-

-4- 6 7 8 9 10 11 57 of the pKa of fluoral hydrate estimated from the kinetic data is 10.04 which is in good agreement with the value 10.01 derived from independent pH measurements. The slope is unity which confirms that the correlation between the rate of oxidation and ionisation of fluoral hydrate is real. It is also obvious from the graph that the ionisation process is still responsible for the rate of oxidation even four pH units away (pH 6) from the pKg, of fluoral hydrate. This is a somewhat surprising result in view of the fact that in this region the fluoral hydrate is almost completely unionised,

(e) Deuterium Isotope Effect

The deuterium isotope effect can be used to indicate whether the C-H bond is broken in the rate determining step. For this reason the fluoral hydrate-l-d was oxidised by permanganate under identical conditions as for the protio compound. Comparison of the rates for the protio and deuterio compounds are presented in Table V.

The presence of a substantial deuterium isotope effect shows clearly that the transition state involves the rupture of the C-H bond. With the exception of the value at pH 6.12 it almost appears that the isotope effect decreases as the rate becomes faster. It cannot be said at present whether this phenomenon is real or if it arises from intrinsic experimental difficulties such as those involved in the accurate determinations of the very slow rates. Wiberg and Slaugh

(99) have predicted a general decrease in the magnitude of the isotope effect for faster reactions, however this statement refers to reactions of compounds with different substituents and not to reactions of the same compound. 58

Table V

Oxidation of Fluoral Hydrate, Protio and Deuterio Analogues,

in the Region pH 6 to 12

Deuterium Isotope Effect

ks>H - 1 pH 1.mole-1.sec-1 l.mole .sec" ^HAD

6.12 I.69 x IO"4 1.50 x 10"5 11.3

7.6l 3.52 x IO'3 2.66 x IO"4 13.7-) •14.4 3.39 3 2.26 10"4 7.6l x IO" x 15.0* 10.20 0.738 0.0743 9.9-j f 10.2 10.20 0.756 0.0735 10.3J

10.70 1.06a 0.116 9.1

All runs at 25.O - O.O30 and p. = 1.0

Concentration of fluoral hydrates and Mn04~ same as given in Table IV.

(a) Estimated from Figure 4.

(f) Activation Parameters for the Oxidation of Fluoral Hydrate in Region pH 6 to 12

The temperature study of the oxidation reaction has been made

in the region where the ionisation of fluoral hydrate was only partial,

and this necessitated a knowledge of variation of the pKa and pH of

the solution with temperature. These have been determined as described

earlier. Thus the values of the energy and entropy of activation

reported herein are the true values from which the pre-equilibrium

step has been separated. The rate constants used in this evaluation 59 of the thermodynamic data are the true rate constants which were derived from the expression:

True ka = kexp *y

Ka [OH-j .

where keXp is the experimental rate constant and is the ionisation constant of fluoral hydrate. It was assumed here that the oxidation reaction occurs between the fluoral hydrate anion and permanganate ion. The evidence for this assumption and a full discussion will be presented later. The activation parameters for the protio and deuterio fluoral hydrates are presented in Table VI and the corresponding plots are shown in Figure 6. The energy and entropy of activation for fluoral hydrate obtained in this investigation are in reasonable agree• ment with those for the phenyltrifluoromethyl carbinols (64,65), forma•

te ion (59,60,61,63) and cyanide ion (67). The entropy value is in

accord with a reaction involving two species of the same charge.

3. Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions

(a) Rate Dependence on Hydroxyl Ion Concentration

Several oxidations of fluoral hydrate by permanganate were

carried out in alkaline solutions of pH 12 and higher. The manganate which is the product of permanganate reduction is stable in this region and reacts with fluoral hydrate only slowly. For this reason a ratio

of fluoral hydrate : permanganate =1:2 was used in the alkaline

region. 60

Table VI

Thermodynamic Data for the Oxidation of Fluoral Hydrate in

the Region pH 6 to 12

CFaCH(0H)g CF^CD(0H)p

Temp. k2lH True k^n _x k2jp _x True

°C. pH l.mole" .sec~ l.mole" .sec" l.mole" .sec" l.mole" .sec"

15.0 IO.33 O.332 0.204

15.0 IO.33 O.328 0.202 16.4 IO.32 0.0345 0.0218

20.0 10.28 0.514 O.3O3 0.0446 0.0299

25.0 10.20 0.738 0.434 0.0806 0.0496

25.0 10.20 0.756 0.444 0.0743 0.0458

25.0 10.20 0.0755 0.0453

30.0 10.12 1.09 0.626 0.109 0.0655

30.0 10.12 0.102 0.0615

35-0 10.06 I.56 0.856 0.179 0.101

40.0. 10.00 2.3O 1.26 0.264 0.152

4o.o 10.00 2.35 1.29 0.254 0.146

Activation Parameters

A H* = 12.4 kcal mole-1

| CF3CH(0H)2 A S* = -18.4 e.u.

A H* = l4.3 kcal mole"1

| CF3CD(0H)2 A S* = -l6.7 e.u. 61

1000/T°K+x 62

It was found that the oxidation rate was markedly dependent on the hydroxyl ion concentration. Rise in rate was observed with increasing base concentration; closer examination reveals that the plot of the rate constants against the hydroxyl ion concentration is linear for both protio and deuterio compounds (Figure 7); the corres• ponding data are shown in Table VII.

Since the rate of oxidation due to the first ionisation of fluoral hydrate is essentially independent of pH in the region above pH 12, the observed increase in the rate may be accounted for by assuming the participation of the dianion of fluoral hydrate in the oxidation reaction. This assumption is substantiated by the fact that at zero base concentration (Figure 7) the rate constant of the protio compound is that which has been estimated for the first ionisation from Figure k. If Pauling's correlation is accepted (100) then a pKa of about 15 would be predicted for the second, ionisation of fluoral hydrate. Assuming that this value is correct in the order of magnitude, it can be seen that even in 1.0 M sodium hydroxide, the extent of ionisation is only slight. Thus it should be possible to correlate logarithm of the rate constant with pH. This plot should give a straight line; moreover the slope must be unity If the relation is real. Also the plot should be well separated from the one for the first ionisation. That all these conditions are satisfied may be seen from Figure l6, presented later in the. text. There can now be little doubt that the dianion of fluoral hydrate is the reactive species in the oxidation reaction in strongly alkaline solutions and the mono-anion

Table VII

Oxidation of Fluoral Hydrate in Strongly Alkaline Region

Rate Dependence on Hydroxyl Ion Concentration

[OH-J k|H x k2D 1 M l.mole .sec" l.mole .sec kH/kp

0.009 k.ok

0.019 5.42

0.039 8.08

0.059 10.93 2.42 4.9

O.O69 13.72 2.47 5.5

0.079 14.80 2.83 5.2

0.099 17.80 3.60 h.9 0.099 18.50

0.119 21.9 4.46 k.9

0.119 23.O 4.40 5.2

O.I58 5.7^

0.158 5.76

1.00 260 - 100

All runs at 25.00 ± O.O30 and u = 1.0

[Mn04"] 0 == 2 [Fluoral Hydrate] 0 = 0.002 M 65 of fluoral hydrate in the active species in the weakly alkaline region.

(b) Deuterium Isotope Effect

It may be noted from Figure 7 and Table VII that the isotope effect kji/kj) = 5.3 which is about half of that observed for the first ionisation. The decrease in the magnitude of the isotope effect, as compared with that for the oxidation reaction due to the first ionisa• tion of fluoral hydrate, may be partly ascribed to a faster reaction; this is essentially in accord with Wiberg and Slaugh's (99) findings.

(c) Influence of Ionic Strength

Several oxidation runs in strongly alkaline solutions were carried out with varying ionic strength. The minimum ionic strength which can be used is limited by the nature of the product, trifluoroacetic acid, which is fully ionised at that point. For this reason only rough estimate of the effect of ionic strength may be obtained. The results are summarised in Table VIII.

Table VIII

Oxidation of Fluoral Hydrate in Strongly Alkaline Solutions Influence of Ionic Strength 1 5.26 x

18.50 1.0 1.0 0.235 17.80 1.0 1.0 0.235

:-7-96 0.2 0.44-7 0.182 5.81 0.1 0.516 0.156

Temp. = 25.00 + 0.03° [oH"j = O.O99 M K2SO4 used as the inert salt. 66

The ionic strength is far outside the region in which the Debye-Huckel

theory should hold and therefore no correlation could be expected.

The correlation with Harty-Rollefson (101) relation is perhaps better,

and a slope of 2 is obtained for the two slow rates. Even though the

agreement is not striking with either theory, the effect of changing

ionic strength is in the direction to be expected for a reaction between two anions.

(d) Oxidation by Manganate

Even though manganate accululates during the permanganate

oxidation in strongly alkaline solutions, it was advisable to

determine the magnitude of the rate constant for the manganate

reaction. The rate was followed by the iodometric-method in essentially

the same manner used for the permanganate oxidations. Linear second

order rate plots were obtained in all but one case where sigmoid type

curve was observed. The rate constants were computed from the

expression: ± x Vo - Vt

k2 = [Fluoral Hydrate] 0 x t Vt - 1/2 Vo

The kinetic data are shown in Table IX. The results show that

oxidation of fluoral hydrate by manganate occurs 6k times slower than

the permanganate oxidation at the same alkalinity. Also the deuterium

isotope effect is the same as for permanganate oxidations. Increase

in alkalinity is followed by increase in the rate from which it may

be concluded that oxidation is occurring by manganate and not by

permanganate which might arise by disproportionation. 67 Table IX

Oxidation of Fluoral Hydrate in Strongly Alkaline Region

Oxidation by Manganate

l.mole-1.sec"1 l.mole"1.sec"1 LQH'IM ksH kgp kH/kp

0.1 0.285 O.O596 4.8

0.5 O.3OO sigmoid curve

[Fluoral Hydrate] o = •'"'0.001 M

= LMn04 Jo = 0.001 M

Temp. = 25.0 + O.O30, ji = 1.0

(e) Activation Parameters for Oxidation of Fluoral Hydrate in Strongly Alkaline Region

Unfortunately it was not possible to determine the true energy and entropy of activation, since the second ionisation constant of fluoral' hydrate is not known and its evaluation by ordinary means is almost impossible. The only effect which could be determined was

the temperature variation of k2 max for the oxidation reaction due to the first ionisation of fluoral hydrate. This quantity which was pre• viously obtained from Figure 4 was simply superimposed on the

log k2/T vs l/T plot in Figure 6 and the values of k2 max read from the graph for the corresponding temperatures. That this procedure is

correct follows from the subsequent agreement. The value of k2 max is obtained at full ionisation of fluoral hydrate. Thus the ionisation process can not have any effect on its value. Moreover, the tempera- 68 ture effect upon the ionisation has been already separated out in

Figure 6 and therefore the temperature variation of k2 max must necessarily follow the same AH* and AS^ as in Figure 6. Thus the activation parameters for the oxidation in the strongly alkaline region still include the temperature effect upon the second ionisa• tion of fluoral hydrate. This effect is impossible to estimate with present data. The results are presented in Table X and XI and the

plot of log (k2 - k2 max)/T versus l/T is shown in Figure 8.

Table X

Oxidation of Fluoral Hydrate in Strongly Alkaline Solution

Activation Parameters for the Frotio Compound

k2 max k2 1st Ionisation 2nd Ionisation -1 -1 Temp. °C l.mole .sec l.mole .sec 13.0 0.547 7.65 16.0 0.699 8.21 20.0 0.962 10.75 25.0 1.29 13.27

30.0 1.99 16.87 30.0 1.99 16.09 35-0 2.72 20.80 35-0 2.72 19.90

All runs at u = 1.0

[Fluoral Hydrate] 0 = 0.001 M

[Mn04"] o = 0.002 M 69

AH* = 6.6 kcal mole-1

AS* =-31.4 e.u.

Table XI

Oxidation of Fluoral Hydrate in Strongly Alkaline Region

Activation Parameters for the Deuterio Compound

ks max k2 1st Ionisation 2nd Ionisation -1 Temp. °C l.mole .sec" l.mole-1. sec"1 15.1 0.0425 1.50

20.0 O.O656 2.02

25.0 0.126 2.47,

30.0 0.190 3.08

30.0 0.190 3.22

35-0 0.286 4.00

35-0 0.286 4.04

Conditions same as in Table X.

AH* = 7-5 kcal mole"1

As* = -31.8 e.u.

The highly negative value of the entropy of activation is again in accord with a reaction between the permanganate ion and the di-anion of fluoral hydrate. On the other hand the small heat of activation suggests that the electron transfer to permanganate is facilitated by the two negative charges on the fluoral hydrate anion. Considering that the A H^ still includes the heat of second ionisation of fluoral 70

FIG 8 Oxidation of Fluoral Hydrate (Di -anion) Variation of rate constant ^vvith temperature

-1*8 -| 1 1 i 3*2 3'3 3 '4 3*5 1000/T°K 71 hydrate, it would be expected that AH^ would have a different value.

It is however, not possible to estimate what the heat of the second ionisation might be. The evidence presented thus far does not allow one to make a conclusion about the mechanism of permanganate oxidation, in fact it is not possible to distinguish from the results discussed this far between one or two electron transfer or a hydride ion shift to permanganate.

(f) Oxidation of Fluoral Hydrate at 78°C

Anomalous Arrhenius- frequency factors were observed by

Bell and collaborators in the bromination of 2-carbethoxycyclopentanone

(102,103,10U). This reaction is also catalysed by the fluoride ion which exhibits the most marked effect. The unusual frequency factors were related to the quantum mechanical proton tunnelling*. Even though no deviation from the Arrhenius plot is observed between 5 to 50.%

Hullett (103) has detected bending of the plot below -10°, in the direction expected for proton tunnelling. At -20°, the fluoride ion catalysed reaction showed an isotope effect of kH/ku = 10 which approached the theoretically predicted value of 16.

The ratio of Arrhenius frequency factors for the permanganate oxidation of fluoral hydrate monoanion was found to be AD/AH =2.2, and the corresponding difference in the activation energies was

*Proton tunnelling is not necessarily concerned only with protons; in fact protons, hydrogen atoms or hydride ions may be considered under this terminology (105). 72 EJJ - Eg = 1.9 kcal/mole. Both of these values are of similar magnitude as those for the bromination of 2-carbethoxycyclopentanone. If the high isotope effect for the oxidation of fluoral hydrate (kn/kj) = 10) could be taken as an indication that proton tunnelling is occurring at room temperature, then oxidation at high temperature should show a much smaller isotope effect. Fluoral hydrate was oxidised at "JQ" and the results are shown in Table XII.

Table XII

Oxidation of Fluoral Hydrate at 78°

kg l.mole"1.sec"1 kp l.mole"1.sec"1 kg/kp

O.OI65 0.00265 6.3 0.0164

TEMP. = 78°, ju = 1, [Mn04"]0 = 0.0408 M

pH = 6.12 at 25°, not measured at 78°

kH/kp = 11.3 at 25°

An isotope effect of 6.3 was obtained at 780. An identical value for the isotope effect at this temperature may be calculated from the existing data for permanganate oxidation of fluoral hydrate monoanion

(Table VI and Figure 6) using the formula:

S* = 2.303R flog kg - log k + AH*] L T h 2.303RTJ.

However in view of Bell's findings (102) these results can not be taken as conclusive evidence that proton funnelling is absent. The frequency factors for the oxidation of the fluoral hydrate di-anion were found to be Ap/An = 0.9 and the corresponding energies of 73 activation - Eg = 0.9 kcal/mole. Both of these quantities are composite values including any pre-equilibrium step; this however has little effect on the ratio of frequency factors or difference of activation energies. Here it is clearly apparent that the frequency factors are the same within experimental error thus making the possibility of proton tunnelling very unlikely. It is also difficult to see how the mechanism of permanganate oxidation would be radically different for the mono- and di-anions of fluoral hydrate. The normal temperature variation of the isotope effect, 11.3 at 25° to 6.3 at

78% may be explained by assuming that bending as well as stretching modes are lost in the transition state (106,107). These values which are smaller than theoretically predicted ones suggest only partial loss of the bending modes in the transition state. k. Oxidation Kinetics in the Region pH 11 to 12.5

It has been mentioned earlier that the kinetics in this region are difficult to interpret due to the disproportionation of manganate.

It is generally agreed that the manganate is unstable in alkaline solutions less than 1 M in hydroxyl ion (3,11,13). Other investiga• tions show that the disproportionation of manganate occurs below pH

11.7 (17). Surprisingly, no information is available about the rate of manganate disproportionation or the region in which this occurs.

The present investigation does not offer any quantitative results; however it attempts to delineate the region in which manganate disproportionation occurs. It may be noted from Table IV that the Ik rate of oxidation of fluoral hydrate rises up to pH 10.9 and then drops with further increase in pH. Wiberg and Stewart have noted a. similar behaviour for formic acid (58). Exactly opposite behaviour is observed in strongly alkaline solutions as may be seen from the data, for the protio compound in Figure 7- It appears that the rates at 0.009 and 0.019 M hydroxyl ion are higher than predicted by the straight line relationship. The relative increase in rate in the strongly alkaline region and a decrease in the weakly alkaline region may be accounted for reasonably well by the manganate disproportiona• tion reaction. It should be borne in mind that the ratio of substrate to permanganate is 3:2 in the weakly alkaline region and

1:2 in the strongly alkaline region. Supposing that permanganate oxidation is being followed in the strongly alkaline region, say at pH 12. Manganate is one of the reaction products and is reasonably stable there. If however some of the manganate disproportionates into permanganate and Mn02, it follows that the permanganate concentra• tion must increase and in consequence the rate must become faster.

In fact, at slightly lower alkalinity; more of the manganate will . disproportionate and the rate will be correspondingly faster. On the other hand in the weakly alkaline region, say at pH 11, the stochiometric concentration of permanganate is fully dependent on the instantaneous disproportionation of any manganate formed. If now the disporportiona- tion of manganate is a slow process (e.g. rate controlled), then the expected stochiometric concentration of permanganate is not building up fast enough. Therefore a decrease in the rate of oxidation should 75 result. Both of the above phenomena have been observed in this investigation. The above arguments are further supported by the observation that in the strongly alkaline region (substrate: permanganate = 1:2) at pH 12 and below, the formation of green manganate is never observed. In fact the purple permanganate colour changes into grey and then slowly into a red colour and finally Mn02 precipitates and permanganate colour appears again, clearly indicating that the manganate has disproportionated. In the weakly alkaline solutions (ratio = 3:2) the situation was reversed. At pH 11 and higher after most of the permanganate was used up in the oxidation reaction, the green manganate was clearly observable for a short time followed by appearance of a red colour and finally by precipitation of

MnOg and appearance of permanganate colour. Again a clear indication that the manganate disproportionation is rate controlled. On the basis of the arguments and observations it may be concluded that any attempt to measure oxidation kinetics in the region pH 11.0 to 12.5 will be dependent mainly on the rate of manganate disproportionation, relative to the rate of oxidation. Thus in the weakly alkaline regions it might be possible to follow the oxidation kinetics beyond pH 11.0 providing the oxidation rate is much slower than the disproportionation rate. Oxidations in the strongly alkaline regions might be extended below the pH 12.5 if the rate of oxidation is much faster than the rate of disproportionation. However the extension is only slight - one cannot hope to obtain a reasonable value for the rate between pH

11.5 to 12.0, where mixed kinetics are observed. It should be pointed 76 out that all results were obtained at ionic strength 1.0 and that the region of manganate disproportionation might be different at lower ionic strengths. Caution should be exercised when interpreting some of the older results which were obtained in the region pH 11 to 12.5-

5- Oxidation of Fluoral Hydrate in Weakly Acidic Region, pH 6 to 0.7

The oxidation of fluoral hydrate by permanganate was further extended into the weakly acidic region. It was thought that it would be possible to detect the reaction of permanganate ion with the unionised substrate. The oxidation rate appears to "flatten out" (becomes essentially pH independent) below pH 6.1, passes through a minimum at pH 2.5 and then rises slightly again at pH 0.7. The kinetic results are presented in Table XIII, and the plot, including the complete region investigated, is shown in Figure 16 presented later in the text.

Table XIII

Oxidation of Fluoral Hydrate in Weakly Acidic Region - pH 6.12 to 0.7

kaD l.mole" .sec -x l.mole"1. sec_ x ka/kD.

,-5 6.12 1.69 x 10'>- 4 1.50 x 10" 11.3

2.51 2.30 x 10-5 7.41 x IO"6 4.0 0.7 4.08 x IO"5 1.10 x IO"5 3.7

1.0 77

It is interesting to note that deuterium isotope effect still persists

even at pH 2.5 where the rate is slowest and is undoubtedly due to the reaction between unionised fluoral hydrate and permanganate ion.

Activation parameters have not been obtained in this region due to the very slow reactions observed here.

6. Oxidation of l,l)l;5,5i5-Hexafluoro-2-propanol

The pKa of this compound has not been measured)however its value may be determined easily from Haszeldine's data (108) for similar alcohols. The estimated pKa is about 10.5 in 50$ ethanol solutions.

It would be expected that the magnitude of the pKa would be slightly

lower in water (109). The corresponding deuterio compound was prepared

so that the isotope effect could be studied. Both protio and deuterio hexafluoro-2-propanols were oxidised in the region pH 8 to 13. Good

second order rate plots were obtained for at least 80$ of the reaction.

The kinetic data are shown in Table XIV. The unit slope agreement

(not shown) is perhaps not as good as for the fluoral hydrate, however

the number of points is not sufficient to draw any definite conclusions*.

The deuterium isotope effect kn/kD = 19.5 at 25° is of some interest

since it appears to be the highest isotope effect observed thus far.

(a) Oxidation of l,l,l;3,3,5-Hexafluoro-2-propanol in Deuterium Oxide

This experiment was performed in a manner described in the

experimental section. Dauterium oxide was used in one run and another

* The rate of pH 10.7 is probably low. This is ascribed to the interferance of the time dependent disproportionation of manganate. 78

Table XIV

Oxidation of l>l>l;3>3t3-Hexafluoro-2-propanol

ksE k2D pH l.mole"1.sec-1 l.mole"1.sec"1 kH/kp

8:07 0.0054

10.15 0.092

10.69 0.100

12.7 O.376 0.0191 20

13.0 0.382 0.0206 19

13.0 O.38O 0.0194

13.7 0.421

All runs at 25.0 - 0.03°C and ju = 1.0

[Alcohol] o = 0.00284 M for pH 8 to 10.7 and

0.0019 M for alkaline runs

[Mn04" ] o = 0.0019 M for pH 8 to 10.7 and

O.OO38 M for alkaline runs control experiment was made with water. Both rates were identical within experimental error.

kp20 = 0-309

kHao = 0.293

Runs made at 25.00 ± 0.03° at u ="1.0 in 0.1 M NaOH

Data are for the protio compound.

The reaction mixture contained 9^$ D2O by volume.

The results indicate that the permanganate can not be abstracting hydride ion or hydrogen atom from the hydroxyl group since appreciable 79 difference in rate should be observed in this case. Further significance of the results will be discussed fully later.

7. Permanganate Oxidation of 2,2,2-Trifluoro-ethanol

Being a primary alcohol, trifluoroethanol will yield trifluoroacetaldehyde upon oxidation, which will immediately hydrate and be further oxidised to trifluoroacetic acid. In order that the kinetics be interpretable, fluoral hydrate must either be oxidised at much faster or much slower rate than trifluoroethanol. If both compounds are oxidised at comparable rates then the kinetics would be quite complicated. Jbrtunately it was found that fluoral hydrate was oxidised at a much faster rate than trifluoroethanol. This necessitated the use of the 1:4 ratio of trifluoroethanol to permanganate. All runs were made in strongly alkaline solution at 25° and ionic strength

1.0. Good second order rate plots were obtained up to 80$ of the reaction. The results are presented below in Table XV. These results compare very well with those for the permanganate oxidation of the primary alcohol HCF2CF2CH2OH (64). It may be seen that the rate of oxidation of trifluoroethanol decreases with decreasing alkalinity which is in accord with its pKa value of 12.37 (HO). This, of course,assumes that the trifluoroethanolate ion is the reactive species in the permanganate oxidation. 80

Table XV

Permanganate Oxidation of 2,2,2-Trifluoroethanol

fOH"] M l.mole-1. sec -l

0.05 1.44

0.1 1.78

0.1 1.76

0.5 2.22

0.5 2.36

Temp. = 25.00 t 0.05°, jx = 1.0

8. Permanganate Oxidation of Chloral Hydrate

Attempts to measure the rate of oxidation of chloral hydrate in strongly alkaline solutions failed. It was found that even at pH

10 the hydrolysis interferes with the oxidation rate. This is rather unfortunate since it prevents any comparison between the rates of fluoral hydrate and chloral hydrate oxidations. Three runs were made below pH 10; however the results are not too precise due to the slow reaction and small reactant concentration. The results are shown in

Table XVI and for comparison the rates for fluoral hydrate oxidation are included. Although the results are not precise, it can be seen that chloral hydrate is oxidised at a faster rate than fluoral hydrate, which is in accord with greater inductive electron withdrawing power of the trif luoromethyl group as compared to the trichloromethyl group.

The rate for chloral hydrate at pH 8.5 is uncertain, since the possibility 81

Table XVI

Permanganate Oxidation of Chloral Hydrate

Chloral Hydrate Fluoral Hydrate 1 1 1 pH ka l.mole" .sec" kg l.mole" sec"

0.7 0.0005 0.00004

6.6 0.0017 0.00049

8.5 O.O85 O.O39

Temp. = 25.00 t 0.03°, p. = 1.0,

[chloral Hydrate] o = I [Mn04~]0 = 0.0015 M

of hydrolysis cannot be completely excluded. On the other hand a good second order plot passing through zero was obtained at that pH, so perhaps the extent of hydrolysis is not too serious.

9. Permanganate Oxidation of Substituted Benzhydrols

Stewart has previously oxidised unsubstituted benzhydrol and its 0(-d analogue in essentially the whole pH region (17). He found an isotope effect kH/kp =6.6 from which it follows that the C-H bond is broken in the rate determining step. 0xygen-l8 experiments have shown that there is no oxygen transfer from the labeled permanganate to the benzhydrol. On the basis of the evidence a mechanism, involving hydride ion transfer to permanganate was postulated. Since these results appeared very promising, it was decided to extend this investigation into the study of the effect of substituents on the rate of oxidation. Several 4,4'-identically substituted benzhydrols 82 were chosen with substituents ranging from dimethlamino to nitro groups. It was found upon closer investigation that all of the substituted compounds had only slight..water solubility necessitating' the use of the spectroscopic method in the kinetic study. The very low solubility prevented any extensive investigation. All runs were made in 0.1 M sodium hydroxide at 25.0 i 0.03° at ionic strength of

0.1. It was found that the oxidation of blank (containing permanganate but no substrate) was appreciable in the case of slow rates. For this reason it was necessary to apply a correction to the kinetic data. This was done simply by adding the optical density of the blank to the optical density due to permanganate absorption and subtracting it from the manganate. This procedure is not too rigorous, however is probably the best under the circumstances, since the oxidation of the blank was neither first or second order. Further complication arose, since the ketone, formed during the reaction, precipitated from the solution at about 50$ reaction. It was possible to observe this visually as well as on the rate plot, which exhibited a sharp break in the straight line. Thus most of the reactions could be followed to only 50$ completion. Ten centimeter cells were tried, containing ten times diluted reactants, however it was found that under these conditions the blank oxidised at a rate comparable to the substrate. This method had to be abandoned. Dichlorobenzhydrol was the most insoluble of all the compounds studied, and for this reason an attempt was made to correlate some rates at the same concentration as for dichlorobenzhydrol. Unsubstituted benzhydrol gave the same rate 83 as at higher concentrations, however the correlation was not so striking for dimethoxy^benzhydrol. Deuterium isotope effect was determined for several of the "benzhydrols. All attempts to measure the rate of oxidation of 4,4'-bis(dimethylamino)benzhydrol and the o/ -d compound failed. This was due to overoxidation which probably occurred at the dimethylamino function. Moreover the rate of oxidation was too fast to be measured by ordinary technique; permanganate appeared to be reduced directly to manganese dioxide without going through the manganate stage. The oxidation products were analysed by ultra-violet spectroscopy, but the absorption which would be due to the Michler's ketone was not observed. On the other hand it was noted that in the solid state, Michler's hydrol

(4,4'-bis(dimethylamino)benzhydrol) was easily oxidised by air, forming a yellow coating on the surface of the crystals, which could be identified as the Michler's ketone. The corresponding 0( -d compound appeared to be oxidised by air at much slower rate. It is unfortunate that the oxidation of the hydroxyl function only could not be accomplished in solution. Kinetic data for the protio and deuterio benzhydrols as well as the isotope effects which were evaluated in some cases are shown in Table XVII. Closer examination of the data reveals several trends. The rate Increases with both electron with• drawing and electron donating substituents. At first glance this is the same trend in an exagerated form, as was observed by Stewart and

Van Der Linden (64,65) in the effect of substituents on the oxidation phenyl-trifluoromethyl carbinols. However the correlation is not so 84

Table XVII

Permanganate Oxidation of 4,4'-identically Substituted Benzhydrols

4,4'-Substituent

and Deuterium [Benzhydrol] o k^ x kzj) Position M x 104 l.mole-1.sec" l.mole"1.sec" kH/kn

H 1.33 10.53

H 1.33 9-8 6.1 H-

H- tf-d 1.33 1.74

Me 1.33 9.40

Me 1.33 8.64 6.8 Me- o( -d 1.33 1.37 J Me- o( -d 1.33 1.28

OMe 1.30 18.2

OMe 1.30 20.2 6.4

OMe- o( -d 1.37 3.0 (a)

Cl 0.40 34 (a)

N02 0.85 329

wo2 O.85 274 (a) 4.9

N02- c< -d 0.43 49

N02- c< -d 0.43 66 (a)

H 0.40 (b) 8.34

H 0.40 (b) 8.77

OMe 0.40 (b) 12.1

0 All runs at 25.0 t O.O3 in 0.1 M NaOH, ju = 0.1, [Mn04"]o = 2 [Benzhydrol] (a) Average of two runs (b) Runs at low concentration 85 simple as it appears. Their rates were compared at full ionisation of

the alcohols which ia certainly not the case here. The increase in the oxidation rate of benzhydrols with both, electron with drawing as well as electron donating substituents may be explained reasonably well on the basis of the difference in stability of the transition state. This point is further expanded in the Discussion section. However since the rates for the undissociated or completely ionised benzhydrols are not known, any quantitative conclusions are therefore absent.

Of interest in the relatively small isotope effect for dinitrobenzhydrol as compared to other benzhydrols. The decrease in the magnitude of the isotope effect with increasing rate is in accord with Wiberg and Slaugh's prediction (99). However this isotope effect

is not strictly compareable since they were obtained at the point where

the benzhydrols are only partially ionised.

10. Substituent Effect on the Bate of Permanganate Oxidation of Compounds of the Type CFaCHRiRg

Even though only a small number of compounds were investigated

it was possible to vary the nature of substituents greatly. In

combination with Stewart and Van Der Linden's data (64,65) a better idea

'about the effect of substituents is obtained. Inspection of the data,

summarised in Table XVIII, reveals that the oxidation rate rises in the

series Ri = H > OH > CF3 and R2 = 0". This trend would be predicted on the basis of the electron properties of the substituents. This effect

is much more marked in the series: undissociated fluoral hydrate <^

fluoral hydrate mono-anion < fluoral hydrate di-anion, where the rate 86

Table XVIII

Effect of Substituents on the Rate of Permanganate Oxidation

of Compounds of the Type CFaCHRiRg

Ri I

CF3 — C — E

1 -1 *2 k2 l.mole" .sec References

OH OH 0.00002 Present work

Ph OH 0.001 64,65

CF3 0" 0.42 1 OH o- 1.3 [ Present work H 0* 2.3 1 m-Br-Ph o- 7.7 (a)

Ph o- 7.9 (a)

p-CH3-Ph o- 8.9 (a) 6k,65

m-N02-Ph 0" 8.9 (a)

p-CH30-Ph 0" 9.6 (a) J

0" 0~ 103 - 104 (b) Present work

Data for protio compounds only.- All values rounded off.

(a) At yu = 0.2 (b) Estimated rate at full ionisation rises by a factor of about 105 and 103 - 104 respectively. This is

expected from the strong electron donating power of the 0" group. The

full significance of these results, in connexion with other evidence, is

elaborated in the Discussion section. 87

11. Permanganate Oxidations in Concentrated Sulfuric Acid

(a) Determination of PKR. of Permanganic Acid

The pKa of permanganic acid .has been determined previously in the system perchloric acid-water (31). In this investigation, potassium sulfate was used for adjusting the ionic strength and sulfuric acid was added to phosphate to adjust the pH in the rate experiments. For this reason the pKg, of permanganic acid was determined in sulfuric acid-water system.

The plot of optical densities (at 458 and 525 mu) versus the

H- function showed the behaviour of typical ionisation curves

(Figure 9). The pKa was calculated from the data at 525 and 458 mu (31) and then by modified Davis-Geissman method (75), using optical data at both wavelengths (Figure 10, 11, 12). Since the plot of log ^B" - ^525 versus H- did not show a slope of unity, the €\ 525 - £ BH logarithms of the extinction coefficients were also plotted against the Ho and HR (JO) functions in order to determine whether the dependence on the H_ function is real. It was found that the slope was essentially the same when Ho or H_ function was used. This is not surprising, since both functions are almost parallel in the region studied. The slope, using HR function was however much less than unity, and for this reason this relation was rejected. The optical data used in the calculations are shown in Table XIX and the pKa, summarised in Table XX. The pKa calculated with the aid of HR function

FfG. 12 pKa of Permanganic Acid (modified Davis-Geissman method) 92

Table XIX

Spectroscopic Data for Ionisation of Permanganic Acid

Optical Extinction Optical Extinction Density Coefficient Density Coefficient

i H2S04 at 525 mu at 525 mu at 458 mu at 458 mu / Neutral O.761 2370 O.O99 308 Solution

Ul.5 0.740 2310 0.102 310

52.2 O.63O 1965 0.107 331

57.6 0.515 1603 0.123 583

62.9 O.362 1178 0.160 h99

67.5 0.212 660 0.186 579

72.5 0.142 443 0.199 620

82.2 0.060 187 0.199 620

97.0 0.041 128 0.158 493

£ Mn04"J = 0.000321 M (in water)

Values obtained in 97.0 $ H2S04 were not used in calculation of the

pKa- Measurements made in one cm cells. Spectrum of Mn04" remains unchanged from aqueous solutions up to 30$ KzSQ*. 93

Table XX

Summary of pK^'s of Permanganic Acid Evaluated With

Aid of Hp, H. and HR (JQ) Functions

Acidity Function Used Wavelength in Evaluation mu - PKa Slope of pKa 525 5.10 0.68 H_

458 5.11 1.06 H-

525 4.57 O.65 H0

458 4.61 I.09 Ho

525 9.15 0.3 HR

458 9.26 0.5 HR

Evaluation of pKa using modified Davies-Geissman method.

Acidity Function H_ Hp HR

- PKa 5.10 4.57 9-17 Slope 0.71 0.71 0.32 was rejected on the basis of the small slope. The pKa calculated with the H- function was accepted on the basis of theoretical considerations.

The low slope (using data at 525 mu) is perhaps due to interference of the ionisation:

+ + HMn04 + H Mn03 + H20 which is presumably taking place at high acidity (4l). This process

may be occurring in 97$ H2S04, where the optical density of k^Q mu decreases after passing through a plateau at 72.3 and 82.2$ H2SO4, 94

(Table XIX). The decrease of optical density is real as was confirmed by several experiments with varying concentration of permanganate and slightly different experimental conditions. The smaller value for the pKa of permanganic acid determined by Symons et al. (31) is not in disagreement with present work since the acidic medium is different in both investigations. Stewart and Lee (ill) have found that the pKa of chromic acid was markedly influenced by the acidic medium

(sulfuric acid, perchloric acid, acetic acid-sulfuric acid),

(b) Permanganate Oxidation of Fluoral Hydrate in 25 to 46$ Sulfuric Acid

Fluoral hydrate was oxidiBed by permanganate in the region

25 to 46$ H2SO4. Iodometric procedure was used to follow the reaction kinetics, as described in the experimental section. Permanganate is reduced to Mn02 which appears to stay in the solution at least for the time of the oxidation run. Mn02 finally precipitates out after long standing. The valency change of permanganate was established titrimetrically by following a reasonably fast oxidation reaction to about 99$ completion. The final titer corresponded to the theoretically predicted value (2/5 Vo), indicating that MnIV is the product of permanganate reduction. The oxidation reaction was found to be of second order, and good second order rate plots were obtained up to 50$ of the reaction. No attempt was made to establish whether the rate is first order in each reactant. Since permanganate decomposes slowly in this region, the decomposition of the blank was followed alongside each run and a suitable correction was applied to the oxidation rates.

The kinetic results of the oxidation of protio and deuterio fluoral 95 hydrates in strong acid are summarised in Table XXI.

Table XXI

Oxidation of Fluoral Hydrate in 26 to 46$ Sulfuric Acid

k2H kaD 1 -1 x $ H2S04 l.mole" .sec l.mole" .sec" WkP 46.5 1.25 0.182 6.8 44.1 O.587 0.0890 6.6 41.8 0.272 0.0428 6.4

54.8 O.O516 0.0054 5.9 25-5 O.OO545

[FLuoral Hydrate! 0 = 1 [ Mn04"] 0 = :0.06l2 M 2

Temp. = 25.0 t O.lfi

The plot of log k2 versus Ho function gave a straight line with the slope of 1.7. The slopes using H_ and HR function were 1.2 and 0.8 respectively (Figure 13). The isotope effect for the oxidation reaction in strong acid is about 6.4 on the average, although the tendency to decrease with decreasing acidity is quite apparent. This may be due to the interference of the oxidation reaction of unionised fluoral hydrate (for which knAp 4) which would affect more the slower rate of the deuterio compound. (c) Activation Parameters for Oxidation of Fluoral Hydrate in 26 to 46$ Sulfuric Acid

The variation of the oxidation rate with temperature was

determined at 4l.8$ H2S04. Any pre-equilibrium step which might be FIG. 13 Oxidation of Fluoral Hydrate

0 in 25 to 46% H2 S04 Relation between rate constant

and o-H_J«-H0 .•-HR

Add -4 for HR scale -1 -I log k2 -2H

-3-

-4

o

-5 -I 1 ' I r -4 0 12 3 H0 H_ HR 97 occurring will be relected in the energy and entropy of activation for the rate determining step. The kinetic data appear in Table XXII

and the plot of the log k2/T against l/T is shown in Figure 14. .

Table XXII

Activation Parameters for Oxidation of Fluoral Hydrate

in 26 to 46$ Sulfuric Acid

Temp. °C l.mole- .sec"1

20.0 0.187

20.0 0.188

20.0 0.189 (a) 25.0 0.272

35.0 O.565 35-0 O.584 40.0 O.927

Data for the protio compound, at 41.8$ H2SO4

(a) Reaction with exclusion of light.

AH* = 12.9 kcal mole"1

AS* = -17.8 cal mole-1 deg"1

The possibility of any light catalysis was eliminated, since the rate for the oxidation reaction in the light was the same as in the dark.

The values of the energy of activation for the oxidation reaction in strong acid is about the same as for the reaction of the mono-anion of fluoral hydrate. The value of the entropy of activation is unexpectedly 98 99 high (and negative) for a reaction of an ion with a neutral molecule or of two uncharged species. However, the oxidation mechanism pro•

bably involves establishing of an equilibrium (HMn04/Mn04~) prior to the rate determining step .which would affect the magnitude of the entropy of activation to some extent. Assuming a maximum value for the rate constant of 10 to 100, would lead to the entropy of activation between - 13 to - 8 e.u. respectively, which would be more reasonable values for the ion/molecule or neutral/neutral species reaction.

12. Summary of the Permanganate Oxidations of Fluoral Hydrate

Permanganate oxidations of fluoral hydrate were investiga• ted throughout a wide region from 46$ sulfuric acid to 1.0 M sodium hydroxide. The plot of rate constants for the protio compound versus pH and H- are shown in Figure 15. The identities of the reactive species are not immediately apparent from this plot. A better picture is obtained when logarithm of the rate constants is plotted against the pH or H- functions. This is shown in Figure 16. The reaction of the di-anion of fluoral hydrate with permanganate may be clearly recognised in strongly alkaline region and that of mono-anion can be seen in the region pH 6 - 11. Unionised fluoral hydrate reacts with permanganate in the region pH 1 - 5 and at higher acidities yet another reaction may be seen. The latter reaction will be analysed in the

Discussion section. FIG. 15 Oxidation of Fluoral Hydrate Relation between rate constant and pH, H_ FIG. 16 Oxidation of Fluoral Hydrate Relation between log of rate constant and pH, H

o CF3CH(OH)2

• CF3CD(OH)2

-5 H. 0 PH 10 15 102

13. Behaviour of MnIV IH Concentrated sulfuric Acid

Since the majority of permanganate oxidations of formic acid

in concentrated sulfuric acid were followed by a spectrophotometric method, it was necessary to establish the absorption pattern of the

reduction products of permanganate. In addition, several other

interesting observations were made, which are discussed below.

In weakly acidic or alkaline aqueous solutions, any

formed during a reaction, precipitates out of solution. Various forms

of hydrated manganese dioxide have been extensively studied (ref. 2

and refs. therein). An aqueous solution of Mn*V which appears to

follow Beer's Law has been recently reported (112). Near pH 7 the

solution has considerable stability and remains clear even after

boiling. The U.V. spectrum shows a peak at about 390 mu (£ = 9000) with a fading absorption extending beyond 660 mu. The spectrum is

attributed to the anion H2Mn04~.

In 6% oleum, permanganate decomposes to Mn*V which forms

stable blue solution (42). The Mn*^* compound is monomeric and exists

probably as the uncharged sulphate.

The present investigation concerns itself with the behaviour

of Mnlv in concentrated sulfuric acid. Permanganate is reduced to

quadrivalent manganese in the oxidations of fluoral hydrate and formic

acid in concentrated sulfuric acid. 5br experimental purposes, the

Mn-"-v solution was prepared by adding 5 ml. of 97$ sulfuric acid to

5 ml. of 0.1 M permanganate solution. It was noted that the 103 permanganate was decomposing only very slowly for about two minutes.

After this time the decomposition became almost violent, accompanied by visible sudden evolution of gas (oxygen). This behaviour is in accord with an autocatalytic reaction, which is known to take place in the Mn02 catalysed decomposition of permanganate. In the concentrated acid the Map? stays in solution and its catalytic effect is probably more marked, than in the case of the Mn02 precipitate. The Mn^ solution prepared in the above manner was suitably diluted for spectroscopic experiments. The species forms a clear brown solution in the region 20 to 70$ sulfuric acid. Beer's Law is obeyed as can be seen from the following data:

[MnIV] M £ X 23Q 0.0001 4250

0.0002 4750

0.0003 4700

IV in 62$ H2S04. The spectrum of Mn in 62$ sulfuric acid is shown in

Figure 17. It resembles the spectrum for the MnIV in alkaline

solution (112), however the extinction coefficients are somewhat lower and the wavelength ( X max) of absorption shifted to shorter wavelengths. When a sample of the MniV solution (62$ sulfuric acid) was dissolved in 97$ sulfuric acid, the brown colour changed slowly

to blue. The spectrum of this blue solution (Figure 18) was identical with that obtained by decomposing permanganate in 3°$ fuming sulfuric

acid, or dissolving the brown MnIV solution or solid Mn02 in fuming

acid. It is also identical with the blue MnIV species reported by

106

Symons (42). In addition the blue Mn*V species is obtained upon addition of solid manganous sulphate to a green solution of permanganate in 97$ sulfuric acid. This is an example of the Guyard reaction taking place in concentrated sulfuric acid. It is quite likely that the slow change of brown MnIV into blue MnIV in 97$ sulfuric acid is a slow dehydration process. This may be envisaged as taking place in the following manner:

+ + Mn(0H)4 + 3H ^ Mn02H + 2 HaO*

+ 1 or MnO(0H)2 + 2H+ v i i ^ Mn02H + HaO" "

Thus the mono- or di-hydrates of manganese dioxide are protonated in a fast step followed by subsequent slow elimination of water. The

+ above postulated species Mn02H might imply that it should be possible to dissolve manganese dioxide in concentrated sulfuric acid. It is of interest to note that aged solid manganese dioxide appears to be insoluble in 97$ sulfuric acid, however it dissolves (with difficulty) in fuming acid to give the blue MnIV species. The reason for this apparently contradictory behaviour is not known at present; it may be perhaps due to the small solubility of aged solid manganese dioxide in 97$ sulfuric acid.

When the brown MnIV solution in 62$ sulfuric acid is heated at about 150° for 15 minutes, the solution slowly changes colour to red-violet. The spectrum of this solution is shown in Figure 19.

Prolonged standing of Mn*V solution at room temperature also produces the same red-violet colour. Assuming the same concentration of

108 IV Mn , the appearance of red-violet colour is faster and more intense with increasing acidity. Thus solutions of ifaIV in 25$ sulfuric acid show appearance of a weak red colour after several weeks, whereas in about 60$ sulfuric acid the reddish-violet colour appears more intensly after several days of standing. The reddish-violet species may be produced by addition of solid manganous sulphate to a solution of MnIV in 62$ sulfuric acid. This reaction appears quite fast. If the above reaction may be interpreted in terms of

MnIV + Mn11 nj •• • ^ 2 Mn111 then the red-violet species is identified as Mn***. Indeed when the red-violet solution is poured into water, colloidal manganese dioxide appears after a short time, which is in accord with the instability of

111 Mn below 6 N H2S04 (113)• Addition of excess of formic acid to the red-violet solution produces a berry red colour which disappears after several days. This may be interpreted in terms of complex forma• tion between Mn111 and formic acid, and a subsequent slow oxidation of this complex (cf. Ref. 16). This is also analogous to the complex formation between Mn11* and pyrophosphates (16). The observation that a stable complex is formed between Mn*1* and formic acid disposes of any possibility of induced oxidation by Mn111, at least in the region of 25 - 70$ sulfuric acid. When the red-violet Mn111 solution is intro• duced into fuming sulfuric acid, the colour does not change.

The blue and brown Mn*V solutions oxidise and sulfur dioxide (added as bisulphite) quite rapidly. The reduction of Mn*V is assumed to proceed to Mn1* stage, since the resulting 109 solutions are colourless. These are presumably the first examples of homogeneous oxidations by Mn*"^.

Ik. Permanganate Oxidation of Formic Acid in 20 to 72$ Sulfuric Acid

The behaviour of formic acid towards permanganate was previously established throughout the weakly acidic to the strongly alkaline regions (55,56,57,58). The reaction of formic acid was recently reinvestigated in the weakly acidic region up to 1.0 M perchloric acid (59,60,6l) where the contribution of the unionised formic acid- permanganate reaction was detected. The present investigation adds further to the knowledge of behaviour of formic acid towards permanganate in the region 20 to 72$ sulfuric acid. Iodometric and spectroscopic methods were used to obtain the rates. All rates in strong acid appear to accelerate with time after about 50$ completion of the reaction.

This is understood in terms of Mn*^ catalysed decomposition of permanganate described earlier. In the region 20 - 30$ sulfuric acid, second order kinetic plots were curved and for this reason pseudo first order kinetics (which gave straight line plots) were used here. The reaction of formic acid with permanganate was established to be first order in each reactant (55 "to 6l) and the same conclusion was reached in this investigation. No attempt was made to obtain rates in the region higher than 72$ sulfuric acid since the dehydration of formic acid might interfere in this region (115,114). 110

(a) Permanganate Oxidation of Formic Acid in 20 to 42$ Sulfuric Acid

In this region the oxidation rate increases with increasing acidity. The rates are listed in Table XXIII and Halpern and Taylor's

(59,60,61) data for the weakly acidic region are included for compari• son.

Table XXIII

Permanganate Oxidation of Formic Acid in 0.1 to 1.0 M Perchloric Acid

(59,60,61) and 20 to 42$ Sulfuric Acid

[HC104] Corrected M H_ H_(a) l.mole" .sec"

0.1028 + 1.10 + 0.93 0.0162 (b)

0.1696 + 0.92 + 0.74 0.0106 (b)

O.2569 + O.74 + 0.60 0.00728 (b)

1. 029 - 0.21 - 0.21 0.00264 (b)

kgD 1 1 i H2S04 S_ l.mole- .sec" l.mole .sec kH/kn

20.3 - 1.03 O.OOO83 2.2 (c) 20.9 1.08 0.00182

28.3 - 1.80 O.OO3O3 O.OOO87 3-5 (c)

31.3 - 2.10 0.00718 0.00175 4.1 (d)

36.0 - 2.55 0.0173 0.00625 2.8 (e)

38.0 - 2.73 0.0282 0.00715 4.0 (d)

41.5 _ 3.08 0.0614 (e)

(a) Corrected for addition of NaC104. Effect of NaC104 assumed same as for Ho (ll6). (b) Halpern and Taylor's data at 25.8°, u=1.0 (59,60,6l). _ (c) Pseudo first order kinetics [HCOOH] 0 = 0.1M, [Mn04 3o = 0.00133 M

(d) Second order kinetics, iodometric method [HCOOHlo = 3/2[Mn04"JJo = 0.01 M

(e) Spectral method [HC00H]O = 3/2 [Mn04" ]0 = 0.001 M Ill

The H, values for Halpern and Taylor's data had to he corrected for the effect of added sodium perchlorate. Since the effect of salt on

H_ function is not known, it was assumed that the effect is the same as for the Ho function (ll6). This is satisfactory for the purposes of present discussion.

The relation of the logarithm of the rate constant to the H_ function is depicted in Figure 20. Unit slopes are obtained for the

HCOOH and DC00H oxidation in the region 20 - k2$ sulfuric acid. Only one point at the H_ value of about - I.03 deviates from the straight line relationship. Halpern and Taylor's data show unit slope for three points, however the region in which this occurs, covers only 0.3

H_ units. The last point lies clearly off the unit slope. This is in apparent contradiction with the straight line relation which was obtained for a plot of rate constants versus l/ [H+] (59). If the rate constants had been plotted against h-, the slight deviation from straight line plot would have been observed, even though the relation is rather insensitive at high acidities. It is not immediately apparent its? •

from Figure if where (and if) the reaction of HC00H/Mn04- occurs. The

problem of determining the rate constant for the reaction HC00H/Mn04~ may be approached in the following way. The unit slopes

(HC00"/Mn04" and HC00H/HMn04 reactions) intersect at H_ = - 0.8 which corresponds to the value of log k2 = 3«5 (See Figure 20), and

k2 = O.OOO316. The total rate at this point is equal to 2 x O.OOO316.

Substracting O.OOO632 from the lowest point of the curved portion of

the plot (log k2 = - 2.8)-yields a value for the rate constant for the FIG.20 Oxidation of Formic Acid in Perchloric and Sulfuric Acids

Plot of logka vs H_ -2-

rt XL CD O -3- Data nef. 59

O HC02H

3 DC02H -4- Extrapolated

~~i— 0 -1 -2 -3 H_ ro 113

reaction HC00H/Mn04" which is estimated as O.OOO96 or k2 & 0.001.

Although the rate constants for the reaction DC00~/Mn04~ are not available, the isotope effect for this reaction is known (kg/kD = 7).

Thus a line with a unit slope may be drawn and its intersection with the straight line of the deuterio compound of present investigation yields

the value for the rate constant, k2 ~ O.OOOO65. Application of the same procedure as above yields a value for the rate constant (for the

reaction DC00H/Mn04") which is estimated as O.OOO67. From these data the isotope effect kg/lCD ^ l.k may be estimated for the oxidation of HCOOH and DC00H by permanganate ion. This is in accord with the isotope effect of about k which is observed in the oxidation of unionised fluoral hydrate with permanganate. Halpern and Taylor predict an isotope effect of 1.06 for the oxidation of unionised formic

acid with permanganate (59). That the presence of Mn04" - HC02H reaction is real, is apparent from the decreasing isotope effect (60).

This was also observed in the present investigation. A plot of rate constant versus h- (shown in Figure 21) depicts straight lines for the protio and deuterio formic acids which appear to intersect at zero acidity. The precision of the data does not allow any estimation of the rate for the unionised formic acid with permanganate which should be observable at the intercept. The estimation of this rate is achieved better from the logarithmic plots, which are based on theoretical considerations. It is of interest that Figure 21 would

_ predict the reaction HC00H/Mn04 at zero acidity, whereas Halpern and Taylor (59)-prediet the same at high acidity. These results are

115 not in contradiction, but follow simply from the rate dependence on acidity, which has reverse effect in both investigations. The nature of the reactive species observed in the present investigation of oxidation of formic acid will be discussed fully later.

(a) Activation Parameters for Permanganate Oxidation of Formic Acid in 20 to k2% Sulfuric Acid

The determination of the temperature dependence of the rate was made in 38$ sulfuric acid, using second order kinetics and the iodometric method. The data for HCOOH are presented below

CT°£ ka

15.4 0.0141

25.0 0.0282

38.8 0.0871

and the plot of the log k2/T versus 1000/T is shown in Figure 22.

The heat and entropy of activation were computed from the plot:

AH* = 13 kcal/mole

AS* = -21 e.u.

Both of these values are composite since they include any pre-equilibrium step. The A H* is comparable with the value obtained for the oxidation of unionised fluoral hydrate in concentrated sulfuric acid, however the entropy of activation is unusually large and negative for a reaction of a neutral molecule with any permanganate species. Assuming a higher value for the rate constant, the entropy of activation would have a smaller magnitude. 116

~A'A A 1 1 1 1— 3*1 3*2 3*3 3*4 3*5 1000/T°K 117

(c) Permanganate Oxidation of Formic Acid in 50 to 72$ Sulfuric Acid

The rate of oxidation of formic acid in 20 to 42$ sulfuric acid followed well the H_ function. Further increases in acidity produce a more rapid rise in the rate than would be predicted by the unit slope. This can be tentatively ascribed to a reaction between formic acid and some species of MnVH. The rates are summarised in

Table XXIV.

Table XXIV

Permanganate Oxidation of Formic Acid in 5° to 72$ Sulfuric Acid

k2H k2D -1 - 1 1 1 $ H2S04 HR l.mole .sec 1.mole- -sec"

52.2 - 7.05 2.3

57.6 - 8.5O 20.1 3.2 5.5 - 0.2

62.9 - 9.67 105 14.2 7.4 + 0.5

67.5 -10.87 272 32 8.5 ± 1

72.5 -12.12 495 51 9.7 i 2

Temp. = 25 - 0.1°C

[HCOOH] 0 = 3/2 [ Mn04" ] 0 = 0.0005 M

The isotope effect increases with increasing acidity, however the values of the rate constants are subject to great error due to the very fast reactions, which can not be conveniently followed by conventional methods. The rates for the protio compound at 62.9, 67.5 and 72.3$ sulfuric acid should be regarded as approximate only. The fact that an isotope effect is observed even in 72$ sulfuric acid 118 eliminates the possibility that the formic acid is being dehydrated to any appreciable extent. The logarithm of the rate constant for

HCOOH appears to follow approximately the HR (JO) function with a slope of about 0.8 (Figure 23). The implication of this relation will be discussed later. An attempt was made to determine the activation parameters for this reaction in 52.2$ sulfuric acid; however a curved line was obtained for the Arrhenius plot. The data are shown below:

Activation Parameters for Permanganate Oxidation of

Formic Acid in 52.2$ Sulfuric Acid

T?C kg

12.0 1.68

25.0 2.3

32.5 3.4

Concentrations same as in Table XXIV.

The values of heat and entropy of activation were estimated from a tangent drawn through the point: k2 = 2.3, Temp. = 25°

AH* = 6-2 kcal/mole

AS* =-38 - 10 e.u.

Both of these values probably include some pre-equilibrium step and might be smaller in magnitude. The curvature of the Arrhenius plot is probably due to the contribution of the reaction at lower acidity which has a high heat of activation. 3H FIG. 23 Oxidation of Formic Acid

in 50 to 72% H2S04 2H Relation of log k, with H

OH

O

-H O HC02H

3 DC02H

•2H

—r 11 9 1 3 7 9. H VO

4 120

DISCUSSION AND CONCLUSIONS

1. Mechanism of Permanganate Oxidation of Fluoral Hydrate in Weakly and Strongly Alkaline Solutions

Fluoral hydrate is much more acidic than the o<-phenyltri- fluoromethylethanols (64,65,66), benzhydrol (17) and substituted benzhydrols mainly because of the great electron withdrawing power of the trifluoromethyl (86) and hydroxyl groups.

The available evidence shows that the kinetics of oxidation of fluoral hydrate by permanganate are first order in each of the reactants. The rate exhibits first order dependence on the hydroxyl ion and is accommodated by a rate law of the form

V = k [CF3CH(0H)2] [Mn04~] [ OH" ] (l)

A positive salt effect is observed for the oxidation of fluoral hydrate in 0.1 M sodium hydroxide at 25° (Table VIII). The above results are consistent with the following reaction path in the region pH 6 to 11

OH 0" 1 K 1 CF3-C-H + OH" v CF3-C"H + H20 (2) I - I OH OH

0" I k CF3— C-H + Mn04" - — > Products (3) I slow

OH

The rate law corresponding to this mechanism is

0"

- a fMn04'3 = k [CF3CH(0H)j [ Mn04~] (4) dt 121

and K = [ CFaCH(OH)~l [ HgOJ (5)

[CF3CH(OH)2] [0H"J

Substitution for the concentration of fluoral hydrate amion into the rate expression yields

- d I MnCuJ = kK [ CFaCH(0H)21 C QH~ 1 [ MnQ4" 1 (6) dt [ HJJOJ

and kexperimental = ^ [0H~1 = HKJ r0H-j (?) where K* is the first ionisation constant of fluoral hydrate.

At complete ionisation:

0" - I

[CF3CH(0H)21 = [CF3CH(0H)J (8) and the rate, which may be expressed as

Rate = k2 [CF3CH(0H)2"j [ Mu04" J (9) becomes independent of the hydroxyl ion concentration, and the

experimental rate constant becomes identical with k.

A similar rate law may be derived forvthe oxidation of fluoral hydrate in the region pH 12 to Ik. 0"

- d T MnQ4" I = k [ CF3- C - H ] f Mn04" ] (10) dt I 0" and the rate may be written as

0"

11 - d r MnQ4-1 = kK [CF3CH(0H)1 [ OH"] f Mn04"J (ll) dt ~X Kw 122 the rate will show only first order dependence on the hydroxyl ion

(Figure 7) since fluoral hydrate is almost completely ionised at pH 12 and above (pKa = 10.0 at 25°, jx = 1, first ionisation).

Activation Parameters.

The activation parameters or the oxidation reaction in the weakly alkaline region were obtained at pH 10.2 and corrected for the ionisation equilibrium (Table VI, Figure 6) assuming that the rate determining step takes place between the aldehydrol anion and permanganate ion :(equ. 3) • The energies of activation are. similar to those observed in other permanganate oxidations (17,58,59,64) and the entropies of activation are reasonable for a reaction between two amion6 (117)• The corresponding uncorrected energy and entropy of activation for the protio compound are 13.4 kcal mole"1 and -14.4 e.u. respectively.

Similarly the energy and entropy of activation for the oxida• tion of fluoral hydrate in the strongly alkaline region were evaluated by assuming that the rate law may be represented by equation (10)

(Tables X and XI and Figure 8). Both thermodynamic quantities are considerably smaller than those for the first ionisation.

Isotope Effect.

A considerable deuterium isotope effect (10 to 14) was observed for the oxidation of fluoral hydrate in the weakly alkaline region. This shows that the mechanism involves breaking of the C-H bond in the rate determining step. 123

The normal isotope effects which result from the loss of CH stretching vibrations in the transition state are much lower at the same temperature (ll8). Several possible explanations of the high isotope effect will be considered. The idea that the CH, CD stretch• ing frequencies are themselves anomalous may be rejected upon the examination of the IR spectrum. It is found that the values of the

CH and CD stretching vibrations appear slightly higher than normal

(2980 cm-1 and 2220 cm-1 respectively). However the ratio

VCH/VCD = 1'35 is normal. In this connexion Stewart and

Van Der Linden (6k) have found normal CH, CD stretching frequencies

1-phenyl-trifluoromethyethanols; the isotope effect of 16 was observed for the.permanganate oxidation of this compound. This value is somewhat higher than obtained in the present Investigation.

The possibility of a cumulative isotope effect arising from some free radical chain branching does not appear likely in view of the excellent second order kinetics observed in this investigation.

The somewhat higher isotope effect is best understood in terms of partial loss of the CH, CD bending modes in the transition state in addition to complete loss of the stretching vibrations

(106,107). The idea becomes more attractive when the possibility of proton tunnelling is discounted on the basis of the normal temperature dependence of the isotope effect. The corresponding isotope effect for the oxidation of fluoral hydrate in strongly alkaline region is only half the size of that observed for the oxidation reaction in the weakly alkaline region. 12k

Substituents Effect.

Further insight into the electronic requirements of the rate determining step may be provided by the effect of substituents on the rate of oxidation. The results in Table XVIII show the trend in the

series CF3CHR1R2. Comparing the rates of fully ionised substrates,

the sequence CF3, 0" < OH, 0" < H, 0" < 0~, 0" shows that the rate

increases with electron donation to the transition state. These results also indicate that the mechanism is the same for the compounds

listed.

Solvent Effect.

The effect of deuterium oxide on the rate of oxidation of hexafluoro-2-propano1 was examined. This compound exhibits almost

identical behaviour towards permanganate as does fluoral hydrate and

the postulates may be applied with confidence to the mechanistic

interpretations of fluoral hydrate oxidations.

The fact that the rates of oxidation in deuterium oxide and water are the same proves that permanganate does not remove hydride

ion or hydrogen atom from the hydroxyl group. In addition these

results suggest that the solvent is not abstracting a proton from the

C-H bond. If this were the case a solvent isotope effect would be

observed since deuterium oxide is a weaker base than water.

Mechanism.

All the above evidence appear to be consistent with a mechanism

involving hydride ion transfer to permanganate. Thus in the weakly 125 alkaline region (pH 6 11) the rate determining step may be represented as follows:

CF3 CF3

- _k I "0-C-^H> Mn04 e (12) C + HMn04 I slow // \ OH 0 OH

= fas + HMn04 + Mn04~ V 2 Mn04= + H (15) and in the strongly alkaline region (pH 12 - 14)

CF3 CF3

Mn04" C + HMn04" (14) I slow 0. 0 0

The presence of the two negative charges on the flural hydrate anion

(equation 14) would facilitate greatly the transfer of hydride ion to permanganatej this is reflected in the low energy of activation

(6.6 kcal mole-1 for the protio compound). The high negative value for the entropy of activation (-51J4 e.u.) is in agreement with a reaction between anions.

Another mechanism which satisfies the kinetic evidence is considered below:

CF3 CF3

C-H + Mn04" C.© - Mn03 (15) I 'H -0 OH •0 OH

CF3

s C + HMn04 + Mn03" / \ /\ OH '0 OH OH 126

The equation (15) represents a triangular transition state (58) whereby the CH or OMn bonds may be broken depending on the course of the reaction. It has been suggested that such triangular intermediate would lead to a small isotope effect (119,120); since this is not the case in the present investigation this mechanism may be considered unlikely. It would be difficult to check for the 018 transfer (from permanganate to substrate) occurring by a path a due to the rapid equilibrium:

CF3 CF3 CF3 CF3

c c JSi c =ji c (16) /'\ /\ /\. HO OH OH _0 OH OH 0 . OH 0 0 as well as to the lack of suitable insoluble salts of trifluoroacetic acid (121).

The evidence does not allow an unequivocal interpretation of the kinetic data. It is kinetically impossible (122) to distinguish between a mechanism involving the aldehydrol anion as a discrete intermediate and a concerted mechanism involving hydroxyl ion and aldehyde hydrate. For this reason some termolecular mechanisms, compatible with the kinetic evidence, will be discussed.

CF3 CF3

2 OH C-H Mn04 > OH - C +HMn04 (17) / \ / \ HO OH OH OH 127

CF3 CF3

+ _ H20 H-C-O" ^MnC^" —lL> E$p + C~0* + Mn04" I I OH OH

CF3

C—0* + Mu04 =22^ Products (l8) I OH

CF3 CF3

_ = OH" ^"^H- C - OH ~\ti04" -Jl*. H2O + L + Mn04 (19) I / \ OH OH OH

CF3 I fast - C« + Mn04 Products / \ OH OH

The mechanism represented hy equation (17) involves a hydride ion transfer to permanganate aided by the participation of hydroxyl ion.

This reaction is kinetically equivalent to the hydride in abstraction from the aldehydrol anion and is otherwise in agreement with the observed kinetic evidence. It would be difficult to test for 018 transfer from the solvent to aldehyde hydrate due to reasons mentioned above (equation l6). Kinetically, bimolecular mechanisms would be favoured over a termolecular collision.

In the reaction (18) the electron is transferred to permanganate whilst a water molecule acts as a proton abstractor. This mechanism does not require a termolecular solute collision and appears reasonable in terms of the kinetic evidence. However, the absence of an appreci- 128 able solvent isotope effect makes this mechanism unlikely.

In connexion with the oxidation of substituted phenyltri- fluoromethyl carbinols (64,65) where the effect of substituents was slight, it appears plausible that the product of the rate determining step of the type shown below would be greatly stabilised by electron withdrawing groups such as nitro or bromo, and to a much lesser extent

0' Cr

Q>-C_— CF3 -(^)=C CF3 (20) by the methoxy or methyl groups. This idea is supported by the results obtained for the oxidation of benzhydrols. The most striking example being the unusually high rate for 4,4'-dinitrobenzhydrol

(k2 ~ 500)• A small rate, comparable to that for hexafluoro-2-propanol would be predicted for a clean hydride ion transfer on the basis of the electron withdrawing effect of the two nitro groups. It would be difficult to devise an experiment to check for proton or hydride ion transfer. Perhaps solvent effects on the oxidation of dinitrobenzhydrol may show the most pronounced effect.

Equation (19) depicts a mechanism whereby permanganate abstracts an electron from the hydroxyl group of the neutral substrate whilst the proton is being transferred to the hydroxyl ion.

Induced oxidation by hydroxyl radicals may be discounted since the alkalinity is much too low for their production. This process may become noticeable in permanganate solution above 5 M sodium hydroxide where hydroxyl radicals are known to be formed (5*0, however 129 it would be difficult to obtain the kinetic evidence in this region.

The ester mechanism, frequently postulated for chromic acid oxidations (123)> is not; in agreement with the kinetics of the permanganate oxidation of fluoral hydrate in the region pH 6 to lh.

Moreover the analogy is lacking since chromic acid oxidation in the alkaline region is not known.

It seems reasonable to assume that resonance stabilisation of an aliphatic ion-radical intermediate would be much less than that for the aromatic species shown in equation (20). Considering that the above evidence is sufficient to reject the mechanisms 15 and the termolecular mechanisms, we conclude that the mechanism of the aliphatic series involves a clean hydride ion transfer to permanganate.

2. Oxidation Mechanism of Fluoral Hydrate in the Region pH 1 to 6

The reaction of neutral fluoral hydrate with permanganate appears to take place in the region pH 1 - 6 (Figures 15 and l6) how• ever contributions from other reactions are detected at the boundaries.

The rate is essentially pH independent throughout the region and may be expressed by the equation:

Rate = k [ Fluoral Hydrate} [Mn04"}

In addition a deuterium isotope effect of about k still persists

(Table XII), indicating that the C-H bond is broken in the transition state. Chloral hydrate oxidises faster than fluoral hydrate

(Table XV) in this region. 130

Two obvious mechanistic possibilities will be considered.

CF3 CF3 I K I /C-U +Mn04" C-H V0H. (21)

HO OH HO 0Mn03 Permanganate ester

CF3 CF3 I k I

V - H + OH" > ^C^ + H2O + Mn03

HO 0Mn03 HO 0

CF3 - k I Hj + Mn04 ^ C + HMn04~ (22)

HO + OH

Equation (21) depicts the formation of a permanganate ester in a pre- equilibrium step and its subsequent decomposition to trifluoroacetic acid and hypomanganate in a rate determining step. The observed kinetics required for the hydroxyl ion to participate in the transition state.

Rate = k [ester] ^0H~] = kK [ Fluoral Hydrate] [Mn04'] (23)

The small concentration of the hydroxyl ion in this region (lO-8 to

10"13 M) makes the ester mechanism very unlikely. A hydride ion transfer to permanganate is shown in equation (22). This mechanism is in agreement with the rest of the kinetic evidence and the substituent effect.

The triangular transition state represented by equation (15) may also be written for the neutral fluoral hydrate. If the ideas of 131

Lewis and co-workers (119,120) about the triangular intermediate are correct, then this possibility cannot be completely ignored on the basis of the small isotope effect observed. This mechanism also accomodates satisfactorily the kinetic evidence, with exception of the substituent effect. The difficulties in the verification of the

0 transfer have been already disuussed.

Further data such as thermodynamic values and solvent effects are needed before postulation of.more definite conclusions.

3. Oxidation Mechanism of Fluoral Hydrate in the Acidic Region Up to 46$ Sulfuric Acid

Preliminary investigation of the mechanism of permanganate

oxidation was attempted for the more strongly acidic region.

A deuterium isotope effect of 6.5 iH observed for the oxidation of fluoral hydrate in 25 to 46$ sulfuric acid. This value

is normal and may be predicted from the loss of C-H stretching modes

in the transition state. The rate exhibits first order dependence on

the hydrogen ion concentration which can be accommodated into a rate

law of the form

+ V = k [fluoral hydrate] [ Mn04"] [ H J (24) In the strong acid the hydrogen ion concentration must be replaced by

the appropriate ho, h_ or hR (jo) functions. Figure 13 depicts the

relation between the logarithm of the rate constant and the Ho, BL

and HR (Jo) acidity functions; the individual slopes of the straight

lines being 1.7, 1.2 and 0.8 respectively. Since any real correlation 132 must be signified by a slope close to unity, the dependence on the Ho function can be rejected since the slope is too high.

The relation with the HR (JO) function necessarily implies an equilibrium of the following type:

ROH + H""1 N R+ + HaO

This equilibrium is applicable toonly one of the reactants, e.g. fluoral hydrate:

+ CF3CH(OH)2 + H CF3— C—H + H20 (25)

OH

The carboilium ion may further react with the permanganate ion to form an ester which will decompose unimolecularily in the rate determining step.

CF3 CF3 I K I 2 HO — C + + Mn04~ v ^ HO — C - 0Mn03 (26) I I H H permanganate ester

CF3 CF3

1 k 1 HO — C — OMn03 -••> Cs + HMn03 I / ^ H HO 0

The rate is expressed by the following equation:

Rate = k [ester] = kKjK2 [Fluoral Hydrate] [Mn04"]hR (27)

Inclusion of water into the transition state leads to the rate law which implies the h_ dependence:

OH CF3

1 k iv - + CF3 C — OMn03 =• C + H30 + Mn03 (28) K / \ - H HO 0

SH20 133 and the rate becomes:

Rate = k [ ester] [ H2o] = kK [ Fluoral Hydrate] j^Mn04"J h_ (29)

The equations (25) and (26) may be more simply written as

OH

+ CF3CH(OH)2 + Mn04- + H =J=== CF3C -r- OMn03 + HJJO (30) I

H

The equilibrium represented by the equation (30) is kinetically undistinguishable from the following:

+ Mn04" + H ^ • • ^ HMn04 (H- dependence) (31)

OH

CF3CH(0H)2 + HMn04 CF3C— 0Mn03 + H20 (32) I H

permanganate ester

Decomposition of the ester in the rate determining step without the participation of water leads to the result in equation (27), and the inclusion of water in the rate determining step gives the result represented by equation (29); e.g. the logarithm of the rate constant being dependent on the HR or H_ respectively.

A direct hydride ion transfer may be formulated for the oxidation of fluoral hydrate by permanganic acid

CF3 CF3

C—H + ' HMnO* —C + H2Mn04" (33)

HO OH HO OH

Even though this mechanism fits the rate law (equation 2k) it is not in good accord with the observed activation parameters. The energy 154 of activation (12.9 kcal mole-1) is similar to those observed in other permanganate oxidations (17,59,64), however the entropy of activation (- 17.8 e.u.) seems unusually negative for a reaction between two neutral molecules (equations 28 and 35), or for a unimolecular decomposition of the permanganate ester (equation (£6), rate determining step). Cn the other hand a cyclic mechanism, similar to that postulated by Rocek and Krupicka (124) for acid catalysed chromic acid oxidations, accounts well for the observed entropy of activation.

CF3 H C- OH CF3

C Mn C + H3Mn04 (54) / \ / \ / \ HO 0-H 0 0 HO 0

In the absence of data on substituent effects it is difficult to decide whether hydride ion (equations 55,54) or proton transfer accompanied by simultaneous gain of two electrons (equation 26) is taking place. Consideration of other more subtle differences in mechanism are unjustified at present for the same reason.

4. Mechanism of Oxidation of Formic Acid in the Region 20 to 42$ Sulfuric Acid

_ HC00H/Mn04 Reaction.

Oxidation of undissociated formic acid by permanganate occurs in a very narrow region in the vicinity of 20$ sulfuric acid (Figure

20); it was previously detected by Halpern and Taylor (59,60,61) at lower acidities. This reaction is not clearly distinguishable due to 135

the overlap of the hydroxyl ion catalysed HCOO"/Mn04" reaction and

acid catalysed HCOOH/HMn04 oxidations; both rates contribute consider•

ably to the overall rate in the region where HCOOH/Mn04~ reaction is important. Separation of the various rates leads to an isotope effect

of 1.4 for the HCOOH/Mn04~ reaction.

The mechanism of this reaction was discussed previously in some detail' by Halpern and Taylor and will not be expanded further here.

HCOOH/HMn04 Reaction.

A rise in the oxidation rate is observed with an increase in

acidity beyond 20$ sulfuric acid. A plot of log k2 against the H- function depicted in Figure 20, shows clearly a straight line with a slope of 1.0. This immediately disposes of any mechanisms dependent

on the H0 or HR (Jo) functions. Further information is known:

(i) An isotope effect of 3.4 is observed in the region 30 - 42$ H2S04,

(ii) the heat of activation is similar to those found in other permanganate oxidation (AH* = 13 kcal mole-1), (iii) the entropy of activation in large and negative (AS* = - 21 e.u.). Both of the activation parameter values are similar to those for the oxidation of fluoral hydrate in the same region.

The correlation of the logarithm of the rate constant with the H_ function, or kg with h- (Figure 21) imply that the reactive

- species are the neutral formic acid and permanganic acid (or Mn04

and HCO2H2*). It also suggests that the relation of log k2 with 136

HR (JO) for the oxidation of fluoral hydrate is probably fortuitous.

Several mechanisms similar to those described for fluoral hydrate in the same region may be written for formic acid (equations

28,33,34). This will not be discussed further here.

5. Mechanism of Oxidation of Formic Acid on the Region 50 to 72$ Sulfuric Acid

A further increase in the oxidation rate observed with increase in acidity above 42$ sulfuric acid.

This reaction exhibits a considerable isotope effect

(kg/kj) 8) which shows that the C-H bond is broken in the rate determining step.

The activation parameters (estimated from a plot which was curved) indicate a small value for the activation energy (6 kcal mole" and a very large negative entropy of activation (-38 e.u.).

Approximate dependence of the log k2 on the HR (Jo) function is found in this region (Figure 23), however it is difficult to fit the observed data into an appropriate rate law. This may be due to

the fact that HC02H/HMn04 reaction still contributes greatly in the region where the increase in the oxidation rate occurs (~ 42$ H2SO4).

It is not possible at present to separate the two rates contributing

to the. overall rate, since the rate due to HC02H.HMn04 reaction can not be determined independently. A tentative mechanism which is in qualitative agreement with the above observations involves a reaction

+ between Mn03 and formic acid. A hydride ion transfer (similar to 137 equations 33,34) or a formation of the permanganate ester and its subsequent decomposition in the rate determining step (similar to equation 26) may be formulated. Both of these possibilities must be regarded as tentative until more evidence is accumulated about this reaction.

6. Summary

The permanganate oxidation of fluoral hydrate (formic acid) in such a great range of acidic and basic solutions has enabled the reactive species to be identified and has led to the postulation of reasonable reaction mechanisms.

The interesting features of the permanganate oxidation are the variation of the isotope effects throughout the investigated region and the great variation of rates for the corresponding mechanisms.

The increase in rate of oxidation in the progressively more alkaline region is due to the change of the nature of the oxidisable substrate; e.g. the rate increases for the series fluoral hydrate < fluoral hydrate mono-anion < fluoral hydrate di-anion, and the rise may be related to the ease of hydride ion abstraction.

On the other hand in the strongly acidic solutions, the VII substrate's identity remains the same whereas the Mn species change. The expected increasing electrophilic character for the

+ sequence Mn03 > HMnO* > Mn04" is reflected in the corresponding rise in the rates of oxidation by these species. 138

11 + If the role of the various Mn^ species, Mn04", HMnO*, Mn03 , has been correctly assigned, it follows that permanganate ion, MnO*", is the weakest oxidising agent of the three.

A parallel between the oxidation by chromic acid and permanganate appears to exist in the acid region; however such a similarity is completely lacking in the alValine region.

The unusual isotope effect kH/kD = 20 for the oxidation of hexafluoro-2-propanol deserves some attention. Its magnitude is in essential agreement with a transition state involving a complete loss of stretching and bending modes (106,107). The examinations of the

IR spectra reveals that the CH, CD stretching vibration (2960 cm-1 and 2170 cm-1 respectively) and bending modes (1377 cm-1 and 1015 cm-1 respectively) are somewhat higher than usual, however the ratio

VCH/ VCD = 1.36 is normal. Perhaps some specific influence of the trifluoromethyl groups might be responsible for this high isotope effect.

Of interest might be Swain's ideas (125) for distinguishing between a proton or a hydride ion transfer, however more data would be required before a full use could be made of his proposals. 139 APPENDIX

N.M.R. Study of Hydroxyl Proton Spin-Spin Splitting

INTRODUCTION

It should be pointed out at the outset, that the coverage of this topic is not intended to be comprehensive or complete. The primary purpose of this study was to develop a method useful to the organic chemist in elucidating the structure of molecules, although it is impossible to escape and separate entirely the theoretical consequences, which have more far reaching effect than described herein. For introductory and more advanced theoretical treatment of the nuclear magnetic phonomenon the reader is referred to, what are now recognised, standard texts on the subject (126,127,128,129,130).

It has been pointed out that the slope and width of nuclear magnetic resonance (absorption) lines are very sensitive to time-dependent processes (ref. 127, chapters 5 and 10). These may vary from intramole• cular exchange and internal rotation to exchange of nuclei between two different chemical positions. Many such processes are known, but pertaining to present discussion a suitable example may be sought in the proton exchange, in a system such as phenol-water or alcohol-acid.

In pure dry ethyl alcohol the proton of the OH group gives rise to a triplet signal due to the spin-spin interaction with the protons of the adjacent methylene groups. When however, small amount of hydrochloric acid is added the OH triplet collapses to single sharp iko signal. This is interpreted in terms of rapid exchange of hydroxyl group proton between neighbouring molecules or perhaps more appropriately in terms of protonation of hydroxyl group and subsequent proton exchange.

If the exchange is slow the spectrum will show two distinct signals, one corresponding to the water and other to the hydroxyl group. If the exchange however is rapid, then only one signal will appear in the averaged position. Broadening of signal will result as a consequence of time averaging processes. The detection of the OH spin-spin coupling multiplets eluded the investigators for a long time, and exchange processes such as mentioned above were blamed as the major cause (131). However careful purification of sample can reduce the rate of exchange and spin-spin final structure of the OH group may be detected, as for example in ammonia (132) and ethanol (133,134).

Zimmerman and Weinberg (133) noted that presence of water had marked effect on the chemical shift of the OH - H2O coalesced line.

Interestingly enough, the spin-spin splitting was still observable when small amounts of water were added.

Sometimes it is possible to retard the exchange process bf addition of a complexing agent, which forms preferential stable molecular complex with the sample. This method was used successfully by Zimmerman and collaborators (135) who have studied the solution of methanol in acetone. The hydrogen bond which is presumably formed between the methanol and acetone increases the life time of the OH protons in enough molecules to reveal fine structure. Additional experiments with methanol show that the internal chemical shift can 141 be changed by varying the acetone concentration. As a consequence alteration in fine structure results.

MacLean and Mackor have used low temperature as means for retarding the proton exchange (136)• Ethanol was protonated by a mixture of HF-BF3 at - 75°' The protonated ethanol shows a predicted multiplet structure which is identified with the species CH3CH20"H2.

Hydrogen bonding appears to play an important role in the detection of the OH spin-spin coupling, and for this purpose few references will be mentioned here, pertaining to the present investiga• tion. Pimentel and McLellan book "The Hydrogen Bond" may be con• sidered as a general reference (137)•

Shift of the hydroxyl group proton resonance of ethanol and methanol was studied as a function of concentration (138,139,140,141); carbon tetrachloride served as the diluent. The shift of the hydroxyl resonance is interpreted in terms of equilibria between hydrogen bonded trimers, tetramers (138,140) or dimers (139,141) of alcohol molecules with the corresponding monomers. Most investigators agree that the monomeric species exist in dilute solutions; the IR and viscosity data are In support of this assumption. It is suggested that hydrogen bonding is a normal dispersion force initiated by the dipole-dipole interaction and larger than most dispersion forces due to the high polarisability of the electrons on the oxygen (138).

The behaviour of a symmetrical hydrogen bond in the hydrogen maleate ion was studied by Forsen (142). Dimethylsulfoxide was used as a solvent since there- is presumably no possibility of interfering 142 proton exchange between solvent and solute. Strong intramolecular hydrogen bond persisted in this solvent as judged by large shifts which were observed for the hydrogen maleate and hydrogen phthallate ions. In contrast undissociated maleic and fumaric acids showed only small shifts and gave similar spectra. On the basis of this observa• tion it was concluded that the intramolecular hydrogen bond is broken up in maleic and in this solvent.

In connexion with the structural determination of fluoral hydrate by n.m.r. a closely related study was made of the equilibrium between acetaldehyde and its hydrated form (l4-3). Changes in spectra were observed with subsequent additions of water or deuterium oxide.

Spectrum corresponding to the aldehyde was clearly distinguishable from that of the aldehyde hydrate. The position of the peaks remained essentially constant only intensities changed depending on the relative amounts of water and aldehyde. Equilibrium constant calculated from the n.m.r. data agreed very well with previous values derived by other methods (l44).

With the aid of theoretical approach, developed by Karplns

(14-5), Lemienx and co-workers were able to analyse rather complex spectra of cyclohexane derivatives and arrive at this conformational structure (146). The variation of the Cw-H splitting in a number of compounds was measured by Lauterbur (ikj). Somewhat more detailed study on formyl type compounds (148) and halomethanes (14-9) led to the interpretation of the increase of the coupling constant JcH with the increasing electronegativity of-the substituents attached to the formyl 143 group, in terms of progressive change in orbital hybridisation. The magnitude of the coupling constant ranges from 173 cps for CH3CHO to

267 cps for PCHO.

Much smaller coupling constants were obtained by Narashiman and Rogers (15O) and Brown, Dickerhof and Bafus (151) for the coupling between the methyl and methylene protons in metal alkyls of the type

M(CH2CH3)X. Ikk

EXPERIMENTAL

1. Technique of n.m.r. Measurements

The experimental procedure appears to play, an important role in the observation of the OH spin-spin coupling and therefore it will be described here in some detail.

Selection of n.m.r. Tubes.

Standard n.m.r. tubes (outside diameter = 5.0 mm,) made of thin glass which are supplied with the A-60 Varian Associates instru• ments were used. Tubes which were not previously used for n.m.r. work were selected, washed with reagent grade acetone and dried in the oven at about 100°. They were provided with a specially made polyethylene cap fitting tightly over the rim of the tube. This arrangement pre• vented any evaporation of the solvent as well as enabled thorough shaking and mixing of components. The specially, selected tubes were always washed with reagent grade acetone only - never with chromic acid or soap.

In case of liquid samples, then were run, whenever possible, as neat liquids first, spectrum was recorded and then a particular solvent was added until OH splitting was observed. If an unequivocal identification of the OH absorption could be made, no other tests were required. However in many instances the OH absorption was obscured ' by overlap of another absorption due to other protons in the molecule, and the OH proton, whether split or not, could not be identified. In 145 this case the following procedure was adopted: The compound under study was first dissolved in one particular solvent, for which the approximate chemical shift of the OH was known. Another solvent was then added dropwise and the changes in the spectra noted. Since the

OH proton suffered greatest chemical shift as compared to C-H protons, the movement of the OH could he followed up to the point where it was sufficiently separated from other interfering absorptions. There it could be easily identified either by its splitting or by collapsing it with the addition of small amount of base - usually 2 jul of O.O5 M

NaOH. Sometimes it was necessary to allow the sample to stand for several hours before the splitting -was observed.

All proton spectra were obtained at 60 Mc on the A-60 Varian

Associates instruments and the F18 spectra at kO Mc on the VkjQO Varian

Associates spectrometer.

Solvents.

Following solvents were used in this investigation: acetone, dimethylsulfoxide (DMS), carbon tetrachloride, water, deuterium oxide, hexadeuterio, acetone, deuterio chlorform and tetrachloroethylene.

All solvents were of reagent grade and were checked for purity on the vapour phase chromatography. They were further purified by distillation and stored in dark bottles over anhydrous magnesium sulphate and small amount of anhydrous sodium bicarbonate. Hexadeuterio acetone was kindly donated by Dr. D. E. McGreer and deuterio chloroform was a gift by Dr. J. P. Kutney. 146

Tetramethylsilane (TMS) was used as the reference sample, in

some cases as an internal, and in others as an external standard. No

"bulk susceptibility correction was made when it was used as an

external reference and for this reason caution should be taken when

interpreting the values of chemical shift.

Most of the compounds used in this investigation were purchased

commercially and purified either by recrystallisation (solids) or by distillation or vapour phase chromatography (liquids).

The sample of diacetone glucose and diacetone galactose were kindly supplied by Dr. G. G. S. Dutton of this department.

Samples of phenyl-trifluoromethyl-carbinols were kindly provided by

Mr. D. G. Lee of this laboratory.

The l,l,l;3,3,3-hexachloro-2-propanol was prepared by method

of Geiger and Usteri and Graenacher (152) in 98$ yield. Recrystallisa•

tion was afforded from petroleum ether (30 - 60°).

Synthesis of other compounds used in the n.m.r. study is

described in the main body of the thesis. RESULTS

1. Structure of Fluoral Hydrate

Fluoral hydrate was prepared as described in the main body of the thesis. However, the melting point, molecular weight and carbon-hydrogen analysis were not in accord with other investigators

(85,89,90,9l)• Infra-red spectrum of the compound prepared in this laboratory (as well as sample purchased from Columbia Company) was however reasonably identical with that obtained previously by others.

In order to elucidate the structure of this compound, it was thought that n.m.r. might yield some additional information. This proved to be the right approach which led eventually to the assignment for the structure of this compound. The n.m.r. spectra of several possible structures of fluoral hydrate were predicted.

The monomeric fluoral hydrate, CF3CH(0H)2 would be expected to give rise to two peaks in the H1 region of intensity 2:1, correspond• ing to two OH protons and one CH proton. In the absence of any CH-OH coupling the CH proton should be split into a quartet due to the interaction with adjacent fluorines. The F19 spectrum should show a doublet, due to the coupling with CH proton. ^-H

The condensation dimer Hv y0 C

/ \ 0 CF3

CF3 OH \ H should also give two peaks in the H1 region, but now these should be Ik8 in a ratio 1:1 corresponding to the two OH and two CH protons. Again

CH protons should he split into a quartet. The F19 spectrum would be

expected to be almost identical with that of the monomer.

For the cyclic trimer (trioxane type) two or possibly more

isomers might be observed. The significant difference, as compared

to the other two possibilities, is the fact that no OH absorption

should be observed in the H1 regions. Similarly it is possible to

predict the spectrum of the linear trimer but this will not be shown here.

Attempts were made to run the spectra in deuterio chloroform,

however the compound was not readily soluble in this solvent. The

proton n.m.r. spectrum of fluoral hydrate in hexadeuterio acetone

(external TMS reference) showed two peaks, one of them a broad

doublet at a low field and the other a broad singlet, at a higher

field, showing some indication of fine structure (Figure 2k). The

areas of both peaks were approximately equal. After standing over•

night in a stoppered tube, the proton resonance spectrum of the sample

was run again, and a considerable sharpening of the spectrum was noted.

A distinct doublet at 6.56 Q (J = 8.5 - 0.2 cps) and a sextet at

5.1 £ (J = k.l i 0.2 cps) can now be clearly recognised (Figure 25).

In addition another small peak at 5.87 <£ can also be sean. Addition

of 10 ^tl of water causes collapse of the doublets into a broad resonance

and the sextet is changed into a quartet (J = k.l - 0.2 cps) (Figure

26). Also other small peaks have appeared at a slightly higher field

to the quartet. On the basis of broadening of the resonance line, Ik9 upon addition of water, the doublet at 6.56 <£~ is identified as being due to the hydroxyl. protons.

The F19 spectrum revealed one large doublet at low field

(J = 4.3 - 0.2 cps) and another doublet of much lower intensity at slightly higher field (Figure 27).

From the point of view of the oxidation kinetics, it was necessary to establish the identity of the species in water solutions.

This was done by adding successive small amounts of water to the sample, (dissolved in de acetone) and observing changes of the proton n.m.r. spectra. Addition of 10 ul of water was already shown in

Figure 26. Further addition of water (39 pi) decreases the intensity of the quartet of this dimer, and another quartet is seen growing at higher field (Figure 28). Clearly the addition of water causes hydrolysis of the dimer into the monomer of fluoral hydrate. Thus it may be safely assumed that the species present in water solution is entirely in the form of the monomer. On the basis of the above results the spectrum of fluoral hydrate is identified as being due to the con• densation dimer

CF3 / H\ 0 C H A / \

H

The splitting of the OH resonance (Figure 25, 6.56 *t> ) into a doublet

151

FIG. 25 H1 Spectrum of Fluoral Hydrate after Standing

7 6 ^ 5 SppwN

H0 152

FIG. 26 H1 Spectrum of Fluoral Hydrate

(+10 yul H2Q)

_ r- 1 7=—

7 6 5 ^PPm FIG. 27 FI

The absorptions at 5.1 £ (Figure 25) are due to the CH protons; the sextet arising from a fortuitous relation of the CF-CH and CH-OH coupling constants, e.g. JCF-CH : JCH-OH = 4.3 : 8.5 (^1:2). The identity of the CH resonance is further supported by the fact that water addition which removes the CH-OH interaction, reduces the CH resonance to the expected quartet (Figure 26). The F19 spectrum

(Figure 27) shows a doublet as would be expected from the interaction of the Fis with the adjacent CH proton. The doublets of small intensity in the F19 and H1 spectra (Figures 2J and 25 respectively) are assumed to be due to the monomeric form of fluoral hydrate, which is formed by hydrolysis of the dimer due to the traces of water present in the sample.

After a lapse of two months the H1 and F19 spectra were

repeated. Sample was dissolved in the usual manner in d6-acetone and the proton spectrum recorded (Figure 29). The OH resonance shows only single peak and the CH issplit into the expected quartet. Two other additional quartets appear at lower and higher field of the CH resonance. The peaks at higher field were previously ascribed to the monomer and the quartet at lower field was assumed to be the trimer.

The absence of the corresponding proton resonance indicates that it is the cyclic trimer. Fluorine spectrum which was run at the same time is shown in Figure 30. Five doublets are clearly visible. The strongest doublet is due to the dimer and the one at highest field to the monomer. The doublet at lower field to the dimer is presumed 155

FIG. 28 H1 Spectrum of Fluoral Hydrate

(39 yu| H2Q)

76 I5 Spprr, ^

157

FIG. 30 * F|c? Spectrum of Fluoral Hydrate 158 due to the cyclic trimer . The remaining two doublets at lowest field are assumed to be due to some polymeric form, obviously present in negligible amounts. From the point of view of oxidation kinetics, the present results are immaterial, since all kinetic solutions were made up from the freshly prepared fluoral hydrate dimer, which con• tained no other species. It is apparent that polymerisation result upon prolonged standing, which is reasonable, since the sample was kept in a closed container over P2O5.

Proton spectra resulting from successive additions of water are shown in Figure 31 (+ 30 jul H2O) and 33 (+40 jul H2O) and correspond• ing changes in fluorine spectra in Figures 32 and 34. The disappear• ance of the dimer and trimer and appearance of the monomer may be clearly recognised. The spectrum after addition of 160 ul of water . is shown in Figure 35 &nd the corresponding fluorine spectrum in

Figure 36. The dimer and trimer are now almost completely hydrolysed to the monomer. This leaves no reasonable doubt that in water solution, the monomeric form is the only species present.

Another sample which was made up in de-acetone in a manner identical to that above was allowed to stand for 10 days. Originally the spectrum was the same as shown in Figure 29 but after 10 days the

OH and CH showed fine structure (Figure 37)• Moreover the OH of the monomer at 5.83 & is now clearly distinguishable. It seems reasonable to assume than any water present in the solvent, absorbed by sample during handling, or standing would hydrolyse the trimer or dimer into the monomer. It also explains why the OH-spin-spin coupling is not 159

FIG. 31 H1 Spectrum of Fluoral Hydrate

03 0 ul H2Q) FIG' 32 Fn Spectrum of Fluoral

Hydrate (t30 /Jl H20)

H ON O FIG. 33 . H' Spectrum of Fl. Hydr

(•40/JI H20)

H2Q

>p p rv\ H, FIG. 34 F11 Spectrum of Fluoral Hydrate

+ ( 40yul H2Q)

50 cps

»• FIG. 35 H1 Spectrum of Fl. Hydn

0160 ul H20) l6k

FIG. 36 F^Spectrum of Fluoral Hydrate

(+160 jul H2Q)

50 cps

166 observed in freshly prepared n.m.r. sample. Fast proton exchange would be expected since the pKa difference of water and fluoral hydrate

(monomer) is about k pKa units. Moreover, the dimer should be more acidic than the monomer. In fact OH of the dimer appears at lower field than the OH of monomer. If however all the water is used up in the hydrolysis of the dimer, than OH splitting should appear - and this is actually the case. This also points out that the hydrolysis is a rate controlled process - this is again confirmed by the fact that the dimer disappears more quickly than the trimer (Figures 32 and 3*0.

A sample of fluoral hydrate obtained commercially was also examined by n.m.r. and found to contain mainly the dimer. Sample was a liquid. Addition of water showed clearly increase in the monomer and decrease in the dimer concentration.

It is now interesting to compare the available evidence on fluoral hydrate. Henne and co-workers (89) report that the hydrate sublimes at about 50°, whereas at least two groups of investigators find that the melting point is 68 - 70° (85,91). Shechter and Conrad report that fluoral dissolves very slowly in water but readily in dilute acids (90), whereas others note that it is readily soluble.

Storage of fluoral results in polymer formation which hydrolyses readily in dilute carbonate solutions but slowly in concentrated acids to give fluoral hydrate (90). The hydrate is readily hydrojkysed by alkali (91). 167 Observations made in the present investigation point out that fluoral reacts rapidly with water even at - 78% in fact at room temperature the reaction is so strongly exothermic that it is almost violent. The attempts to determine the melting point on an standard

Fisher type apparatus failed. Almost any melting point could be obtained, depending only how warm was the stage at the start. Determination of melting point was finally made in sealed tube and both, the protio and deuterio compounds melted at essentially the same temperature. However this melting point was about 35° higher than that reported previously.

Repeated melting and cooling of the samples did not alter the melting point. The alkaline hydrolysis of fluoral hydrate (monomer) was checked in the kinetic experiments and it was found that at room temper• ature, (ionic strength ip.^1, O.15 M NaOH) the hydrolysis is negligible for at least 1200 sec. The IR of the commercial sample of fluoral hydrate (claimed to be monomer) is identical with that reported by

Husted and Albrecht (91), except for one peak which is due to impurity.

The proton and fluorine n.m.r. spectrum show later that the commercial

sample consists mostly of the dimer. The impurity is also seen on the n.m.r. spectrum.Mslecular weight, C,H analysis, melting point, IR and n.m.r. of the protio and deuterio compounds prepared in this labora•

tory show conclusively the dimer structure.

The deuterio compound (fluoral hydrate dimer) showed only one peak in the proton region of n.m.r., identified as the hydroxyl group

(Figure 38). Position of the peak was about the same as for the protio dimer in accordance with expected acidity. Appearance of a

singlet is expected due to the quadrupole moment of deuterium which 168 will result in broadening of any possible splitting. The concentration of the deuterio compound was 197 mgm/0.3 ml. acetone (protio 140 mgm per 0.3 ml acetone) and the peak areas, corrected for weight, were same - for both protio and deuterio compounds (aera deuterio/area protio = O.95.). This identifies the deuterium compound also as the dimer, even though the evidence here is not so complete as for the protio compound.

2. Structure of Chloral Hydrate

Structure of this componnd is, of course, very well known, but it makes an interesting comparison with structure of fluoral hydrate.

Compound was sufficiently soluble in deuterio chloroform, which is in contrast to fluoral hydrate. The n.m.r. spectrum is very

simple, showing only two peaks (Figure 39), the smaller peak at 5.5£

is easily identified as the CH proton and the greater peak at 4.87 £

is the OH resonance. The ratio of the 0H:CH is 2-7:1- Small amount of water which is probably present in deuterio chloroform causes the

OH absorption to be greater. Theoretical ratio should be 2:1.

Spectrum in acetone shows a doublet at 6.02 £ with coupling constant

of 6.7 t 0.2 cps. A triplet at higher field (5.04 ) has a coupling

constant 6.6 ± 0.2cps (Figure 40). Integration of areas shows ratio

of doublet to triplet as 5.8:3 respectively. The spectrum shows

clearly that the structure of chloral hydrate is CCl3 OH

Hx OH FIG. 38 H' Spectrum of 170

F!G. 39 171

FIG. 40 H1 Spectrum of Chloral Hydrate Acetone

I V \ '6

H0 172

Addition of water (which removes the CH-OH interaction) causes broadening of the OH resonance and the CH absorption collapses into a

singlet (Figure 41).

3- Investigation of OH Spin-Spin Coupling in Alcohols

The encouraging success with fluoral and chloral hydrates, where it was possible to observe the OH fine structure, precipitated

further study of the OH spin-spin splitting in alcohols. It was hoped that this would lead to development of a method for identification of primary and secondary hydroxyl groups in a molecule. Some of the examples chosen are of the simple type, AB or AX systems - such as benzhydrol and hexachloro-2-propanol, however others are considerably more complicated ABXn type systems, e.g. isopropyl alcohol. Some samples were run neat (pure liquids) or in various solvents. The OH resonance was identified in the usual way by collapsing it with addition of base, and observing the changes in the CH proton resonance.

The OH coupling constants and other data of compounds investigated are shown in Tables XXV and XXVI, presented later in the text.

In most of the spectra, considerable dilution with the solvent was required before the OH splitting was observed. The following spectra are shown: isopropyl alcohol (Figure 42)j l,l,lj3,3,3-hexa- fluoro-2-propanol (Figure 43, OH splitting absent, Figure 44, OH splitting present), phenyltrifluoromethyl carbinol (Figure 45). In the case of hexafluoro-2-propanol considerable simplification of the

CH spectrum results due to the fortuitous relation between the JCF-CH 173

F IG. 41 FIG.42 H' Spectrum cf Isopropyl Alcohol 175

FIG 43 H' Spectrum of CF3CHOHCF in Acetone FIG. 44 H" Spectrum of CF3CHOHCF3 in Acetone-DMS FIG. 45 H' Spectrum of PhCHOHCF in DMS

Ho 178 and JcH-OH coupling constants (Figure kk) (cf. fluoral hydrate dimer,

Figure 25). k. Solvent Influence on the Hydroxyl Resonance

In some instances the OH proton resonance overlaps the CH absorption, in which case no spin-spin splitting is observed, or it occurs too close to the CH resonance so that the system is of the AB type and then theoretical calculation is needed to predict the number of lines and intensities. It was found advantageous in some cases to dissolve sample in one solvent and then shift the hydroxyl proton into a suitable position where the splitting could be observed. This procedure finds greatest application in the study of natural products which will be discussed later, where the spectrum is rather complicated and it is not always possible to observe the OH resonance (split or not).

In order to find out which way the OH proton shifts in a particular solvent, a relatively simple system of isopropyl alcohol was chosen. The spectrum of isopropyl aloohol is well known and will not be reproduced here. For purposes of comparison the various

absorptions are listed below: OH at 5.11 £ , CH at 3.9k % and CH3 at

1.15 £ (neat), all with respect to the internal TMS reference. A particular solvent was then added and the OH shift was noted after each addition. Final position of the OH, corresponding to high dilution is reported below: in acetone 3.38 £ (split), in 179 dimethylsulfoxide 4.32 % (split), in tetracbloroethylene 1.63^ (not split), all & values with respect to the internal TMS reference.

Thus the relative shifts of the OH from its initial position in pure sample to the final position at almost infinite dilution are as follows: dimethylsulf oxide 0.79 £> , acetone 1.73 ^ and tetrachloro- ethylene 3.48 £> ; all solvent shifted the OH resonance to higher field.

It is reasonably obvious from these results, that sample dissolved in tetrachloroethylene will have 0** appearing at high field. This OH resonance can now be shifted to low field with addition of acetone or dimethylsulf oxide, to a position where it is split, or where it can be identified. Application of this method will be shown in connexion with structure of natural products.

5. Study of Hydroxyl Proton Splitting in 2,2,2-trifluoroethanol

Isopropyl alcohol was chosen in the previous study of solvent effects as a typical example of a secondary alcohol. Here, trifluoro• ethanol was chosen as an example of a.primary alcohol. Some interest• ing results were obtained which are discussed below.

Sample was run first as a pure liquid, and then various solsrents were added and the shift of the OH resonance observed. The spectrum of 'the pure liquid showed OH singlet at 5.10 and CH quartet at

3.93 £ f with respect to internal TMS reference. The coupling constant of the quartet JcH-CF = 8.9 ± 0.2 cps. Addition of dimethylsulf oxide produced shift of the OH to the lower field. This is surprising in comparison to the isopropyl alcohol, where the OH 180 was shifted to a higher field by this solvent. After addition of large amount of dimethylsulf oxide, the OH resonance was found at

5.98 £ and was split into a triplet with J = 6.5 cps (Figure 46).

Similarly the CH resonance shows eight clearly distinguishable lines with same splitting J = 6.5 cps. The coupling between the two peaks at highest field is somewhat higher (7.8 cps) and also the intensity is higher than expected. This is caused by unfortunate interference of a side bond of the solvent which may be seen growing in that position before the CH resonance splits. Otherwise the spectrum is in excellent accord with that predicted theoretically - AM2X3 system.

Carbon tetrachloride was tried as the next solvent, however trifluoroethanol is virtually insoluble in this solvent. For this reason, solutions of carbon tetrachloride saturated with trifluoro• ethanol were used. At this dilution (essentially infinite) the OH resonance was now shifted to higher field (2.98^ ), with respect to internal TMS reference. The OH is again split into a triplet, but now the coupling constant is J = 7-2 - 0.2 cps (Figure 4-7). This is different from the dimethylsulfoxide case where the coupling constant was J = 6*5. cps. Indeed this is no mistake, since also the CH resonance shows essentially five lines with only barely discernible fine structure. Moreover the CH coupling constant is also 7.2 -

0.2 cps. This is certainly no coincidence, in fact it is obvious that the trifluoroethanol species in carbon tetrachloride is different to that in dimethylsulfoxide. An explanation is needed in order to account for this phenomenon. FIG. 46 H' Spectrum of Tr if Iuoroethano in DMS overlapping s id e ban d

ITN PP H CD H, FIG. 47 H' Spectrum of Trifluoroethanol in CCI4

H0 It is generally accepted that the alcohol species exist at infinite dilution in the monomeric form (138,139,140). Thus the assumption can be made, that in carbon tetrachloride, the trifluoro ethanol molecules exist as monomers. It is very doubtful that any extensive hydrogen bonding will occur between this solvent and the alcohol. On the other hand, the different OH coupling constant, which is observed in dimethylsulfoxide is consistent with the tri• fluoroethanol species strongly hydrogen bonded to the solvent (cf. ref. 135)• This assumption is reasonable, since dimethylsulfoxide is a much stronger base than trifluoroethanol.

The spectrum of trifluoroethanol in tetrachloroethylene is identical with that in carbon tetrachloride. The only exception is that the OH absorption occurs at 2.4-3 & (internal TMS reference).

In all cases the CH resonance remains essentially constant within

+ 0.02

6. Application of OH Spin-Spin Splitting to Natural Products and Derivatives.

A relatively simple case was chosen, in order to demon• strate that it is possible to observe the OH spin-spin splitting even in more complicated structures.

Diacetone-glucose (l,2j5,6-diisopropylidene»D*glucose) was selected since it contains only one secondary hydroxyl, which should be identifiable as a doublet. Spectrum was first obtained in acetone (Fig. 48) and even though it appeared that one of the several doublets which were observed might be the hydroxyl, no 18U

FIG. 48 H1 Spectrum of Di acetone-glucose (Ring protons) 185

clear-cut conclusion could be reached. Addition of 49 ul. of water

shifted the hydroxyl peak into a position where it could be iden•

tified (Pig. 49) as a doublet. Addition of 2 yul. of 0.05 M sodium

, hydroxide collapsed the hydroxyl resonance completly and the same

time, corresponding change in the CH resonance was observed, which

enabled the identification of the proton coupled to the OH group

proton. This spectrum is not shown here. The OH splitting was also

clearly observed in tetrachloroethylene (OH appearing at high field

was well separated from all other resonances) and the shift of OH to

lower field was produced by addition of dimethylsulfoxidej this spectrum

is not shown here.

As another example, citronellol was chosen as a representa•

tive of a group of naturally occuring compounds containing primary

hydroxyl groups. The spectrum of the pure liquid is shown in Fig. 50.

Assignements of various peaks appear directly in the illustration.

At higher resolution it is possible to see the fine structure of the

vinyl proton resonance (triplet) caused probably by secondary coupling

with the trans methyl group. Similarly, one of the low field methyl

groups shows coupling to the vinyl proton of about 1 cps (not shown

in Fig. 50). The OH resonance in pure liquid is not split. Addition

of dimethylsulfoxide however, splits the OH proton into a clearly

identifiable triplet and the corresponding CH2 resonance appears as

a quartet (Fig. 51). The quartet formation is again explained on the

basis of fortuitous combination of the JCH2-CH2 couPlinS constants.

These two examples indicate clearly that it is possible to 186

FIG. 49 t H1 Spectrum of Diacetone-g 1 ucose (Ring protons and OH) Hp

OH 1 If 1 [ 1 r Nr

H0 FIG. 50 < Hl Spectrum of Citronellol (neat)

CH3

w.—-

189 detect the hydroxyl groups in more complicated structures and also

identify such hydroxyl groups as primary or secondary.

7. Conformational Analysis of Diacetone-galactose.

The spectrum of diacetone-galactose (l,2;3,4-diisopropylidene-

D-galactose) was attempted in tetrachloroethylene, with the pri• mary purpose in mind of identifying the primary hydroxyl which

is present in this molecule. This has not been accomplished,

however it was possible to identify all other protons in the mole•

cule and establish the conformational structure.

The structure of diacetone-galactose is represented below

in the conventional Haworth's formulation.

1

The individual hydrogens and methyl groups are labeled in this for•

mula. The n.m.r. spectra are shown in Figures 52 and 55.

Analysis of Spectra.

Assuming that the spectrum is not unnecessarilly complicated,

then essentially only two clearly defined doublets should be observable. 190

FIG. 52 H1 Spectrum of Diacetone-ga lactose (Ring protons) C r •>

A B

V

4 6 ppn^

H 191

FIG. 53 H Specrum of Diacetone galactose (Ring protons) E 192

One for the and the other for the C6 protons. One doublet (J =

4.9 cps), marked A in Fig. 52, appearing at low field is identified with reasonable certainty as the (l46), which is presumably the most acidic proton in the molecule (excluding the OH proton). The absorption B is a doublet with large J = 8.0 cps, each peak being

split into a further doublet having small J = 2.1 cps. Clearly not H2.

The absorption C has an appearance of a distorted triplet, having fine structure, however the area corresponds to two protons. It seems that the apparent triplet is in fact composed of two overlapping doublets, one having large J = 8.4 cps (small J = 1.6 cps) and the other having large J = 5.0 cps and small J = 2.0 cps. Thus now a

temtative assignement can be made that B is H3 or H4 and C contains

H2 and H3 or Ei. It is difficult to decide whether H3 or H4 should appear at low field, however this is of no consequence in the con•

formational analysis. Having tentatively identified H2, H3 and

H4 it now remains to establish the position of H5 and the methylene

group at C6. The H5 should be split by H4 and C6 protons. The methylene group should show a doublet. A closer examination of

Fig. 53 reveals presence of a distorted triplet (D) and distorted doublet (E). The doublet has a coupling constant J = 5.5 cps and is identified as the methylene group. The triplet with large J= 5.0

cps and small J- 2.0 cps is assumed to be the H5. The doublet is distorted more than expected, which might indicate that the OH group is overlapping the methylene group. The individual coupling 193 constants may be summarised as follows:

Jl-2 = 5 cps

J2_2 = 2 cps

8 C S J3>4 = P

2 C S Jk-5 = P

J 5-C6= 5 cps.

Using now the data of Karplus (153) and noting that the only unusually large coupling constant is J3_4 = 8 cps, it follows that the angle between H3 and H4 should be in the order of zero or

180°. By the same considerations, the angle between H2 and H3 or

H4 and H5 should be in the vicinity of either 60 or 120°, and the angle between H1-H2 about kO to 140°. Four possible structures were considered; chair and a boat form with the six-membered ring oxygen down and a chair and a boat form with the oxygen up (see formula below). Both boat forms were rejected since they did not fit the data. On the other hand both chair forms fitted the data equally well, and the decision between them was not clear-cut.

However additional information may be derived from the position of resonance of the methyl groups of the acetone residues.

Three peaks may be clearly recognised at high field (not shown here), corresponding to two equivalent and two non equivalent methyl groups.

Non equivalence may be caused by a combination of a steric and spatial electronic interaction of the methyl groups with the oxy• gen atoms in the molecule. For this reason, the chair form with the oxygen down is favoured over the other form, since it apparently 194 fits the data better. The conformational structure is shown below

U CH-a \ X

It is now apparent that the methyl group I can interact stericaly as well as electronicaly with the hydroxyl group which will be orien• ted in the direction of the two oxygens (five membered ring) and hydrogen bonded. Also the methyl group III would be able to inter• act with the six-membered ring oxygen. It is also apparent that methyl groups II and IV should be equivalent since they cannot interact with the ring. The difficulty with which the OH resonance could be shifted with addition of small amounts of DMS, H2O or acetone, might imply that the OH Is hydrogen bonded internally - e.g. to the two oxygens of the five membered ring. This would not be possible in the chair conformation with the six-membered ring oxygen up. The evidence points out that diacetone-galactose exists in tetrachloroethylene in the chair form as shown above. This structure is also consistent with the fact that intramolecular hydrogen bonding would be expected to increase in an inert solvent, such as tetrachloroethylene. It is also expected that in polar 195 solvents, such as acetone or dimethylsulfoxide , the chair form with the oxygen up will he preferred, since now the hydroxyl group will have an opportunity to hydrogen bond directly with the solvent, rather than internally.

Studies are in progress on the conformation of diacetone- glucose in tetrachloroethylene, acetone and dimethylsulfoxide. Pre• liminary results indicate that all four methyl groups are non equi• valent In tetrachloroethylene, whereas in acetone three methyl groups are equivalent and one appears at a lower field than the other three.

This again indicates that diacetone-glucose exists in different confor• mations in these solvents.

8. Effect of Substituents upon the Magnitude of the OH coupling.

In recent years, several studies have been made of the influence of substituents upon the magnitude of the C1^-H spin-spin coupling (lV7,l48,l49), however, it appears that a comparable study of the OH spin-spin splitting constants has not been made up to the present.

The coupling constants of some primary and secondary alcohols are summarised in Tables XXV and XXVI. No quantitative correlation of the data will be attempted, since the investigation was not extensive enough.

Upon examination of the Tables, several trends become appa• rent: (i) The magnitude of the coupling constants of the aliphatic compounds increases with increasing electronegativity of the substi- 196 tuent, (ii) The coupling constants of the aromatic series (phenyl- trifluormethylcarbinols) appear to he relatively insensitive to the phenyl ring substituents. There is no apparent reason for the diffe• rent effect of the substituents in the aromatic and aliphatic compounds.

One possible explanation might be seen in terms of the spatial electronic interaction of the substituent with the hydroxyl proton in the aliphatic compounds; this would not be possible in the aromatic compounds where the para substituent is far away from the.hydroxyl group. The phenomenon may also be satisfactorily accounted for on the basis of changing hybridisation, which changes with the substi• tuent (148). 197

Table XXV

Spin-Spin Coupling of the Hydroxyl Proton in

Secondary Alcohols of the Type R1CHOHR2

TMS Compound R2 JgH-OH Reference Solvent

CH3CHOHCH3 CH3 CH3 3.8 Internal C2Cl4

CI3CCHOHCCI3 CCI3 CCI3 7.8 External Acetone

F3CCHOHCF3 CF3 CF3 7.7 Internal DMS

Diacetone-glucose- 5.0 External C2Cl4

Diacetone-glucose 4.8 External Acetone

Fluoral Hydrate Dimer CF3 OR 8.5 External Acetone

CF3CH(OH)2 CF3 OH 7.2 External Acetone

CCI3 CC13CH(0H)2 OH 6.5 External Acetone PhCHOHPh Ph- Ph 3.5 External CJJCU

4,4' -Dimethyl P-CH3R1 p'-CHsPh 3-7 External Acetone Benzhydrol

PhCHOHCF3 CF3 Ph 5.5 Internal DMS

-CH3OR1 p-CH3OPhCHOHCF3 CF3 P 5-5 Internal DMS

£ -naphtyl-CHOHCFg CF3 f$ -naphtyl 5-5 Internal DMS

p-ClPhCHOHCF3 CF3 p-ClPh 5.5 Internal DMS 198

Table XXVI

Spin-Spin Coupling of the Hydroxyl Proton in

Primary Alcohols of the Type RCHgOH

TMS Compound R JCH-02 Reference Solvent

CH3OH H 4.8 Acetone

CBgCHaOH CH3 4.8 — —

CF3CH2OH CF3 7.2 Internal CCI4

CF3CH2OH CF3 6.5 Internal DMS n-butyl alcohol n-prpp. 4.8 External Neat (154) iso-butyl alcohol i-prop. 4.5 External Neat (154)

Citronellol R-CH2 5-0 Internal DMS

PhCH20H Ph 5-5 External Neat (154) 199

DISCUSSION AMD CONCLUSION

Most topics were discussed directly in the body of the text, however, certain aspects will be mentioned here briefly. The most important of these is the ability to detect the OH spin-spin splitting.

It has been demonstrated previously (132,133,134,154) that the coupling of the mobile proton with the adjacent nuclei may be clearly detected in carefuly purified samples. This could be taken as an indi• cation that autoprotolytic equilibria are not important in the detec• tion of the spin-spin coupling.

Absence of the coupling in a neat sample, but its appearance upon dilution with solvent may be understood in terms of the fast pro• ton exchamges, e.g. acid - base equilibria. Trace amounts of acid in the sample will cause the collapse of the splitting. Addition of the solvent, with greater basic properties than the sample, results in the competition between the protonation of the sample and the solvent and as a consequence, the splitting appears, e.g. neat citronellol and citronellol in dimethylsulfoxide. Dilution also decreases the collisional probability between the sample and the acid (base) impu• rity and thus increases the probability of appearance of the splitting

(e.g. neat trifluoroethanol and trifluoroethanol in carbon tetrachloride).

The choice of solvent appears to play and important role in the detection of the OH spin-spin splitting. Of the many conventional n.m.r. solvents known, dimethysulfoxide appears to have certain advan• tages in that it is a good solvent for a variety of compounds, in some 200

cases better than CDC13, CCI4 or acetone. Only a small amount need be added and the OH splitting is observed. As a consequence the con• centration of the sample is still great enough and the spectrum is su• fficiently intense and sharp. In most cases studied herein, dimethyl• sulfoxide proved to be the best solvent to bring up the OH splitting.

Tetrachloroethylene does not appear to have been used as an n.m.r. solvent. Several of its properties may be quoted here.

It is easily obtainable in a pure state and its solvating power is

greater than that of CDC13 or CCI4. It has no absorption in the proton region of the n.m.r. Many successeful runs have been made in this solvent and the OH splitting clearly observed (e.g. diacetone-glucose and other alcohols). Moreover, the OH resonance in this solvent occurs at high field and a shift of the OH to a low field may be produced by addition of strongly hydrogen bonding solvent, such as dimethylsulfo• xide. In this way the OH absorption may be identified, and in some

cases the systems ABX^ or ABaXn may be studied (cf. ref. 155).

The hydrogen bond phenomenon has not been solved in this investigation, however the ability to observe two distinct species of trifluoroethanol in different solvents might provide a basis for further extensive studies. 201

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155. P.L. Corio, R.L. Rutledge and J.R. Zimmerman, J. Molecular Spectr., 3_, 502 (1959). time sec x FIG. 2 Oxidation of Fluoral Hydrate 3*0 3*1 3*2 3*3 3*4 3*5 1000/ T°K

FIG. 5 Oxidation of Fl uoral Hydrate Lfnear relation between pH

k2 and log r*—— *2max -k2

o

/ slope = ro

6 7 8 9 10 11 PH -3*2- \ FIG. 6 \ Oxidation of Fluoral Hydrate B (Mono-anion) \ Variation of rate constant \ with temperature -3*4- \

-3*6- >> \ \ I- OJ U) o -38-

O CF3CH(:OH L \ \ x = -0*1 -4*0- y =+0*9 •\ \

• CF3CD( v :OH)2 V \

-4'2- 3*1 3'2 3*3 3*4 3*5 3*6 1000/T°K+x

FIG 8 Oxidation of Fluoral Hydrate (Di-anion) Variation of rate constant with temperature

3*2 3*3 3*4 3*5 1000/ T°K FIG. 10 pKa of Permanganic Acid (using €Xs25)

FfG. 12

pKa of Perman9anic Acid (modified Davis-Geissman method) •roi

<3 A V <• i

<3

ro 2 FIG. 13 Oxidation of Fluoral Hydrate o- in 25 to 46% H2 S04 Relation between rate constant

and o - H_, 3- H0 , • -HR

-1-1 -4 for HD scale

I o g k2

-2-1

-3-

-4-

-5 -r— i 0. 1 2 -4 Ho H HR 3-2 3"3 1000/T 3'4 FIG. 15 Oxidation of Fluoral Hydrate 7H Relation between rate constant and pH, H.

54 u (A

E

3H

2H

H

-5 pH 1& 5

FIG. 16 Oxidation of Fluoral Hydrate FIG. 17

FIG. 20 Oxidation of Formic Acid in Perchloric and Sulfuric Acids

Plot of logk2 vs H.

nef. 59

; o HCO2H

3 DCQ2H

— ---Extrapolated

0 -1 H_ 1500 -4'4 I i - -—i—• i ' ——r- 3*1 3*2 3*3 3*4 3*5 1000/T°K

FIG. 25 H1 Spectrum of Fluoral Hydrate after Standing

7 6 ^ 5 £ pp**

H0

FIG. 28 H1 Spectrum of Fluoral Hydrate

(39 JJ\ H20)

'5 i —»• FIG. 29 H1 Spectrum of Fluoral Hydrate

FIG. 31 H1 Spectrum of Fluoral Hydrate

(+30 jul H2Q)

'6

H, FIG 32 F1* Spectrum of Fluoral

Hydrate (+30 >ul H20) FIG. 33 H1 Spectrum of Fl. Hydr

(+40/JI H20)

H20

1

II J

J '

6 5

(+40/J I HPO) FIG. 35 H1 Spectrum of Fl. Hydr.

(+160 jul H20) FIG. 36 F^Spectrum of Fluoral Hydrate

(+160 jul H20) FIG. 37 , H' Spectrum of Fluoral Hydrate I after Standing FIG. 38 H' Spectrum of

CF3CD(OH)2

7 ' r6 £

Ho FIG. 39 Hl Spectrum of Chloral Hydrate in CDC 13 FIG. 40 H' Spectrum of Chloral Hydrate in Acetone

H0 FIG. 41 Hl Spectrum of Chloral Hydrate in Acetone

(+10 jul H20) FIG. 42 FT Spectrum of Isopropyl Alcohol FIG 43 H' Spectrum of CF3CHOHCF in Acetone

H0 FIG. 44 H'Spectrum of CF3CHOHCF3 in Acetone-DMS

6 5 4 Sppt* 3 DMS = Dimethylsulfoxide 73 * FIG. 45 Hl Spectrum of PhCHOHCF in DMS FIG. 46 Hl Spectrum of Trifluoroethanol

in DMS overlapping / sideband

V

'5 pp IT! H, FIG. 47 H' Spectrum of Trifluoroethanol in CCI4 FIG. 48 H1 Spectrum of Diacetone-glucose (Ring protons)

1

Pi f / ' 1 1 r

5 4 & ppm Ho FIG. 49

H1 Spectrum of Diacetone-glucose (Ring protons and OH)

Ho H0 FIG. 51 H* Spectrum of Citronellol FIG. 5 2 H' Spectrum of •galactose (Ring protons) C i—**»

A B FIG. 5 3 H Spec rum of Diacetone galactose (Ring protons)

"" E

HDO

S i deband D 11

H,