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Mechanism of the chromic acid oxidation of alcohols. Lee, Donald Garry 1963

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MECHANISM OF THE CHROMIC ACID OXIDATION OF ALCOHOLS by DONALD GARRY LEE B.A., University of Saskatchewan, 195# M.A., University of Saskatchewan, I960 A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF DOCTOR OF PHILOSOPHY in the Department of CHEMISTRY We accept t h i s thesis as conforming to the required standard THE UNIVERSITY OF BRITISH. COLUMBIA March, 1963 In presenting t h i s t h e s i s i n p a r t i a l f u l f i l m e n t of the requirements for an advanced degree at the U n i v e r s i t y of B r i t i s h Columbia, I agree that the L i b r a r y s h a l l make i t f r e e l y a v a i l a b l e for reference and study. I furth e r agr.ee that per-mission for extensive copying of t h i s t h e s i s for s c h o l a r l y purposes may be granted by the -Head of my Department or by h i s representatives» I t i s understood that copying.'or p u b l i -c a t i o n of t h i s t h e s i s for f i n a n c i a l gain s h a l l not be allowed without my w r i t t e n permission. Department of C^^^^yyvNui^^^Tc^ The U n i v e r s i t y of B r i t i s h Columbia,. Vancouver 8, Canada. Date 1y^&-i£ 1 / & 3  PUBLICATIONS Rearrangement Studies with C^4, X. Ethanol-2-C 1 4 from the I r r a d i a t i o n of ethanol-1-C 1 4 with Co 6 0 Gamma Rays. Can. J. Chem. 38. 2 3 1 5 (1960) C.C. Lee and D.G. Lee 2 - P h e n y l e t h y l - I o d i d e from the I r r a d i a t i o n of 2-Phenylethyl-l-C 1 4 Iodide with Cobalt-60 Gamma Rays. Radiation Research 17, 536 (1962) C.C.. Lee, D.G. Lee, and I.S. Bhardwaj Rearrangement Studies with C 1 4, XIII. The Thermal Decomposition of l-C^ 4-2-Butyl Chloro-sulphite. Tetrahedron 18, 1395 (1962) C.C. Lee J.W. Clayton, D.G. Lee and A.J. Finlayson The University of B r i t i s h Columbia FACULTY OF GRADUATE STUDIES PROGRAMME OF THE " FINAL ORAL EXAMINATION FOR THE DEGREE OF DOCTOR OF PHILOSOPHY of DONALD GARRY LEE B.A. , The University of Saskatchewan, 1958 M.A. , The University of Saskatchewan, 1960 MONDAY. APRIL ,29th^ 1963, AT 9s30 A.M. IN ROOM 261, CHEMISTRY BUILDING COMMITTEE IN CHARGE Chairman: F.H. Soward W.A. Bryce R.E. Pincock W.R. Cullen R. Stewart C.A. McDowell D.S. Scott External ExaminerF.H. Westheimer Harvard University MECHANISM OF THE CHROMIC ACID OXIDATION OF ALCOHOLS ABSTRACT The r e l a t i v e rates of chromic acid oxidation of eight aryltrifluoromethylcarbinols were determined i n 77.27o acetic acid medium. The primary deuterium isotope effects of these alcohols ranged from 7.40 for j D-tolyltri ;fluoromethylcarbinol to 12.93 for 3.5 dinitrophenyltrifluoromethylcarbinol at 25° . The rates f i t t e d a Hammett plot best when modified <T + substituent values were used and a modified "rho" value of -1.01 was observed. The rates of oxidation of several halogenated-2-propanols as well as their corresponding 2-d ana-logues were also determined in 77.27, acetic acid and i n 50.17, s u l f u r i c acid. Results indicated that whil the isotope effect varied with changes in the nature of the alcohol being oxidized, i t was constant for a particular alcohol i n solutions of varying a c i d i t y . The thermodynamic parameters for the oxidation of two different alcohols under a variety of conditions were determined and found to be approximately the same, A H# = 9 kcal/mole, AS# = -40 cal/mole deg. An extensive investigation of the oxidation of alcohols i n s u l f u r i c acid solutions ranging from 5 to 957o was conducted and a spectroscopic study was made of the behaviour of chromium VI in this entire region. A Q - * value was obtained for the CF3 group and this was used to determine the "true rate constant", k = k 0bs/ K> f° r the chromic acid oxidation of a series of primary alcohols,, A Taft (j* - p* plot then yielded a p* value of -0,92 for the rate de-termining step of this reaction. A spectroscopic investigation of the v a r i a t i o n of the pKa of chromic acid with changes i n the i d e n t i t y of the mineral acid solvents was made, and the re-suits were applied to explain observed variations i n the rate of oxidation of isopropyl alcohol i n solu-tions of' different mineral acids a l l having the same ac i d i t y . The results are best explained i n terms of a uni-molecular decomposition of a chromate ester v i a a c y c l i c t r a n s i t i o n state. GRADUATE STUDIES F i e l d of Study: Physical Organic Chemistry Topics in Physical Chemistry Topics i n Organic Chemistry Physical Organic Chemistry Inorganic Reaction Mechanisms Organic Reaction Mechanisms R.F. Snider-J.A.R. Coope R„A. Bonnett R, Stewart J. Kutney R. Stewart R.E„ Pincock J. Halpern R.E, Pincock Related Studies: D i f f e r e n t i a l Equations Light S.A. Jennings J.J. Veit ( i i ) ABSTRACT The r e l a t i v e rates of chromic a c i d oxidation of eight a r y l t r i f l u o r o m e t h y l c a r b i n o l s were determined i n 77.2$ ac e t i c acid medium. The primary deuterium isotope e f f e c t s f o r these alcohols ranged "from 7.40 f o r £-tolyltrifluoromethylcarbinol to 12.93 f o r 3,5-dinitrophenyltrifluoromethylcarbinol at 25°. The rates f i t t e d a Hammett plot best when modified <5N+ substituent values were used and a modified wrho t t value of -1.01 was observed. The rates of oxidation of several halogenated-2-propanols as well as t h e i r corresponding 2-d analogttas were also determined i n 77.2$ ace t i c a c i d and i n 50.1$ s u l f u r i c a c i d . Results indicated that while the' isotope e f f e c t varied with changes i n the nature of the alcohol being oxidized, i t was constant f o r a p a r t i c u l a r alcohol i n solutions of varying a c i d i t y . The thermodynamic parameters f o r the oxidation of two d i f f e r e n t alcohols under a v a r i e t y of conditions were determined and found to be approximately the same, = 9 kcal/mole, = -40 cal/mole deg. An extensive i n v e s t i g a t i o n of the oxidation of alcohols i n s u l f u r i c a c i d solutions ranging from 5 to 95$ was conducted and a spectroscopic study was made ( i i i ) of the behavior of chromium VI i n t h i s entire region,, A <5~* value was obtained f o r the CF^ group and t h i s was used to determine the "true rate constant", k = k Q l D 5/K, f o r the chromic acid oxidation of a •if- $ series of primary alcohols. A Taft - 6 p l o t then yielded a £ value of -0.92 f o r the rate determining step of t h i s reaction. A spectroscopic investigation of the v a r i a t i o n of the pKa of chromic acid with changes i n the i d e n t i t y of the mineral acid solvents was made, and the r e s u l t s were applied to explain observed variations i n the rate of oxidation of isopropyl alcohol i n solutions of d i f f e r e n t mineral acids a l l having the same a c i d i t y . The H Q a c i d i t y function f o r n i t r i c and phosphoric N acid systems with a 1:1 molar r a t i o of sodium perchlorate was also determined. ACKNOWLEDGMENTS The author wishes to express his sincere appreciation to Dr. Ross Stewart f o r his guidance during the course of t h i s research and f o r h i s constructive c r i t i c i s m during the preparation of t h i s t h e s i s . Grateful acknowledgment i s also made to the University of B r i t i s h Columbia f o r the award of two Teaching Assistantships (1960-61 and 1961-62), and to the National Research Council of Canada f o r the award of a Student-ship (1961-62). (v) TABLE OF CONTENTS Page I. HISTORICAL INTRODUCTION 1. Chromic Acid^Oxidation of Organic Compounds 1 2. Chromic'Acid Oxidation of Alcohols 3 3. Primary Kinetic Isotope E f f e c t s 24 I I . SCOPE OF THE PRESENT INVESTIGATION 36 I I I . EXPERIMENTAL 39 1. Materials ( i ) Alcohols 39 ( i i ) Inorganic Reagents 53 ( i i i ) Indicators 54 2. K i n e t i c Methods 55 (i) Iodometric Method 56 ( i i ) Spectrophotometric Method 64 3. Consideration of the Acid Chromate-Bichromate Equilibrium 64 4. Product Analysis 65 5. Determination of pKa Values f o r HgCrO^ 67 6. A c i d i t y Function f o r the Phosphoric Acid-Sodium Perchlorate System 72 7. A c i d i t y Function f o r the N i t r i c Acid-Sodium Perchlorate System 79 Determination of the Chromate Ester Equilibrium Constant f o r 2-Methoxy-1 ethanol and 2 , 2 , 2-Trifluoroethanol 80 (vi) IV. RESULTS AND DISCUSSION 1. Oxidation of Fluoro Alcohols i n 77.2$ Acetic Acid 2. Oxidation of some Halogenated 2-propanols i n 50.1% S u l f u r i c Acid i 3. Variation i n the Magnitude of Isotope E f f e c t s with Changes i n A c i d i t y 4» Thermodynamic Parameters 5. Oxidation i n Concentrated S u l f u r i c Acid Solutions 6. Determination of the 6 Value f o r the CF^ Group 7. Determination of the "True Rate Constant", k2 = kobs/ K* f o r t h e C h r o m i c Acid Oxidation of some Primary Alcohols 8. A c i d i t y Functions f o r N i t r i c and Phosphoric Acid Solutions with Added Sodium Perchlorate 143 9. Variation i n the pKa of Chromic Acid with Changes i n the Identity of the Mineral Acid Solvent 149 10. Variation i n Rate of Chromic Acid Oxidations with Changes i n the Identity of the Mineral Acid Solvent 166 V. CONCLUSION 175 VI„ SUGGESTIONS FOR FURTHER RESEARCH 181 VII. APPENDIX 133 VIII. BIBLIOGRAPHY 134 Page 85 101 107 109 113 136 138 ( v i i ) TABLES I. Va r i a t i o n i n Theoretical Isotope E f f e c t s 28 with Temperature Hi Data f o r Rate Plots 59 I I I . A c i d i t y of Solutions Containing a 1:1 Molar Ratio of H 3.P0 4 to NaClO^ 74 IV. A c i d i t y of Solutions Containing a 1:1 Molar Ratio of HNO^  to NaClO^ 75 V. Data f o r Figure 10 84 VI. Oxidation of Fluoro alcohols, i n 77.2% Acetic Acid 86 VII. V a r i a t i o n i n Rate of Oxidation of Phenyltrifluoromethylcarbinols i n 77.2$ Acetic with H Q 101 VIII. Rates of Oxidation of Substituted 2-Propanols i n 50.1$ S u l f u r i c Acid 104 IX. Deuterium Isotope Effects f o r the Oxidation of Isopropyl Alcohol i n Perchloric Acid Solutions of,Varying A c i d i t i e s 108 X. Heats and Entropies of Activation f o r the Oxidation of Substituted 2-Propanols Under Various Conditions 110 XI. Rates Used i n Calculation of Thermodynamic Parameters 111 XII. E x t i n c t i o n C o e f f i c i e n t s of Chromium VI i n Su l f u r i c Acid Solutions 117 XIII. Oxidation of CF^CROHCH^ by Chromium VI i n S u l f u r i c Acid Solutions at 25.0°C. 123 XIV. Oxidation of CH3CROHCH3 i n S u l f u r i c Acid Solutions 124 XV. Oxidation of 1 , 1 , 1 , 3 , 3 , 3-hexafluoro - 2 -propanol i n S u l f u r i c Acid Solutions 125 ( v i i i ) XVI. Oxidation of Substituted Methanols i n 3.32M HCIO^ i; 139 XVII. Calculation of the "True Rate Constants'*, k 0 = k , /K f o r some Substituted 2 O D S ' Methanols i n Aqueous Perchloric Acid Solutions 141 XVIII. A c i d i t i e s of Solutions Containing a 1:1 Molar Ratio of Phosphoric Acid and Sodium Perchlorate 144 XIX. A c i d i t i e s of Solutions Containing a 1:1 Molar Ratio of N i t r i c Acid and Sodium Perchlorate 145 XX. Comparison of the pKa of HgCrO^ i n Aqueous Solution of Various Mineral Acids and the Point of Slope Change i n a Plot of log k v.s. H o f o r these Acids 151 XXI. Spectral Data f o r the Acid Chromate Ion i n Various Mineral Acids 153 XXII. Spectral Data f o r Chromic Acid i n Various Mineral Acids 154 XXIII. Rate of Oxidation of Isopropyl Alcohol i n Mineral Acids of a Given A c i d i t y (H Q - -2.5) 155 XXIVo Rates of Oxidation of Isopropyl Alcohol i n S u l f u r i c Acid Solutions 169 XXV. Rates of Oxidation of Isopropyl Alcohol i n Perchloric Acid Solutions 170 XXVI. Rates of Oxidation of Isopropyl Alcohol i n Phosphoric Acid Solutions 170 XXVII. Rates of Oxidation of Isopropyl Alcohol i n Hydrochloric Acid Solutions 171 XXVIII. Rates of Oxidation of Isopropyl Alcohol i n N i t r i c Acid Solutions 172 (ix) XXIX. Rates of Oxidation of Isopropyl Alcohol i n Solutions Containing a 1:1 Molar Ratio of H 3P0 4 to NaClO^ 173 XXX. Rates of Oxidation of Isopropyl Alcohol i n Solutions Containing a 1:1 Molar Ratio of HNO, to NaClO. 173 (x) FIGURES 1. Typical Second Order Rate Plots f o r the Oxidation of Substituted P h e n y l t r i -fluoromethylcarbinols i n 77*2% Acetic Acid 57 2. Typical F i r s t Order Rate Plots f o r the Oxidation of Isopropyl Alcohol 58" 3. pKa of H 2Cr0 4 i n H 2S0 4 69 4. pKa of H 2CrO^ i n HCIO^ 69 5. pKa of H 2Cr0 4 i n HC1 70 6. pKa of H 2CrO^ i n HN03 70 7. pKa of H 2Cr0 4 i n H3PO^-NaC104 71 8". Ionization of Indicators i n rLPO, -NaCIO, Solutions -> * ^ 77 9. Ionization of Indicators i n HNO^-NaClO, Solutions j> 10. Chromate Ester Equilibrium Constants #3 11. Hammett c r + ~ <£ Plot f o r the Chromic Acid Oxidation of Substituted P h e n y l t r i f l u o r o -methylcarbinols i n 77$ Acetic Acid 8*9 12. Hammett cr - € Plot f o r the Chromic Acid Oxidation of Substituted P h e n y l t r i f l u o r o -methylcarbinols i n 77$ Acetic Acid 90 13. Relationship Between the Rate of Oxidation of Trifluoromethylcarbinols i n 77$ Acetic Acid and the Primary Deuterium Isotope E f f e c t 96 14. Linear Relationship Between Rate and A c i d i t y f o r the Oxidation of Phenyltrifluoro-methylcarbinol i n 77.2$ Acetic Acid 100 15. Taft <T*- 6* Plot f o r the Oxidation of Substituted 2-Propanols i n 50.1$ H 2S0^ 106 (xi) 16. Thermodynamic Parameters 1 1 2 17. Extinction C o e f f i c i e n t s f o r Chromium VI in S u l f u r i c Acid Solutions 116 18. U. V. Spectra of Chromium VI i n H9SO, Solutions * * 118 19. Attempted Determination of the pKa f o r p_-Toluic Acid i n Acetic Acid-Sulfuric Acid Media 120 2 0 . Relationship Between H and the Rate of Oxidation of three 2-Propanols i n Su l f u r i c Acid Solutions 1 2 2 2 1 . Oxidation of l , l , l - T r i f l u o r o - 2 - p r o p a n o l i n 5 7 . 6 $ H 2 S 0 ^ 1 2 9 2 2 . Oxidation of 1 , 1 , 1 - T r i f l u o r o - 2 - p r o p a n o l i n 6 3 . 9 $ H 2 S 0 ^ 1 3 0 2 3 . Oxidation of l , l , l - T r i f l u o r o - 2 = p r o p a n o l i n 6 8 . 5 $ H 2 S 0 4 1 3 1 24. Oxidation of l , l , l - T r i f l u o r o - 2 - p r o p a n o l i n 7 6 . 4 $ H 2 S 0 4 1 3 2 2 5 . Oxidation of l , l , l - T r i f l u o r o - 2 ~ p r o p a n o l i n 8 5 . 7 $ H 2 S 0 ^ 1 3 3 2 6 . Oxidation of l , l , l - T r i f l u o r o - 2 - p r o p a n o l i n 9 5 . 3 $ H 2 S 0 4 1 3 4 2 7 . Relationship Between K o b g f o r the Oxidation of Substituted Methanols and the cT^ Value f o r the Substituents 1 4 2 2 8 . Relationship Between the True Rate Constant, k 2=k Q b s/K and <s** Values 1 4 2 2 9 . Comparison of the H o Function f o r a 1 : 1 Molar Ratio of H^PO^ and NaClO^ with that f o r H^PO^ 1 4 6 ( x i i ) 30. Comparison of the H Q Function f o r a 1:1 Molar Ratio of HN03 and NaClO^ with that f o r HNO^  147 31. U l t r a v i o l e t Spectrum of Chromium VI i n 3 .50 M HCIO^ 156 32. U l t r a v i o l e t Spectrum of Chromium VI i n 3.10 M H 2S0 4 157 33. U l t r a v i o l e t Spectrum of Chromium VI i n 3.60 M HC1 158 34. U l t r a v i o l e t Spectrum of Chromium VI i n 3.00 M H 3P0 4 159 35. rtpKan of HCrS0 ? 165 36. »pKa w of HCr0 3Gl 165 37. Chromic Acid Oxidation of Isopropyl Alcohol i n Aqueous Solutions of Mineral Acids 16$ I. HISTORICAL INTRODUCTION 1 . CHROMIC ACID OXIDATION OF ORGANIC COMPOUNDS While chromic acid has long been used as an o x i d i z i n g agent i n preparative organic chemistry i t i s only i n comparatively recent years that k i n e t i c or mechanistic investigations have been conducted. Since the f i r s t k i n e t i c studies performed by Westheimer and Novick ( 1 ) on the chromic acid oxidation of isopopyl alcohol, investigations have been carried out on the mechanism of oxidation of hydrocarbons ( 2-M), aldehydes ( 12-17), ketones ( 1 8 , 19) carboxylic acids ( 2 0 , 2 1 ) , ethers ( 2 2 ) , o l e f i n s ( 2 3 ) and glycols ( 2 4 - 2 6 ) , as well as alcohols. The l a t t e r group w i l l be discussed i n d e t a i l l a t e r . In addition to the two most commonly used solvents, water and acetic acid, the reaction has been studied i n acetic anhydride ( 1 5 ) , propionic acid ( 2 3 ) , propionic anhydride ( 2 3 ) , benzene ( 2 7 ) , pyridine ( 2 7 , 2 8 ) , aqueous acetone ( 2 9 ) , and concentrated s u l f u r i c acid ( 3 0 , 2 1 , 3 1 ) . While the d e t a i l s of the reaction vary considerably from one type of compound to another and from one solvent to another several very general features can be observed. In a l l cases studied the reaction was found to be a c i d catalyzed, with the rate being proportional to some power of ho ( 1 2 , 2 0 , 3 D . Also 2 i n a l l studies reported to date but one i t has been observed that electron withdrawing groups attached to the organic substrates slow down the rate of reaction while electron donating groups increase the rate (20, 27, 32). The one exception i n t h i s case i s the oxidation of substituted benzaldehydes i n 91% a c e t i c a c i d where Wiberg and M i l l found a H r h o " value of +1.02 (12). Roc'ek, on the other hand, has reported that the chromic acid oxidation of a series of substituted acetaldehydes gives a "rho*" value of -1.2 (14, 33). F i n a l l y , f o r a given a c i d i t y i t has been found that reactions i n acetic acid are always f a s t e r than the corresponding reactions i n aqueous solution (34), and that increasing the proportion of a c e t i c a c i d i n mixed a c e t i c acid-water solvents 1 usually r e s u l t s i n a corresponding increase i n rate (35, 36). Since the investigations of t h i s reaction have been so extensive and of such a large v a r i e t y no attempt w i l l be made i n t h i s thesis at a comprehensive review of a l l the l i t e r a t u r e . In-stead attention w i l l be directed to the oxidation of alcohols by t h i s reagent, and only such material that d i r e c t l y pertains to t h i s reaction w i l l be included. 3 2. CHROMIC ACID OXIDATION OF ALCOHOLS The l i t e r a t u r e on t h i s reaction was reviewed by Westheimer (37) i n 1949, and again by Waters (38) i n 1958o In addition to these two review a r t i c l e s Wiberg (2) and Kwart (39) have both recently published extensive resume's of the l i t e r a t u r e pertaining to t h i s subject. In a series of investigations over a period of some seventeen years, Westheimer and coworkers have made a det a i l e d study of the mechanism of the chromic acid oxidation of isopropyl alcohol. The re s u l t s they have obtained indicate that the reaction i s f i r s t order i n H + concentration i n d i l u t e acid solutions and second order i n H + concentration i n more a c i d i c solutions (1); that the i n i t i a l reactions involves a 2-electron change (40); that i n d i l u t e a c i d i c aqueous solutions the only active o x i d i z i n g species i s HCrO^~ (1); that isopropyl alcohol i s oxidized much more ra p i d l y than diisopropyl ether (22); that the rate determining step involves f i s s i o n of the oc-carbon-hydrogen bond ( k H / k D =6.6) (41); that although s i m i l a r i n many other respects the reaction i n 86.5$ ace t i c acid i s 250 times as f a s t as i n water of a correspond-ing a c i d i t y (34); and that the reaction i s fas t e r i n a c i d i c solutions of Do0 than i n solutions of Ho0 4 containing the same concentration of acid ( k D QA^ 0 = 2.44 f o r the portion of the reaction which i s f i r s t order i n H concentration and kp QA^ Q = 6.26 f o r the portion of the reaction which i s second order i n H + concentration) (22). Thfcsedata together with the observation that a chromate ester of isopropyl alcohol could be prepared i n water and extracted into benzene, and further that addition of water or bases such as pyridene caused the ester to decompose to acetone and chromium I I I (42) led Westheiraer to propose a mechanism i n which the isopropyl chromate ester was an intermediate. . Ka HCrO^" + H 30 + ^==^ H 2 C R 0 4 + H 2 ° ( A H 2Cr0 4 + R2CHOH R 2CH0Cr0 3H + H 20 (b R 2CH0Cr0 3H + H^O* ^==± RgCHOCrOjHg + H20 (c k. 1 R 2CHOCr0 3H + HgO ' RgC - 0 + C r 1 V + H 30 + (d k 2 R 2CHOCr0 3H 2 + + HgO * RgC = 0 + C r 1 V + H 30* (e This mechanism seemed to amply explain a l l the observed features of the reaction and Westheimer was also able to use i t to explain why chromic a c i d oxidations were f a s t e r i n acetic acid medium than i n water and why addition of chloride ion decreased the 5 rate of reaction (34)o He suggested that the position of the equilibrium i n equation (a) would be s h i f t e d to the r i g h t due to a lowering of the d i e l e c t r i c constant i n a c e tic acid and that the position of the equilibrium i n equation (b) would be shifted to the r i g h t due to a simple mass action e f f e c t . These two e f f e c t s together were assumed to be s u f f i c i e n t to overcome any reduction i n the rate of ester decomposition caused by a lowering of the d i e l e c t r i c constant. Westheimer's f i r s t suggestion that the pKa of chromic acid would be d i f f e r e n t i n acetic acid than i n aqueous medium bears further consideration. I f the proper H_ a c i d i t y functions were used the pKa of H 2CrO^ should be the same in both media. However i t i s possible that chromic acid does not e x i s t as the same species i n these two solvents, i . e . an acetate ligand may be incorporated into the coordination s h e l l of the chromium or a mixed anhydride may be formed. One would then be comparing the pKa's of two d i f f e r e n t species such as H0Cr020H and HOCrOgOAc. These two species might e a s i l y have d i f f e r e n t pKa's. In f a c t , Klaning and Symons (43) have recently measured the equilibrium constant f o r the formation of such a mixed anhydride; HCrO, + AcOH AcOCrO. + H o0 . 6 K was found to be 4.5. A simple c a l c u l a t i o n then shows that i n 36.5$ acetic acid 90$ of the t o t a l Chromium VI present i s in the form of the mixed anhydride. Hence, i f H0Cr020Ac were a more active o x i d i z i n g reagent than H 2CrO^ the rate enhancement found i n acetic acid would be explained. Further, i n explaining the reduction i n rate caused by addition of chloride ions Westheimer found i t necessary to invoke the formation of chlorochromate ions (34). C l C r 0 3 " + H 30 + i = ^ C l C r 0 3 H + H 20 (f) ClCr0 3H + R2CH0H R 2CH0Cr0 2Ci + H 20 (g) R 2CH0Cr0 2Cl + H 20 > R 2C = 0 + CrOgCl" + H^O* (h) He contended that while incorporation of chloride i n t h i s way would not have much e f f e c t on the position of the e s t e r i f i c a t i o n equilibrium i t would lower the pKa of the a c i d and thus s h i f t the equilibrium i n (f) as compared with (a) to the r i g h t . Although there would also be an e f f e c t on reaction (h) i t was thought that i t would not be of as great a magnitude as the e f f e c t on the pKa of the a c i d . The ester mechanism has also been used by Wiberg and M i l l (12), and by Graham and Westheimer (13) to account s a t i s f a c t o r i l y f o r the k i n e t i c r e s u l t s observed when benzaldehyde i s oxidized by chromic 7 acid. The rate law i n t h i s case was found to be, V = k [RCHO] [HCrO^ ""] h Q and a n r h o w value of +1.02 was observed. The mechanism proposed involved a pre-equilibrium ester formation by reaction of one molecule of ^ C r O ^ with a molecule of aldehyde. HCrO^"" + H + ^ = ± H 2 C r 0 4 0 OH ArCH + H oCr0, Ar-A-OCrO-jH 2 k | 3 H OH l Ar-C-OCKUH + H o0 * ArC0 oH + Cr i 3 2 2 H IV As has been previously mentioned t h i s i s the only reported instance i n which electron a t t r a c t i n g substituents increase the rate of oxidation, and there has been considerable discussion concerning i t i n the l i t e r a t u r e . Rocek (33)» i n l i g h t of the negative r»rho*n value of -1.2 which he observed f o r the oxidation of four substituted acetaldehydes has suggested that the positive n r h o " observed by Wiberg and M i l l might be due to the fac t that i t i s the aldehyde hydrates which are being oxidized. I f the "rho"for the hydration equilibrium were strongly 8 p o s i t i v e i t would cancel the negative "rho w f o r oxidation. In reply to t h i s c r i t i c i s m Wiberg and Evans (2) have estimated from a comparison with the permanganate oxidation of aldehydes that the "rho H f o r hydration i s approximately unity. This implies that the "rho" f o r the r e s t of the reaction i s almost zero. In addition these authors point out that i t i s possible to reconcile a negative "rho" value with a rate determining proton abstraction step i n two ways. F i r s t l y , i f the activated complex strongly resembled the products, resonance s t a b i l i z a t i o n of the acid by electron donating groups would lead to an enhanced rate. Secondly, i t must be borne i n mind that i n addition to a proton loss two electrons must also be transferred to the oxidant, and i t i s conceivable that the e f f e c t of substituents could be mainly concerned with the electron transfer. The exact fate of the chromium IV species produced by these mechanisms i s as yet unknown. Westheimer o r i g i n a l l y postulated that one chromium IV species would combine with a chromium VI species to give two chromium V species which could then react further with alcohol by 2-electron changes. Cr VI + Cr IV Cr v + R-CHOH * R 2C = 0 + Cr I I I 9 Evidence favoring t h i s type of a sequence has come from a study of the oxidation of phenyl-t-butylcorbine.1 by chromic a c i d (44). The products of t h i s reaction are phenyl t-butyl ketone, benzaldehyde and t-butylalcohol. Since the y i e l d of cleavage products (benzaldehyde and t-butylalcohol) was sharply reduced by addition of Ce or Mn to the reaction medium i t was postulated that the cleavage products resulted from an oxidation of the phenyl-t-butylcarbinol by intermediate Cr^ and Cr"^ species. In the presence of Mn + + or Ce"'" the intermediate chromium species would be reduced before they could p a r t i c i p a t e i n a cleavage reaction. Further, the maximum y i e l d of cleavage products i n the absence of C e + + + or Mn + + was about 67$ in d i c a t i n g that the most l i k e l y active intermediate was C r V and not Cr* V, I f the l a t t e r species were responsible f o r the cleavage reaction a maximum y i e l d of only 33$ would be expected. Subsequent work by Kaplan (45) on the competitive oxidation of 2-propanol and 2-propanol-2-t has indicated that rupture of the oc-carbon-hydrogen bond i s also the rate-determining step f o r the oxidation of 2-propanol by the intermediate chromium species. More recently Wiberg and Richardson (16) have V made a study of the oxidation of aldehydes by Gr and 10 Cr . Since these chromium species are unstable under reaction conditions they were studied by two new and probably very useful techniques, namely the competitive oxidation of a mixture of organic compounds and the oxidation of a compound which w i l l give d i f f e r e n t products f o r d i f f e r e n t reaction paths. The general r e s u l t derived from the competitive studies was that the oxidation by the intermediate chromium species exhibited a "rho r t value of the same sign and of approximately the same magnitude as that observed f o r chromium VI oxidation (+0.45 f o r C r I V - C r V c.f. + 0.77 f o r C r ^ 1 ) . This indicates a p o s s i b i l i t y that both the VI intermediate species and the Cr may react by s i m i l a r mechanisms. Results obtained from the oxidation of triphenylacetaldehyde on the other hand, indicate that about one-third of the reaction proceeded v i a a path that did not involve an ester intermediate. Although the r e s u l t s of many subsequent i n -vestigations such as substituent effects on oxidation rates of a group of arylmethylcarbinols (27) and the r e l a t i v e oxidation rates of a x i a l and equatorial cyclanols (46, 47) were also explained i n terms of the ester mechanism, Rofe'ek and Krupic'ka i n 1958 proposed a d i f f e r e n t mechanism involving hydride ion transfer from the oc-carbon atom with simultaneous 11 proton loss by the hydroxyl (31). R R R R \ / C = 0 + Cr IV Their reasons f o r favoring t h i s mechanism as opposed to the ester mechanism were f i v e folds (i) the rate of oxidation of hydrocarbons, where no ester formation i s possible, i s also much greater i n acetic acid than i n aqueous medium of the same a c i d i t y ; ( i i ) the log of the rate constant was found to be l i n e a r with H Q i n d i c a t i n g , according to the Zucker-Hammett hypothesis (4$), that the t r a n s i t i o n state did not contain a molecule of water; ( i i i ) the e f f e c t of substituents on the rate of oxidation of alcohols could be more e a s i l y -explained on the basis of a rate determining hydride ion transfer; (iv) the c a t a l y t i c e f f e c t of pyridine reported by Helloway, Cohen, and Westheimer (42) could not be duplicated; and (v) a hydride abstraction mechanism more e a s i l y explained why c y c l i c alcohols with a x i a l hydroxyls were more r e a d i l y oxidized than t h e i r equatorial epimers. Each of these points requires further elaboration. The increase in rate of oxidation of hydrocarbons i n a c e t i c acid as compared to aqueous solutions of the same a c i d i t y could be due, as has been discussed, to the formation of a d i f f e r e n t more active species of chromium VI i n this medium. Or i t could be due 12 simply to the f a c t that the reaction i s general acid catalyzed. In addition Wiberg and Evans (2) have observed that i n the oxidation of hydrocarbons the rate of reaction i s dependent on the t o t a l chromium VI concentration, while the rate of oxidation of alcohols and aldehydes i s dependent only on the concentration of HCrO^° . This implies that quite a d i f f e r e n t mechanism i s operative i n the case of hydrocarbons. In f a c t Wiberg and Evans have suggested that the mechanism may be d i f f e r e n t f o r d i f f e r e n t hydrocarbons with diphenylmethane reacting v i a hydrogen atom abstraction while p-methoxydiphenylmethane and pp'-dimethoxydiphenylmethane react v i a hydride abstraction. Wiberg and Foster (3) have also reported that the oxidation of (+)-3-methylheptane i n 91$ ace t i c a c i d gives (+)-3-methyl-3-heptanol with 70-80$ retention of configuration. The mechanism they proposed f o r t h i s reaction involved hydrogen abstrac-VI V tio n by Cr with the r a d i c a l and Cr recombining to IV form an ester of Cr which subsequently decomposes by chromium-oxygen bond cleavage to give o p t i c a l l y active products. More recently Rocek (4) has b r i e f l y summarized some features of the chromic a c i d oxidation of hydro-carbons which indicate that in some cases the i n t e r -mediate appears to have properties s i m i l a r to a 13 carbonium ion while i n other instances i t appears to have the c h a r a c t e r i s t i c s of a free r a d i c a l . Accordingly he has proposed a mechanism which envisages the i n t e r -mediate as a mesomeric hybrid of a carbonium i o n - C r ^ complex and a ra d i c a l - C r complex. R^CH + H^CrO^ +V IV R-C-0-CrO^H, <—>R^C+ 0-Cr0oH. 3 3 4 3 3 4 IV R 3C-0-Cr0 3H 4 -> R3C-OH + Cr IV Since the Zucker-Hammett hypothesis has not been found to be un i v e r s a l l y applicable (48), Rocek and KrupiSka's observation that the log of the rate constant was l i n e a r with H Q loses much of i t s s i g n i f -icance . Further Graham and Westheimer ( 1 3 ) have pointed out that i f one writes out the entire rate expressions f o r both the ester mechanism and the hydride ion mechanism they are i d e n t i c a l except for possible differences i n the a c t i v i t y c o e f f i c i e n t s for the two activated complexes. Roc'ek ( 3 2 ) has however been reluctant to accept these arguments, contending that a l i n e a r r a t e - a c i d i t y r e l a t i o n s h i p f o r the ester mechanism required an u n l i k e l y s i m i l a r i t y between the terms f H 2 C r 0 4 fROH f R O C r 0 3 H 2 + — H 2 0 and B 14 while the hydride ion mechanism requires a more l i k e l y s i m i l a r i t y between the terms f H 2Cr0 4 fROH a n d fH 3Cr0 4" 1"—ROH f B H + With respect to t h i s point i t should be mentioned that Bunnett (49) has recently developed a method of t r e a t i n g the r e l a t i o n s h i p between reaction rates, H o values and the a c t i v i t y of water i n such a way that one can estimate semi-empirically the number of water molecules involved i n the t r a n s i t i o n state of a re-action. According to t h i s treatment reactions which occur by hydride abstraction with no p a r t i c i p a t i o n of H 20 generally have negative ttww values i n the range of -2 to - 7 , while reactions i n which water acts as a base to abstract protons generally exhibit a p o s i t i v e "w* value of +3 or greater. Although the treatment i s f a i r l y new and has not yet been widely tested Bunnett has nevertheless substantiated i t with a wide v a r i e t y of reactions and i t would appear that the treatment can be applied with considerable confidence. Thus Kwart (39) has recently recast Rocek and Krupicka's data i n terms of t h i s new t r e a t -ment and found that a p l o t of (log k + H Q)v.s. log ajj Q gave a straight l i n e with w = +o.49. Since the 15 t » w » value was not found to be negative Kwart contended that Rocek and KrupiSka's mechanism involving hydride transfer with no p a r t i c i p a t i o n of water was i n c o r r e c t . However, since the wwrt value was not as p o s i t i v e as most proton transfer reactions involving n u c l e o p h i l i c p a r t i c i p a t i o n of water i t could not e a s i l y be recon-c i l e d with the ester mechanism either. Instead Kwart claimed t h i s to be evidence i n favor of a recently proposed c y c l i c ester mechanism involving i n t e r n a l proton transfer ( 5 0 ) . / C v ^ 0 R. . 0 . / 0 R T V ^ C ^ C r > -4 > ^ ^ C r ^ — > ^ C= 0 + G r I V R ^ \ // \ R / 4 \ / / \ R X HO OH H - - - 0 OH Although i t i s extremely d i f f i c u l t to predict how such a mechanism would respond to Bunnett's treatment Kwart has presented other evidence (which w i l l be discussed l a t e r ) that gives more weight to such a mechanism. The report by Westheimer et a l . ( 2 2 ) of the comparative rates of oxidation of isopropyl alcohol i n DgO and ^ 0 also must be taken into account. A consideration of the comparative i o n i z a t i o n constants of strong acids i n DgO and H^ O indicates that the rate of chromic a c i d oxidation should be enhanced i n DgO as compared to ^ 0 . The magnitude of the rate increase, a l l other factors being equal, should be about 2 . 5 f o r 16 the reaction under conditions where i t i s f i r s t order HK 2 i n H and (2.5) = 6.25 under conditions where i t i s + second order i n H . This i s i n good agreement with the factors of 2.4 and 6.3 found by Westheimer. These r e s u l t s are not i n accord with Rocek's hydride transfer mechanism since i t would involve a simultaneous 0-D bond rupture when the reaction was carried out i n D 20. Since such a 0-D bond rupture has been shown to r e s u l t i n a decrease i n the rate of bromine oxidation of 2-propanol by a fa c t o r of 1.49 (77) one would not expect as large a rate increase as Westheimer observed. Further, the large difference i n the rate of oxidation between isopropyl alcohol and diisopropyl ether i n -dicates that a d i r e c t hydride abstraction without 0-H or 0-D cleavage i s u n l i k e l y . Perhaps the strongest point i n Rocek and Krupicka's c r i t i c i s m of the ester mechanism i s the observation which Kwart and Francis (27) had previously made, that the rate constants f o r the oxidation of a series of arylmetSiLylcarbinols f i t t e d a Hammett rel a t i o n s h i p with a n r h o n value of -1.01. If the r a t e -determining step involved proton abstraction a negative charge would be b u i l t up on the oc-carbon i n the t r a n s i t i o n state. Such a t r a n s i t i o n state would predict a p o s i t i v e "rho" value. However i f the re-action went by hydride transfer the carbon would bear 17 a positive charge i n the t r a n s i t i o n state and the re-action would be expected to exhibit a negative "rho r t value. Kwart and Francis have j u s t i f i e d the negative rho value with the ester mechanism by suggesting that the substituents have a much stronger e f f e c t on the p o s i t i o n of the ester equilibrium than they do on the rate-determining step. i . e . The complete rate expression f o r the ester mechanism at moderate a c i d i t i e s i s Kwart and Francis claimed that the substituents would have a greater e f f e c t on Ke than on k 2. However, Klaning and Symons (43, 52) have shown that the 'position of the chromate ester equilibrium i s not greatly affected by changes i n the structure of the alcohol, In addition, r e s u l t s w i l l be presented l a t e r which indicate that a Taft <r — £ plot for the true rate constant (k 2 = k o* 9 S/Ke) s t i l l exhibits a negative "rho*" value. Although t h i s observation raises grave doubts as to the v a l i d i t y of Kwart and Francis" o r i g i n a l explanation Westheiraer (22) and Wiberg (2) have recently presented an a l t e r n a t i v e explanation which can be reconciled to the ester mechanism. By a comparison with the permanganate oxidation of aromatic aldehydes which Wiberg and Stewart (53) have shown to involve a rate-determining 18 proton removal step, as well as a negative "rho n value, Westheimer and Wiberg have argued that the substituents have a greater e f f e c t upon the withdrawal of two ! electrons by the oxidant, than upon the removal of the proton which accompanies the reaction. Since Rocek and Krupicka's report that pyridene does not catalyze the oxidation of isopropyl alcohol Westheimer and Chang (54) have re-examined t h e i r evidence and have come to the conclusion that the re-action i s , i n f a c t , not catalyzed by the addition of pyridene. The rate c a t a l y s i s f i r s t reported by Holloway, Cohen and Westheimer (55) was found to be in error due to the method used in determining a c i d i t i e s of perchloric acid solutions containing dissolved pyridinium perchlorate on one hand and sodium perchlorate on the other. The f i n a l objection to the ester mechanism which Rocek and Krupicka raised concerned the r e l a t i v e rates of oxidation of epimeric s t e r o i d a l alcohols. In the hydroxy-chalestane series i t i s found that regardless i f the hydroxyl i s i n the 1, 2, 3, 4, 6 or 7-position the epimer which contains the a x i a l hydroxyl i s oxidized by chromic a c i d at a f a s t e r rate. (46). Since a x i a l hydroxyls are more hindered than equatorial hydroxyls Barton (47) had suggested that the f a s t e r 19 rate f o r the a x i a l epimers resulted from a rate-deter-mining step involving nucleophilic attack of the oc-hydrogen by water. However, i t i s also well known that a x i a l hydroxyls are much more d i f f i c u l t to e s t e r i f y (47) and t h i s would tend to s h i f t the p o s i t i o n of the ester equilibrium to the l e f t . Hence there exists a paradox between t h i s explanation and the suggestion made by Kwart and Francis (27) that the p o s i t i o n of the ester equilibrium determines the r e -l a t i v e rates of oxidation of a series of arylmethyl-carbinols. Roc'ek's hydride transfer mechanism, on the other hand, more e a s i l y explains the observed rates of oxidation of such epimeric hydroxycholestanes. Another explanation f o r the r e l a t i v e rates of oxidation of epimeric hydroxycholestanes, as well as those f o r a series of alkylated cyclohexanols i s that the a x i a l alcohols are oxidized f a s t e r because forma-t i o n of the ketone i s favored by r e l i e f of s t e r i c s t r a i n (46, 56). Such an explanation assumes that the t r a n s i t i o n state f o r the reaction has considerable carbonyl character. In t h i s way s t e r i c s t r a i n i n the tetrahedral ground state would be reli e v e d on passing to a t r i g o n a l t r a n s i t i o n state. However, Kwart (57) has objected to t h i s explanation, pointing out that there i s no p a r a l l e l between the rates of 20 oxidation of c y c l i c alcohols and the s t a b i l i t y of the r e s u l t i n g carbonyl products. For example, cyclopentanol i s oxidized only s l i g h t l y f a s t e r than cyclohexanol while there i s a vast difference i n the r e l a t i v e carbonyl s t a b i l i t y of cyclopentanone and cyclohexanone (39), and with 7-hydroxyborane which on oxidation gives a ketone with a very small valence angle the rate i s only reduced to one-third that of cyclopentanol. The a l t e r n a t i v e mechanism proposed by Kwart and Francis (58) envisages a chromate ester t r a n s i t i o n state i n which proton abstraction occurs by way of a concerted c y c l i c process. 0 Acceleration of the rate of oxidation due to s t e r i c e f f e c t s i s a t t r i b u t e d to r e s t r i c t i o n s i n the freedom of rotation of the chromate ester, i . e . i n those re-actions i n which acceleration due to s t e r i c e f f e c t s i s observed the rotation of the ester i s r e s t r i c t e d i n such a way that the acceptor oxygen spends more time i n the v i c i n i t y of the hydrogen and thus i n -creases the s t a t i s t i c a l p r o b a b i l i t y that a reaction w i l l occur. Kwart and Francis have successfully 21 applied t h i s mechanism to explain the r e l a t i v e rates of oxidation of a series of alkylated cyclohexanols (57) as well as a large number of 1, 2-diols (24). The main c r i t i c i s m of t h i s mechanism has come from Roc'ek (32). His f i r s t contention that such a mechanism should not give a good log k v.s. H Q p l o t i s not well founded, as has been indicated i n previous discussion. On the other hand, his second objection that one would not expect a reaction which required proton transfer to the chromium part of the ester to be accelerated by protonation of t h i s very part, i s a v a l i d c r i t i c i s m of the mechanism. However, a s l i g h t modification of the mechanism to allow f o r i n t e r n a l hydride ion transfer as opposed to i n t e r n a l proton transfer would overcome t h i s objection. • In addition the observed negative "rho" value f o r the reaction would also be more e a s i l y explained on the basis of such an i n t e r n a l hydride ion transfer mechanism. The work of Kwart, Lord and Corey (24) on the oxidation of 1, 2-diols, as well as that reported by Chang and Westheimer (25) and by Rocek and Westheimer (26) has proven to be very i l l u m i n a t i n g because the observed r e s u l t s indicate very strongly that chromate esters function as intermediates i n t h i s reaction. Rocek (32) was the f i r s t to point out that i f a 1, 2-22 d i o l gave r i s e to a c y c l i c ester, the t o t a l amount of e s t e r i f i e d alcohol would be much greater than in the case of a monohydric alcohol and t h i s should lead to enhancement of the rate. Although Roc'ek i n his i n i t i a l i n v e s t i g a t i o n f a i l e d to f i n d any acceleration i n the rate of oxidation of ethylene g l y c o l as compared to ethanol i t has since been found by Chang and Westheimer (25) that pinacol i s more rapidl y oxidized than pinacol monomethyl ether. In addition i t has been observed that exo-exo-2,3-camphanediol and endo-endo-2,3-camphanediol are both more r e a d i l y oxidized than exo-endo-2,3-camphane-d i o l (24), and that 1,2-cis-cylopentanediol and 1,2-cis-cyclohexanediol are both oxidized f a s t e r than the corre-sponding trans-compounds as i s the case for c i s and trans-l,2-dimethyl-l,2-cyclopentanediol (26). In each instance i t has been concluded that the c i s - d i o l s are oxidized at a f a s t e r rate because they can form c y c l i c chromate esters, while the geometry of the corresponding trans compounds prevents formation of such c y c l i c esters. In conclusion, i t can be seen from a review of the l i t e r a t u r e concerning chromic acid oxidations of alcohols that there exists four d i f f e r e n t suggested mechanisms fo r t h i s reaction: (i) The o r i g i n a l Westheimer ester mechanism which assumed the formation of a chromate ester i n a rapid 23 r e v e r s i b l e equilibrium step followed by proton transfer to an external nucleophile such as water i n the r a t e -determining step (37). HCrO^" + R2CH0H ~ — * R 2CHCr0 4~ + H 20 R 2CHCr0 4° + H 20 >RgC - 0 + H^O* + C r I V ( i i ) A hydride ion transfer mechanism proposed by Rocek and Krupicka (31); R ? H OH R \ s<S ^ 1 + / \ xv C ^Cr > C = 0 + C r i V This mechanism has been shown to be quite improbable by studies i n D 20 (13) and by the r e s u l t s obtained i n the oxidation of 1,,2-diols which involve p a r t i c i p a t i o n of the chromate ester (24). ( i i i ) A modification of the ester mechanism which postulates a c y c l i c intramolecular proton transfer i n the ester molecule (5#); HCrO^" + R2CH0 -=i R 2CHCr0 4" + H20 It .Cr—OH R A > \ - 0 ^ C r 1 7 / X C o / R H R 24 (iv) A c y c l i c intramolecular hydride transfer which L i t t l e r (51) has very b r i e f l y indicated as a p o s s i b i l i t y . 3. PRIMARY KINETIC ISOTOPE EFFECTS Primary deuterium isotope e f f e c t s have often been used as an aid in the elucidation of the mechanism of chemical reactions (59). Theoretical predictions i n -dicate that i f the rate c o n t r o l l i n g step of a reaction involves f i s s i o n of a carbon-hydrogen bond the rate of the reaction should be slower i n cases where one of the heavier isotopes of hydrogen (deuterium or tritium) forms the bond than when the l i g h t e s t isotope (protium) forms the bond (60). (While t h i s reasoning can as eas i l y be applied to cases where hydrogen i s bonded to elements other than carbon, i t w i l l be s u f f i c i e n t to consider only carbon-hydrogen bonds i n t h i s discussion.) A t y p i c a l example of the use of thi s e f f e c t as a diagnostic t o o l i n mechanism studies can be found i n the chromic acid oxidation of secondary alcohols. One of the e a r l i e s t postulated mechanisms of t h i s reaction suggested that the rate-determining step involved IV 25 hydride transfer from the hydroxyl group of the alcohol (61). Although there existed at the time some evidence f o r the p l a u s i b i l i t y of such a reaction i t was shown to be incorrect when Westheimer and Nicolaides (37) observed that the reaction exhibited a primary deuterium isotope e f f e c t . This indicated that the rate-determining step involved carbon-hydrogen bond f i s s i o n and led Westheimer to propose the previously discussed ester mechanism f o r t h i s reaction. As has been indicated by B e l l (62a) and others there are two main reasons why reactions involving carbon-deuterium bond f i s s i o n should be slower than corresponding reactions involving carbon-protium bonds; (i) the r e l a t i v e l y large mass difference between these two isotopes and ( i i ) the p o s s i b i l i t y of deviations from c l a s s i c a l mechanics leading to quantum mechanical tunnelling. The mass difference manifests i t s e l f mainly i n a difference i n zero-point energies f o r the C-H and C-D bonds, i . e . the lowest energy l e v e l f o r any bond corresponds to 1/2 ht> , where h i s Planck's constant and ti i s the v i b r a t i o n a l stretching frequency of the bond in question. The v i b r a t i o n a l stretching frequency i s i n turn inversely dependent on the square root of 26 the reduced mass according to the Hooke's law expression, 7) = 1 l""k where k i s the force constant. For 2 TJ c J u examplej i f one considers two molecules A-H and A=D the r a t i o of the v i b r a t i o n a l stretching frequencies f o r the protium and deuterium bond "is given by mD(mA+mH) mH(mA+mD) For large molecular weight molecules m^  i s so much greater than m^ and m^  that the terms i n the brackets cancel leaving one with the r a t i o ^A-H . mn ST -mH Hence, from the observation that the carbon-hydrogen stretching frequency f o r the oC-hydrogen i n 2-propanol i s at 2900 cm?'*" (63) one would predict that the carbon-deuterium frequency i n 2-deuterio - 2 -propanol would be at 2900 x 1/1.414 = 2050 cm"1. This agrees favorably with the observed value of 2100 cm~1. (64) and gives a difference i n the zero point energies for these two bonds of 1.15 kcal/mole. The r a t i o k^/k^ can then be calculated by appl i c a t i o n of the theory of absolute rates which predicts that the s p e c i f i c rate of a reaction should be given by the following expression (65a); 27 k = KT J * _ -E Q/RT * h F x F y e 0 Where K i s Baltzmans constant, T i s the absolute temperature, h i s Planers constant, F # i s the p a r t i t i o n function f o r the activated complex, F and F are the x y p a r t i t i o n functions f o r the reacting molecules, R i s the gas constant and E Q i s the zero point energy of the bond being broken i n the reaction,, I f one neglects possible differences i n the energies of the correspond-ing t r a n s i t i o n states the r a t i o k H / k n should be e A E ^ ^ e Thus, f o r the case of oxidation of isopropyl alcohol where AE was calculated to be 1.15 kcal/mole t h i s predicts a primary k i n e t i c isotope e f f e c t of 6.9 at 25°C. This i s i n agreement with value observed experimentally (41). Since temperature appears i n the denominator of the exponential used to calculate the isotope e f f e c t i t can be seen that smaller effects should be observed at higher temperatures while the converse should be true at lower temperatures. The maximal isotope e f f e c t s predicted f o r various temperatures are l i s t e d i n Table I. 28 TABLE I VARIATION IN THEORETICAL ISOTOPE EFFECTS WITH TEMPERATURE Temperature (°C) o 10 25 8.34 50 7.14 75 6.75 100 5.56 125 5.00 150 4.54 * These values were obtained from t h e i r reciprocals l i s t e d i n reference (66). It i s , however, possible to observe e f f e c t s to be much smaller than those l i s t e d i n Table I, even when the ratei-determining step of a reaction involves carbon-hydrogen bond f i s s i o n , since a l i n e a r t r a n s i t i o n state such as may have a "symmetrical'* mode of vi b r a t i o n involving movement of H as well as A and B (67). I f the hydrogen moves the deuterio t r a n s i t i o n state w i l l have a lower zero point energy than the corresponding protio t r a n s i t i o n state and the o v e r - a l l r e s u l t w i l l be a A—H---B 29 decrease i n the observed isotope e f f e c t . Only i n the case where the so-called "symmetrical" v i b r a t i o n in= volves no motion of the hydrogen atom w i l l the maximal predicted isotope e f f e c t be observed,, Thornton has recently suggested an application of these ideas to show whether the t r a n s i t i o n state i s product-like or re a c t a n t - l i k e (68). It has also been suggested by Melander (66) that an a l t e r n a t i v e explanation of small isotope e f f e c t s may be found i n a consideration of the carbon-hydrogen bending frequencies. I f the bending frequencies were to increase on passing to the t r a n s i t i o n state the isotope e f f e c t due to differences i n the zero point energies of the stretching frequencies would be p a r t i a l l y cancelled. In addition i f the bending frequencies were decreased i n the t r a n s i t i o n state i t would lead to the observation of unusually large isotope e f f e c t s . There have, i n f a c t , been several reports of reactions e x h i b i t i n g isotope e f f e c t s of a magnitude greater than the values l i s t e d i n Table I. These i n -clude the chromic acid oxidation of phenyl-tert-butylcarbinol i n 86.5$ acetic acid ( k H / k D =8.5 to 14.4 at 0°C) (44), the reaction of methyl r a d i c a l s with toluene ( k H / k n = 8.9 at 122°C) (69), ( k H A D = 9.9 at 120°C) (59), the reaction of l i t h i u m cyclohexylamine 30 with ethylbenzene ( k H / k D = 12- 3 at 49.9°C) (70), the abstraction of hydrogen atoms from acetic -oC-d^ ac i d by methyl r a d i c a l s ( k H / k D = 8.3 at 120°C) (59), the abstraction of hydrogen atoms from cyclohexane-d^ by methyl r a d i c a l s ( k H / k D » 6.9 at 120°C) (59), and the oxidation of aryltrifluoromethylcarbinols with basic permanganate (k^/k D = 16 at 25°C) (71). The largest primary deuterium isotope e f f e c t s reported to date have been observed i n the oxidation of f l u o r i n a t e d alcohols by basic permanganate (71, 72). For t h i s reaction values of k^/k D from 12.0 to 19.5, depending on the substrate oxidized and on the condi-tions of the reaction, have been observed at 25°C. While tu n n e l l i n g has been ci t e d as a possible explanation f o r the large e f f e c t s i t has also been suggested that they might be due to a loss of the bending carbon-hydrogen motion i n the t r a n s i t i o n state, or perhaps to an unusual consecutive process which produces a cumulative isotope e f f e c t . It i s of i n t e r e s t , however, to note that the value of 16.1-.1 obtained f o r the permanganate oxidation of m~bromophenyltrifluoromethylcarbinol, p _ - t o l y l t r i -fluoromethylcarbinol and phenyltrifluoromethylcarbinol at a pH of 13.3 i s in good,agreement with the theoret-i c a l isotope e f f e c t of 17.2 predicted for the case 31 where the carbon-hydrogen bond loses both stretching and bending frequencies i n the t r a n s i t i o n state (66) „ Further Stewart and Van der Linden (71) have pointed out that the rate-determining step of the reaction involves i n t e r a c t i o n of two anions and that t h i s could lead to a t r a n s i t i o n state containing two widely separated molecules. In the involvement of such an extended t r a n s i t i o n state may l i e an ex° planation of the loss of bending frequencies i n t h i s reaction. Another possible explanation of unusually large deuterium isotope e f f e c t s i s the operation of a phenomenon commonly c a l l e d "proton tunnelling". In f a c t , B e l l (62b) has stated that, "In any k i n e t i c treatment (of isotope effects) which i s s u f f i c i e n t l y refined to take into account the zero-point energy of the t r a n s i t i o n state i t i s not j u s t i f i a b l e to neglect the tunnel e f f e c t . " This e f f e c t a rises from the quantum mechanical conclusion that a l l p a r t i c l e s have, besides a corpuscular nature, a wave-like nature with a De Broglie wavelength of X = h/mv where m i s the mass of the p a r t i c l e and v i s i t s v e l o c i t y . Since the mass of the p a r t i c l e appears i n the denomi-nator of t h i s expression i t i s to be expected that the e f f e c t w i l l be more pronounced f o r l i g h t e r 32 p a r t i c l e s ? i n p a r t i c u l a r i t should be greater f o r protons than f o r deuterons. Further, i t i s well known that p a r t i c l e s with a wave-like nature do not obey the laws of c l a s s i c a l mechanics and that there i s a f i n i t e p o s s i b i l i t y that a quantum mechanical p a r t i c l e w i l l be found i n regions, which according to c l a s s i c a l theory, are completely forbidden (73). Thus c l a s s i c a l l y the only way that a chemical reaction can occur i s f o r a p a r t i c l e to hurdle the po t e n t i a l energy b a r r i e r . However, quantum mechanical consid-erations lead to the conclusion that i t i s possible f o r a p a r t i c l e of low mass to also penetrate the ba r r i e r . Since the mass of a proton i s only one=half as great as that of a deuteron t h i s e f f e c t could also contribute to the isotope e f f e c t of a reaction i n -volving proton transfer. (The same considerations would also apply to reactions involving hydride ion, or hydrogen atom transfers.) Evidence of a contribu-t i o n from t h i s source to the observed isotope e f f e c t has come from the work of B e l l , Fendley and Hulett (74, 75) and Shiner and Smith (76). Before concluding t h i s discussion of primary k i n e t i c isotope e f f e c t s two other aspects of t h i s f i e l d should be mentioned b r i e f l y ; ( i ) the v a r i a t i o n of isotope e f f e c t s with rates i n a series of analogous 33 reactions and ( i i ) the suggested use of substituent e f f e c t s on isotope e f f e c t s to d i s t i n g u i s h between proton and hydride transfer reactions ( 7 7 ) . With respect to the former feature i t has been observed i n at l e a s t three reactions (the side chain halogenation of toluene ( 7 8 ) , the chromic acid oxidation of some methyl carbinols ( 3 9 ) , and the decarboxylation of #-keto acids i n benzene (79) ) that isotope e f f e c t s become larger with decreasing reaction rates. The reason f o r t h i s r e l a t i o n s h i p would seem to be that the reactions which have the highest a c t i v a t i o n energy also have the most symmetrical t r a n s i t i o n state. For example, i n the c h l o r i n a t i o n of substituted toluenes studied by Wiberg and Slaugh (78) the t r a n s i t i o n state i s probably l i n e a r CI H C and since the isotope e f f e c t i s smaller than seven i n a l l cases the so-called "symmetrical* v i b r a t i o n must involve considerable motion of the hydrogen. However since the isotope e f f e c t varies from 3.22 to 4 . 8 6 to 5 .08 f o r the reaction of N-Bromosuccinimide with £-methoxytoluene-d^, toluene-d-^, and p_-chloro-toluene-d-p respectively, i t has been suggested that the t r a n s i t i o n state i s modified as one moves along t h i s series to involve l e s s and less movement of the 34 hydrogen i n the "symmetrical" v i b r a t i o n ( 6 7 ) . The second suggestion that substituent e f f e c t s on isotope e f f e c t s can be used to d i s t i n g u i s h between proton and hydride transfer reactions has been ad-vanced by Swain and coworkers ( 77 , 7 9 ) . They have presented t h e o r e t i c a l reasons f o r suspecting that the bonding i n the t r a n s i t i o n state of a proton transfer reaction i s "weaker, longer, more io n i c and more p o l a r i z a b l e " than the bonding i n the t r a n s i t i o n state of a hydride ion transfer. A somewhat over-s i m p l i f i e d statement of the t h e o r e t i c a l considerations that led these authors to propose t h i s difference i n t r a n s i t i o n state bonding f o r the two reactions i s that i n a hydride transfer between electron d e f i c i e n t centers the two electrons of the hydride can be accommodated i n the lowest c - o r b i t a l , while f o r the case of proton transfer between two electron r i c h Ij'ceBfeers an antibonding o r b i t a l must also be occupied. This leads, f o r proton transfer reactions, to a more polarizable t r a n s i t i o n state which should be'more sensit i v e to substituent e f f e c t s . As supporting experimental evidence these authors c i t e the carbon-hydrogen and oxygen-hydrogen isotope e f f e c t s f o r the oxidation of 2-propanol and l-fluoro - 2-propanol by bromine. The carbon-hydrogen isotope e f f e c t i s 35 almost the same f o r both alcohols while the oxygen-hydrogen isotope e f f e c t i s 38$ greater f o r l - f l u o r o - 2 -propanol, suggesting'that the mechanism involves hydride transfer from the carbon with simultaneous proton abstraction from the oxygen (77). Also they present data showing that the isotope e f f e c t ob-served in the decarboxylation of several substituted benzoylacetic acids i s variable, i n d i c a t i n g that the rate-determining step i n t h i s reaction involves proton and not hydride transfer. Although t h i s theory could eventually prove to be a useful diagnostic t o o l one cannot make extensive use of i t u n t i l more experimental evidence i n i t s support has been presented. Further, i t should be noted that L i t t l e r (51) has recently offered a b r i e f c r i t i c i s m of the mechanism suggested by Swain and coworkers f o r the bromine oxidation of alcohols„ and that the r e s u l t s of the decarboxylation experiments could as e a s i l y be incorporated into the more general theory which predicts a p a r a l l e l between the magnitude of the isotope e f f e c t s and the a c t i v a t i o n energies in an analogous seri e s of reactions (59). 36 II SCOPE OF THE PRESENT INVESTIGATION The discovery by Stewart and Van der Linden (71) that unusually large isotope e f f e c t s were observed when a group of aryltrifluoromethylcarbinols were oxidized by permanganate i n basic medium gave the i n i t i a l impetus to t h i s investigation,, In l i g h t of t h i s discovery i t was decided to determine i f extra-ordinary isotope e f f e c t s were also observed when these alcohols were oxidized by chromic acid. An i n i t i a l i n v e s t i g a t i o n revealed that although the isotope e f f e c t s observed i n the chromic acid oxidations were not as large as those reported f o r permanganate oxidations the magnitude varied considerably when d i f f e r e n t substituents were introduced i n t o the aromatic nucleus. In order to study t h i s phenomena in more d e t a i l i t was decided to synthesize some add i t i o n a l fluorinated secondary alcohols, both aromatic and a l i p h a t i c . Isotope e f f e c t s f o r the oxidation of various fluoroalcohols i n 77.2$ acetic acid and 50.1$ s u l f u r i c acid were investigated, a l -though due to s o l u b i l i t y d i f f i c u l t i e s with s u l f u r i c acid only a l i p h a t i c alcohols could be used i n t h i s medium. During the investigation of these isotope e f f e c t s i t became apparent that some of the f l u o r o -3 7 alcohols possessed properties such that they could be used e f f e c t i v e l y to t e s t certain postulates concerning the mechanism of chromic a c i d oxidations ? p a r t i c u l a r l y i n very a c i d i c regions. Hence a study of the oxidation mechanism i n concentrated s u l f u r i c acid solutions was undertaken, Further f the publication by Klaning and Symons ( 4 3 ) of a large number of equilibrium constants f o r equilibriums of the type ROH + HCrO^~ ROCrO^" + H20 led to a decision to determine the "true" rate constant (k 2 = k o b s/Ke) f o r the chromic acid oxidation of a number of primary a l i p h a t i c alcohols. F i n a l l y , i t was observed that i f one plotted log kobs a S a i n s t H 0 f o r t n e oxidation of isopropyl alcohol in perchloric acid solutions of various concentrations a sharp change i n the slope of the p l o t was observed at approximately the same point that B a i l e y v Carrington ? Lott and Symons (SO) had determined to be the pKa of H 2CrO^ in t h i s medium. Further i n v e s t i g a t i o n revealed that the pKa of chromic acid varied considerably when i t was determined i n d i f f e r e n t mineral acids and that i n each case the change in slope of the plot occurred a-t approximately the same point as the observed pKa. The detailed r e s u l t s and experimental procedures along with a discussion of the implications that can 38 be derived from the r e s u l t s of t h i s investigation w i l l be presented i n the following pages of th i s thesis„ 39 III EXPERIMENTAL 1. MATERIALS (i) Alcohols: A l l l i q u i d fluoro and chloro-alcohols were p u r i f i e d by use of an Aerograph Model A-90P Chromatograph before use i n k i n e t i c experiments,, A ten foot Ucon Polar column was found to be most s u i t -able for these compounds. The alcohols used i n t h i s i n v e s t i g a t i o n conveniently f i t into three groups; those avai l a b l e commercially, those previously syn-thesized i n t h i s laboratory and those prepared during the course of t h i s i n v e s t i g a t i o n . The following alcohols were obtained from commercial sources: isopropyl a l c o h o l t 1» 3-dichloro~ 2-propanol, l-chloro-2propanol i ) ethanol, 2-methoxy-ethanolp 2-chloroethanol, methanol, n-propanol (Eastman Organic Chemical Company) 1 , 3-difluoro-2-propanol f 1 Pl„l-trifluoro-2-propanol, l ? l f l -trifluoro-2-methyl-2=propanol (Columbia Organic Chemical Co.), 1,l sl , 3 » 3 ? 3-hexafluoro-2-propanol (obtained as a g i f t from S. Andreades, E„I. du Pont de Nemours and Co.„ Research Department Wilmington ? Delaware). The following fluoro alcohols had previously been prepared i n these laboratories (81, 82): p_-40 methoxyphenyltrifluofomethylcarbino1, p _ - t o l y l t r i -fluoromethylcarbinol, £~tolyltrifluoromethylcarbinol-cc»dj phenyltrifluoromethylcarbinol, p h e n y l t r i f l u o r o -methylcarbinol-oc ~d, m-bromophenyltrifluoromethyl-c a r b i n o l 9 m-bromophenyltrifluoromethylcarbinol-oc-d and m»nitroptenyltrifluoromethylcarbinol. During the present i n v e s t i g a t i o n several of these compounds were re-synthesized and good agreement with a l l the reported physical constants was found except i n the case of m«= nitrophenyltrifluoromethylcarbinol where the melting point was found to be 51.2-52.2°C instead of the reported 47-48°C. The preparations of other alcohols which were synthesized e s p e c i a l l y f o r t h i s i n v e s t i g a t i o n are given below. 2-Propanol»2-d. Following the procedure of Leo and Westheimer (41) dry acetone (5 gm., 0.086 moles) i n 20 ml. of dry ether was added dropwise to a s t i r r e d s o l u t i o n of li t h i u m aluminum deuteride (1.05 gnu, 0.025 moles) i n 5 0 ml of dry ether at -80°C„ The mixture was allowed to warm slowly to room tempera-ture and the complex was decomposed with d i l . s u l f u r i c a c i d . The heterogeneous reaction mixture was then d i s t i l l e d through a 12" glass column and the azeotrope of 2-propanol-2 -d and water which d i s t i l l e d at 70-90°C was c o l l e c t e d . This was dissolved i n ether and dried over anhydrous magnesium s u l f a t e . After evaporation of the ether the r e s u l t i n g clear solution was d i s -t i l l e d from 0.2 gm. of calcium oxide to give 2-propanol-2-d (3.1 gm., 0.052 mole, 60%) b.p. = 81-82.3°C. l-Fluoro-2-propanol. Chloroacetone (41.2 gm., 0.5 mole) was added dropwise to a well s t i r r e d mixture of potassium hydrogen f l u o r i d e (87 gm., 1 mole) and 60 ml. of diethylene g l y c o l i n a 250 ml. three necked f l a s k f i t t e d with a dropping funnel, a s t i r r e r and a s t i l l head; the f l a s k at the time being immersed i n an o i l bath at 160°C. As the addition proceeded a col o r l e s s l i q u i d d i s t i l l e d o f f at 80-95°C. This was col l e c t e d , dried over anhydrous magnesium sulfate and r e d i s t i l l e d through a 12" glass column to give fore run (3 gm.) b.p.50-75°C, fluoroacetone (9.9 gm., 0.13 mole, 26$) b.p. - 75-80°C, v ( l i q u i d ) 1730 c n f ^ C = 0) and unreacted chloroacetone (32.6 gm.) b.p. 80-120°C. Lithium aluminum hydride (1 gm., 0.026 mole) i n 50 ml. of ether was added dropwise to a s t i r r e d s olution of fluoroacetone (5.0 gm., 0.068 mole) i n 20 ml. of ether at -80°C. The mixture was allowed to warm slowly to room temperature and f i n a l l y re-fluxed gently f o r 30 min., a f t e r which the complex 42 was decomposed with d i l . s u l f u r i c a c i d . The aqueous phase was separated and extracted twice with 20 ml. portions of ether. These extracts were added to the ethereal phase and dried over anhydrous magnesium s u l f a t e . Evaporation of the ether gave a dark o i l which on d i s t i l l a t i o n yielded l-fluoro - 2-propanol (1 gm., 0.013 mole, 19$) b.p. = 103 - 104°C } ( l i t . (83) 103-105°C) l-Fluoro ° 2-propanol - 2-d. Fluoroacetone was reduced with l i t h i u m aluminium deuteride i n the manner de-scribed above f o r the preparation of l-fluoro= 2 -propanol. Both compounds had exactly the same b o i l i n g point and gave i d e n t i c a l retention times on the vapor phase chromatograph. 1,3-Dichloro -2°propanol -2-do l,3°-Dichloro=2-propanol was oxidized by chromic acid i n the manner described by Conant and Quayle (84). The crude product,, 1,,3~ dichloroacetone, was extracted from the aqueous solution with ether, dried over anhydrous magnesium sulfate and d i s t i l l e d under reduced pressure (b.p. = 43°/0.6 mm.). One gram of t h i s material was d i s -solved i n ether and p u r i f i e d by vapor phase chromatography to insure that i t would be completely free of the s t a r t i n g alcohol. The melting point of the f i n a l product was 45-46°C (lit ( 8 5 a ) 4 5°C). This 43 pure ketone was reduced i n ether with a s l i g h t excess of l i t h i u m aluminum deuteride to give 1,3-dichloro-2-propanol-2 -d i n 90$ y i e l d . The product was characterized by comparing i t s retention time on a vapor phase chromatography column with that of 1,3°° dichloro-2-propanol under i d e n t i c a l conditions and through a comparison of the nuclear magnetic resonance spectra of these two compounds. 1- Chloro-3»fluoro-2-propanol. A s t i r r e d mixture of epichlorohydrin (250 gm.), potassium hydrogen f l u o r i d e (235 gm.) and diethylene g l y c o l (250 gm.) was refluxed f o r four hours. The crude product was then d i s t i l l e d under reduced pressure (b.p. 50-120/40 m m0) and then r e d i s t i l l e d through a 12" column at standard pressure. The f r a c t i o n d i s t i l l i n g at 140-155°C was col l e c t e d and a small portion was p u r i f i e d by vapor phase chroma-tography. n D 2 0 = 1.4320(lit.(83 ) n D 2 g = 1.4269). The nuclear magnetic resonance spectra of t h i s product was also consistent with the assumed structure. 1 n3°Difluoro-2-propanol-2-d. Commercial l , 3 - d i f l u o r o ~ 2- propanol was oxidized to 1,3-difluoroacetone with chromic ac i d i n the manner described by Bergmann and Cohen (83). The crude product was continuously ex-tracted from the reaction mixture with ether, dried over anhydrous magnesium sulphate and d i s t i l l e d under 44 reduced pressure. 7) at 1750 cm" . One gram of t h i s material was p u r i f i e d by vapor phase chromatography and reduced i n ether with a s l i g h t excess of l i t h i u m aluminum deuteride to give 1 r 3-difluoro-2=propanol i n 70$ y i e l d . The product was characterized by comparing i t s retention time on a vapor phase chromatography column with that of 1,3~difluoro~2»propanol under i d e n t i c a l conditions and through a comparison of the nuclear magnetic resonance spectra of these two compounds. 1,1,1.3 3 i,3"Hexafluoro°2°propanol°2-d. Ten gm0 of potassium dichromate was dissolved i n 500 ml. of 97$ s u l f u r i c a c i d contained in a three necked f l a s k f i t t e d with a dropping funnel, a r e f l u x condenser, a nitrogen bubbler and a magnetic s t i r r e r . The top of the r e f l u x condenser was connected by means of rubber tubing to four traps. The f i r s t and fourth traps i n the series were cooled i n l i q u i d nitrogen while the second and t h i r d were cooled with a mixture of dry i c e i n acetone<> The second and t h i r d traps also contained an excess of li t h i u m aluminum deuteride dissolved i n dry ether. The solution was s t i r r e d continuously throughout the preparation and a gentle flow of nitrogen through the system was maintained. l,l,lj3 r 3 f , 3-Hexafluoro-2-propanol (5.0 gm., 0.030 mole) 45 was then added dropwise to the chromic acid solution and,the col o r l e s s gas which came o f f was c o l l e c t e d i n the f i r s t trap. After a l l the alcohol had been added the contents of the f l a s k were warmed to 60°C f o r 2 hrs. When the contents of the f i r s t trap were then allowed to warm slowly to room temperature the gas (1,1,1,3,3,3-hexafluoroacetone) bubbled slowly into the l i t h i u m aluminum deuteride solution i n traps two and three. Any gas not reacting with the l i t h i u m aluminum deuteride was collected i n trap four and recycled. F i n a l l y , the ethereal solutions i n traps two and three were allowed to warm slowly to room temperature and the complex was decomposed with dil„ s u l f u r i c a c i d . The aqueous solution was separated and extracted continuously with ether f o r 48 hours. The ethereal solutions were then a l l c o l l e c t e d , dried over anhydrous magnesium sulf a t e and d i s t i l l e d through a 12 n column. Aside from a large amount of ether the main f r a c t i o n c o l l e c t e d was 1,1,1,3,3,3Thexafluoro-2-propanol-2-d (3.2 gm., 0.019 mole, 63$). This compound had exactly the same retention time on a vapor phase chromatography column as 1,1,1,3,3,3-hexafluoro-2-propanol supplied by S. Andreades of E.I. du Pont de Nemours and Co. and nuclear magnetic spectroscopy indicated that i t contained l e s s than 4$ protium 46 alcohol. The properties of t h i s compound correspond cl o s e l y with those reported by Dr. M. Mocek who had previously prepared 1,1,1,3,3,3-hexafluoro~2=>propanol~ 2->d by a s i m i l a r but independent method i n these laboratories (72). m-Methylphenyltrifluoromethylcarbinol. Commercial m°Bromotoluene (Eastman Organic Chemicals)(49 gm., 0.286 moles) dissolved i n 40 ml. of dry ether was added to magnesium turnings (7 gm., 0.286 moles) under anhydrous conditions. After the reaction has been started by warming the mixture s l i g h t l y 70 ml. of dry ether was added and at the completion of the reaction a further 150 ml. of ether was added. Then t r i f l u o r o a c e t i c a c i d (10.5 gm., 0.092 moles) i n 25 ml. of ether was added dropwise over a period of two hours. After r e f l u x i n g f o r one hour the complex was decomposed i n 10$ s u l f u r i c a c i d . The product (m-tolyl t r i f l u o r o -methyl ketone) was extracted from the aqueous solution with ether, dried over anhydrous magnesium sulfate and d i s t i l l e d . B.p. = 58°C/8 m m. Y i e l d ; 8.5 gm., 0.045 moles, 49$. "Z^ m a x. ( l i q ° ) 1 7 1 V c m " , 1 ( C = 0 ) ° This ketone (6.0 gm., 0.032 moles) was dissolved i n 20 ml. of dry ether and lithium aluminum hydride (O .38 gm., 0.01 moles) in 20 ml. of dry ether was added slowly under anhydrous conditions. After 47 r e f l u x i n g f o r one hour the complex was decomposed with d i l . s u l f u r i c a c i d and the product was extracted with ether, washed with water and dried over anhydrous magnesium s u l f a t e . D i s t i l l a t i o n yielded 3.5 gm (0.0184 moles, 57$) of m-tolyltrifluoromethylcarbinolo B.p. - 95°C/14 m m. (Founds C, 57.0$; H, 3.6$. C 9H ?OF 3 required C, 57.4$; H, 3.7$.) £-tert-Butylphenyltrifluoromethylcarbinol. Starting from £-ter_t-butylbroiaobenzene (b.p. = 104-107°C/14 m m„ , l i t . ( 8 7 ) 103/10 m m.) p - t e r t - b u t y l p h e n y l t r i f l u o r o -methylcarbinol (m.p. = 75.0 - 75.5°C) was prepared by the same procedure used to synthezise m-methylphenyl-trifl u o r o m e t h y l c a r b i n o l . (Found: C, 61.6$; H, 6.13$. C12 H15 O F3 r e < l u i r e s : c> 62.0$; H, 6.52$.) B-Naphthyltrifluoromethylcarbinol. Magnesium (2„4 gm., 0.1 moles) was placed i n a flame dried three necked f l a s k f i t t e d with two dropping funnels, a r e f l u x condenser and a magnetic s t i r r e r . Methyl iodide (0.5 gm) dissolved i n 10 ml. of dry ether was added through the f i r s t funnel and a f t e r the reaction had started fi-bromonaphthylene (Eastman Organic Chemicals) (13 gm.j O.O63 moles) dissolved i n 50 ml. of dry ether was added slowly. The mixture was refluxed f o r one-ha l f an hour a f t e r which 100 ml. of dry ether was added. This was followed by the dropwise addition of 48 t r i f l u o r o a c e t i c acid (1.82 gm., 0.016 moles) dissolved i n 20 ml. of dry ether. After the addition was complete the solution was refluxed f o r two hours and then de-composed with cold d i l . s u l f u r i c a c i d , extracted with ether and dried over anhydrous magnesium s u l f a t e . The product (0-naphthyl trifluoromethyl ketone) was d i s t i l l e d under reduced pressure (b.p. = 88°C/1.1 m m.) and c r y s t a l l i z e d from low b o i l i n g petroleum ether. %) (CHC13)1706 cm"1(C - 0). Yield; 5.0 gm., 0.013 moles, 80$, This ketone (2.5 gm., 0.011 mole) was dissolved i n 50 ml. of dry ether and added slowly to ethereal lithiu m aluminum hydride (0.25 gm,, 0.0066 moles) under anhydrous conditions. The viscous solution was--refluxed f o r one hour, decomposed with d i l . s u l f u r i c a c i d and extracted with ether. The ethereal extracts were washed with d i l u t e sodium bicarbonate, dried over anhydrous magnesium sulfate and evaporated. The re-s u l t i n g o i l was r e c r y s t a l l i z e d from petroleum ether y i e l d i n g a f i n e white c r y s t a l l i n e compound. M.p = 83-84°C. A mixed melting point with naphthylene gave a depression of 20°C. (Found: C, 63.4$; H, 3.64$. Required f o r C 1 2H 9OF 3; C, 63.76$; H, 4.01$.) £ Chlorophenyltrifluoromethylcarbinol. Starting from l-bromo-4-chlorobenzene (Eastman Organic Chemicals) 49 £-chlorophenyltrifluoromethylcarbinol (b.p. = 117-119°C/ 22 mm.) was prepared by the same procedure used to synthesize m-tolyltrifluoromethylcarbinol. (Found: C, 44.9$; H, 2.63$. Required for CgHgOF^Cl: C, 45.6$ H, 2.87$). m-Nitrophenyltrifluoromethylcarbinol-oc -d. m-Nitrophenyl trifluoromethyl ketone was prepared by the method re-ported by Van der Linden (81). This ketone was reduced, as described below, with l i t h i u m diethoxyaluminodeuteride using the procedure developed by Brown and Tsukamoto (88). Ethanol (1.84 gm., 0.040 moles) i n 25 ml. of ether was slowly added to a s t i r r e d solution of l i t h i u m aluminum deuteride (0.87 gm., 0.0201 moles) dissolved i n 50 ml. of dry ether at -80°C. This entire solution was then added slowly to a s t i r r e d solution of m-nitrophenyl trifluoromethyl ketone (9.05 gm., 0.041 moles) i n 50 ml. of dry ether at -80°C. After the addition which took 1.5 hours the mixture was allowed to warm slowly to room temperature and the complex was decomposed with cold d i l . s u l f u r i c a c i d . The aqueous solution was extracted with ether; the ethereal extract was dried over anhydrous magnesium sulfate and then evaporated leaving a brown o i l which d i s t i l l e d at 122°C/0.4 mm. This product was re-50 c r y s t a l l i z e d several times from benzene and petroleum ether to give pale yellow c r y s t a l s . M.p. = 51.5-52.5°C. (Found: C„ 43.1$; H, 2.00$, N, 6.28$. Required f o r CgH 5DN0 3F 3: C, 43.2$; H, 2.72$, N, 6.32$.) 3 fl5°Dinitrophenyltrifluoromethylcarbinol. Fuming n i t r i c acid (Sp. gr. 1.15) was added dropwise to a s t i r r e d solution of trifluoroacetophenone (Columbia Organic Chemicals) (38.4 gm., 0.216 moles) i n 150 ml. of concentrated s u l f u r i c a c i d . External cooling was applied to maintain a temperature of 60-80°C during the addition. The mixture was then heated on a steam bath f o r four hours. This was followed by cooling and the addition of another 40 ml. of fuming n i t r i c a c i d . The mixture was again heated on a steam bath f o r four hours and f i n a l l y at 130-150°C f o r four hours. The reaction mixture was then poured into crushed ice and extracted with ether. The ethereal extract was dried over anhydrous magnesium sulfate and evaporated leaving a yellow o i l which was d i s t i l l e d under reduced pressure. The d i s t i l l a t i o n yielded 2 gm. fore-run, 10.6 gm. m nitrophenyl trifluoromethyl ketone (b.p. = 90°C/.95 m m.) and 3,5-dinitrophenyl trifluoromethyl ketone (b.p. = 120-127/.95 m m.) (32.4 gm., 0.123 moles, 57$). " l ) m a X o (film) 1733 cm"1 (C - 0). Hydrolysis of t h i s compound with d i l . sodium 51 hydroxide yielded 3»5-dinitrobenzoic acid (m.p = 203-205°C, l i t . ( 8 5 b ) 204-205°C). Attempted r e c r y s t a l l i z a t i o n of the crude 3?5= dinitrophenyl trifluoromethyl ketone from most common solvents such as ether, petroleum ether and chloroform yielded a stable white s o l i d which showed no carbonyl band i n the infrared, but exhibited peaks at 3490 and 3270 cm"1, (m.p. = 85 .0 -$5.5°C) . It was concluded that t h i s compound was 3 , 5-dinitrophenyl t r i f l u o r o -methyl ketone hydrate, (Found: C, 32.52$; H, 2 . 13$ ; N, 9 . 22$ . Required f o r C g H ^ O ^ ; C, 34.01$;H, 1 .78$; N, 9 , 94$ .) The nuclear magnetic resonance spectra of t h i s compound was also consistent with the assignment of t h i s structure. This compound was also e a s i l y hydrolyzed to 3 , 5-dinitrobenzoic acid i n d i l . sodium hydroxide. Dehydration of 3 , 5-dintrophenyl trifluoromethyl ketone hydrate by phosphorous pentoxide at 50°C yielded the corresponding ketone (m.p. 77-78°C) . l ) m a Y (CHC1-)1737 cm"1 (C = 0 ) . (Found: C, 3 5 . 8 $ ; H, 1.28$ ; N , 10 .27$ . Required f o r C g H ^ O ^ : c , 36.4$;H, 1 .14$ ;N , 10 .60$ . ) R e c r y s t a l l i z a t i o n of the crude 3 , 5-dinitrophenyl trifluoromethyl ketone from ethanol yielded a white c r y s t a l l i n e compound (m.p. = 70-71°C) which showed no 52 absorption i n the carbonyl region of the infrared spectra nor any i n the region from 3200 to 3500 cm~^. where the hydrate absorbs. (Found: C, 29.85$; H, 1.53$; N, 10.96$; 0, 26.35$; Ff. 0.00$; M.W. = 243.) This compound was r e s i s t a n t to basic hydrolysis and as yet i t s exact structure has not been determined. 3 , 5-Dinitrophenyltrifluoromethylcarbinol was prepared from crude 3»5-diriitrophenyl trifluoromethyl ketone by reduction with l i t h i u m tri-t-butoxyalumino° hydride i n diglyme (89). t e r t - B u t y l alcohol (12.33 gm., 0.166 moles) dissolved i n 25 ml. of dry diglyme was added dropwise to a s t i r r e d solution of l i t h i u m aluminum hydride (2.11 gm.; 0.055 moles) i n 50 ml. of dry diglyme. This complete solution was then added dropwise to 3 j, 5-dinitrophenyl trifluoromethyl ketone (14.6 gm., 0.056 moles)in 50 ml. of dry diglyme at -80°C. After the addition was complete the solution was warmed slowly to room temperature, decomposed i n d i l u t e s u l f u r i c a c i d and extracted with ether. The ethereal extract was d r i e d over anhydrous magnesium sulfate and evaporated leaving a yellow o i l . D i s t i l l a t i o n of the o i l yielded 3>5-dinitrophenyltrifluoromethylcarbinol (b.p. = 108-174°C/8 mm.), (9.2 gm., 0.0346 moles, 63$). This product was further p u r i f i e d by r e c r y s t a l -53 l i z a t i o n from carbon t e t r a c h l o r i d e . The r e s u l t i n g f i n e white c r y s t a l l i n e compound gave an inf r a r e d spectrum s i m i l a r to m-nitrophenyltrifluoromethyl-carbinol and a nuclear magnetic resonance spectra consistent with the assumed structure. (M.p. = ?6-77°C)„ (Found: C, 36.16$; H, 1.72$, N p 10.54$. Required for C g H ^ O ^ : 36.10$; H, 1.89$; N, 1 0 . 5 3$J 3 , 5-Dinitrophenyltrifluoromethylcarbinol- cC°d. This compound was prepared by the reduction of 3 r 5-dintro-phenyl trifluoromethyl ketone with lithium t r i - t -butoxyaluminodueteride as described above. (M.p. = 76.77°C.) (Found: C, 36.07$; H, 1.78$ f 10.78$. Required f o r CgHgDN^F^: C, 36.00$; H, 1.89$; N ? 10.50$,) The nuclear magnetic resonance spectrum was also consistent with the assignment of t h i s structure. ( i i ) Inorganic Reagents Chromic acid stock solutions were made up by d i s -solving weighed amounts of potassium dichromate (Mallinckrodt, A n a l y t i c a l Reagent Grade, Primary Standard) i n d e f i n i t e volumes of d i s t i l l e d water. S u l f u r i c acid solutions were prepared from commercial 95-98$ s u l f u r i c acid (Fisher C e r t i f i e d Reagent). The concentrations of these solutions were determined by t i t r a t i n g weighed amounts of each solution with 54 standardized sodium hydroxide. Perchloric a c i d solutions were prepared from commercial 70$ perchloric a c i d (B.D.H.). These solutions were also standardized by t i t r a t i o n with sodium hydroxide solutions of known concentration. N i t r i c a c i d solutions were prepared from commercial 73$ n i t r i c a c i d (B.D.H. Micro A n a l y t i c a l Reagent), and standardized with sodium hydroxide. Hydrochloric acid solutions were prepared from commerr c i a l 35.4$ hydrochloric a c i d (B.D.H. Analar), and standardized with sodium hydroxide. Phosphoric acid solutions were prepared from commercial 85$ O-phosphoric acid (Fisher C e r t i f i e d Reagent), and standardized with sodium hydroxide. Sodium Perchlorate. Commercial sodium perchlorate (G. Frederich Smith Chemical Co. Columbus Ohio), was r e c r y s t a l l i z e d several times from water and dried for 10 days at 120°C. Acetic Acid. G l a c i a l acetic a c i d (Nichols Chemical Co.) was p u r i f i e d by d i s t i l l a t i o n under reduced pressure from potassium dichromate and s u l f u r i c acid (12). B.p. = 41°C/25 m m. ( i i i ) Indicators. The indicators used i n t h i s i n -vestigation (p_-nitroaniline, o - n i t r o a n i l i n e , 4-chloro-2 - n i t r o a n i l i n e and p_ -nitrodiphenylamine) were a l l 55 obtained commercially and r e c r y s t a l l i z e d from ethanol to constant melting points. The melting points were a l l found to agree favorably with the values reported i n the l i t e r a t u r e (90). £-Toluic acid was obtained commercially (Eastman Organic Chemicals) and r e c r y s t a l l i z e d several times from hot water (m.p, - 178-179, l i t . ( 8 5 c ) 179.6) 2. KINETIC'METHODS. Oxidation rates were obtained by following the disappearance of C r ^ 1 during the course VI of a reaction. The concentration of Cr was deter-mined either iodometrically or spectrophotometrically. In cases where the alcohol was present i n several f o l d and the f i r s t order rate constants were obtained from the slope. These f i r s t order rate constants were converted to the corresponding second order rate constants by d i v i d i n g through by the concentration of the a l c o h o l . In cases where the alcohol and C r ^ 1 were present in stoichiometric proportions a plot of the re c i p r o c a l of [pr^^] against time yielded the second order rate constant d i r e c t l y . With the exceptions of oxidations i n very concentrated or very d i l u t e solutions good str a i g h t l i n e plots to at lea s t 70$ completion were obtained by both methods f o r a l l rate constants reported. excess a p l o t of against time was prepared 56 (See Figs. 1 and 2, and Table 2.) In very d i l u t e acid concentrations i n i t i a l reaction rates were used since i t was found that the pH increased s l i g h t l y as the re-action proceeded. In very concentrated a c i d (68.5$ h^SO^ or greater) i t was found that the o p t i c a l density of the solution at 349 mu . increased f o r a few minutes a f t e r addition of the alcohol and then decreased exponentially to give a good f i r s t order p l o t . This has been interpreted as evidence f o r ester formation by a slowly established equilibrium p r i o r to the rate-determining step and w i l l be discussed i n d e t a i l l a t e r . Typical k i n e t i c runs were carried out as follows: ( i ) Iodometric method. F i f t y ml. of a 4.035 x 1CT3M ace t i c a c i d solution of phenyltrifluoromethylcarbinol were mixed with 20 ml. of 71$ perchloric a c i d and + o thermostated at 25.0-.1C f o r 10 min. A f a s t delivery pipet was then used to add 4.94 ml. of 0.01361 M potassium dichromate to t h i s solution and 4.92 ml. aliquots were withdrawn at i n t e r v a l s . The aliquots were quenched by adding them to 15 ml. of 5$ sodium bicarbonate solution containing a large excess of potassium iodide. Two ml. of 6M s u l f u r i c acid were added and the solution was set aside to develop f o r three minutes. The I^ formed was then t i t r a t e d with standardized t h i o s u l f a t e using Thyodene as an i n d i c a t o r . 1 0 0 0 T i m e (sec.) 3 0 0 0 4 0 0 0 1 1 I L _ Fig.1 Typ i ca l r a t e p lo ts f o r t h e o x i d a t i o n of a r y l t r i f l u o r o -m e t h y l c a r b i n o l s in 7 7 . 2 % a c e t i c ac id . i - - o - r • . 5 0 T i m e (min) Fig.2 Typ ica l ra te p lo ts f o r the ox ida t i on of a lcoho l in 0 . 4 6 3 M H C I 0 4 . 150 i sop ropy l 03-59 TABLE II DATA FOR PLOTS IN FIGURES 1 AND 2 (a) Oxidation of £=methylphenyltrifluoromethylcarbinol in 77.2$ acetic a c i d . H Q - -2.82, T - 25.0oC. Time (sec.) M i l l i l i t e r s of l / [ c r V : H (1/M) 6.64 x 10-3 M.Na^S.O. 90 3.43 648 165 3.21 683 245 2.92 762 320 2.74 814 400 2.50 888 470 2.40 926 530 2.24 994 610 2.16 1030 670 2.02 1100 760 1.89 1177 850 1.75 1272 960 1.66 1340 (b) Oxidation of phenyltrifluoromethylcarbinol i n 77.2$ ace t i c a c i d . H Q = -2.28, T = 25°C. Time (sec.) M i l l i l i t e r s of 1/ K3r V*l -(1/M) 6.42 x 10-3 M Na 2S 20 3 90 3.92 588 453 3.13 870 60 815 2.65 870 1175 2,30 1005 1525 2,05 1125 1920 1.81 1275 2265 1.64 1410 2650 1.48 1500 3250 1.30 1775 4155 1.10 2100 (c) Oxidation of m^browophenyltrifluoromethylcarbinol i n 77.2$ ac e t i c a c i d . H Q - -2.82, T = 25.0°C. Time (sec.) M i l l i l i t e r s of 1 / f c r 1 ^ (1/M) 6.64 x 10"3 M Na 2S 20 3 L J 70 3.88 574 490 3.58 621 915 3.18 700 1320 2.93 759 1890 2.61 852 2345 2.38 936 2915 2.23 996 3440 2.01 1108 3900 1.88 1185 61 (d) Oxidation of m-nitrophenyltrifluoromethylcarbinol in 77.2$ acetic a c i d . H Q = -2.82, T = 25.0°C. Time (sec.) M i l l i l i t e r of l / | c r V ^ ( 1 / M ) 6.65 x 10-3 M Na 2S 20 3 85 3.99 558 465 3.80 585 1670 3.38 657 3090 2.96 750 4260 2.66 834 5580 2.41 922 6810 2.20 10090 7995 2.00 11100 9120 1.81 12280 10140 1.73 12830 (e) Oxidation T = 25.0°C of isopropyl alcohol i n O.463 M HC10,. 0 r -r , X = 4440 A, [alcohol] = 0.129M. Time (min.) Optical Density 2 + log O.D. Observed Corrected 0.0 .765 .765 1.384 0.5 .756 .756 1.879 1.2 .750 .750 1.875 2.1 .741 .740 1.869 3.0 .738 .737 1.868 5.1 .720 .716 1.855 9.78 .682 .677 1.831 62 17.1 .621 .615 1.789 20.5 .602 .595 1.775 26.7 .559 .550 1.740 33.3 .511 .500 1.699 38.5 .481 .470 1.672 45.9 .440 .426 1.629 58.3 .378 .363 1.560 66.2 .344 .328 1.516 71.7 .323 .306 1.486 80.0 .294 .275 1.439 89.5 .260 .240 1.380 93.9 .248 .228 1.358 106.5 .212 .191 1.281 116.0 .192 .170 1.230 121.8 .178 .155 1.190 133.5 .155 .130 1.111 146.1 .139 .114 1.057 159.8 .118 .093 0.969 (f) Oxidation of 2=propanol-2-d i n 0.463 M HCIO^. T = 23°C, \= 4440A, [alcohol] - 0.127 M. Time (min.) Optical Density 2 + log 0.D Observed Corrected 0 .763 .763 1.883 1 .758 .758 1.880 2 .753 .753 1.877 12 .740 21 .730 31 .728 48 .695 66 .674 81 .657 102 .630 120 .610 151 .579 182 .541 256 .470 287 .450 311 .429 373 .380 396 .365 436 .340 450 .330 632 .235 653 .227 685 .212 .740 1.869 .730 1.863 .728 1.862 .694 1.841 .672 1.827 .655 1.816 .626 1.797 .605 1.782 .574 1.759 .535 1.731 .460 1.663 .440 1.644 .419 1.622 .367 1.565 .351 1.545 .325 1.512 .315 1.498 .217 1.337 .210 1.322 .193 1.286 64 Standardization of the t h i o s u l f a t e was accomplished by-following an i d e n t i c a l procedure with deletion of the alcohol, ( i i ) Spectrophotometric Method, Three ml. of 3.42 x lO'V. V'^7- solution in O.463 M HCIO^ were pipetted + o into a one cm. s i l i c a c e l l and thermostated at 25 .0- .1 C i n the c e l l compartment of a Beckmann model D.U. spectrophotometer, Isopropyl alcohol (.03 ml., 3.92 x 10"°^ moles) was then added from a m i c r o l i t e r syringe f i t t e d with a Chaney adapter. The solution was mixed >' o thoroughly and the o p t i c a l density at 4440 A was determined at i n t e r v a l s . The o p t i c a l density at t h i s ] JI wave length was corrected f o r Cr absorption by use of a previously prepared c a l i b r a t i o n curve and a plot of l o g [ c r ^ against time was prepared. (For VI oxidations u t i l i z i n g a lower concentration of Cr the peak at 349^ could be used and no correction to the observed o p t i c a l density was required.) 3. CONSIDERATION OF THE ACID CHROMATE-BICHROMATE EQUILIBRIUM Westheimer and Novick (1) have established that the a c t i v e chromium species i n the oxidation of isopropyl a l c o h o l i s HCr0^~ and that t h i s species e x i s t s i n equilibrium with bichromate. 2HCr0 4" C r 2 ° 7 = * H 2 ° 65 The magnitude of the equilibrium constant i n d i l u t e aqueous solutions i s such that the amount of chromium i n the bichromate form i s appreciable only at f a i r l y high concentrations and Westheimer has observed that in d i l u t e solutions (0.0005 - 0.005M) the rate i s proportional to the gross C r ^ 1 concentra-t i o n (37). Although no further inv e s t i g a t i o n of t h i s equilibrium was conducted during the present work, an attempt to minimize any deviations a r i s i n g from t h i s e f f e c t was made by using concentrations of C r ^ 1 that were always lower than 0.005M and i n cases where rates were used f o r comparison purposes attempts were made to keep the C r ^ 1 concentration constant throughout the whole s e r i e s . 4. PRODUCT ANALYSIS It i s well known that chromic a c i d oxidizes secondary alcohols smoothly to the corresponding ketones with l i t t l e or no carbon-carbon bond f i s s i o n (37). Since the concentration of reactants was low i n a l l of the k i n e t i c experiments conducted i n t h i s i n v e s t i g a t i o n , i t was found to be extremely d i f f i c u l t to i s o l a t e and i d e n t i f y the products of these reactions. However, when the concentrations of the reactants were increased to a preparative scale i t was found that both the aromatic and the a l i p h a t i c secondary alcohols gave 66 almost quantitative y i e l d s of the corresponding ketones 0 The good second order rate plots obtained also indicate that few side reactions are taking place. In the case where pseudo f i r s t order rate plots were obtained by using large excesses of alcohol i t i s very probable, as Best, L i t t l e r and Waters (19) have indicated, that the i n i t i a l product does not react further with oxidant. This i s so because the reacting alcohol i s always present i n much larger concentrations than any possible products. Oxidation of primary alcohols y i e l d as i n i t i a l products aldehydes, which are l i k e l y to be more re-active than the alcohols themselves. Thus, s t r i c t l y speaking, one can only calculate the rate of oxidation of a primary alcohol i f the rate of oxidation of the corresponding aldehyde i s known. This problem has been considered by Roc"ek (32) who found that i f a large excess of alcohol was used the rate constants calculated from the i n i t i a l slope of the log [ c r ^ v.s. time curve agreed well with the rate obtained from the equation f o r consecutive reactions (rates for the aldehydes being known). In t h i s investigation the rate constants f o r primary alcohols were a l l obtained from the i n i t i a l slope of the l o g v.s. time plot since the rates of oxidation of the corresponding aldehydes under these conditions were 67 not known. 5. DETERMINATION OF PKa VALUES. FOR ^CrO^ For the equilibrium HCrO^" + H + ^ ^ H 2 C r 0 4 the pKa i s defined i n the following way: Ka = a H C r 0 4 ~ * H + = C H C r 0 4 " ' ° H * fHCrO " f H + ~aH2Qr0k C H 2 C r 0 4 f H 2 C r 0 4 P K a = l o g \ ™ k - l o g aH + fHCrO, CHCr0 4~ f H 2 C r 0 4 In d i l u t e solutions the l a s t term on the r i g h t becomes pH. In more concentrated acid solutions i t i s commonly denoted as H_ (91). Q i . e . H m = log H 2 C r 0 4 - pKa CHCr0 4" From t h i s equation i t can be seen that a plot CH 2CrO, of 1 P against H should give a straight l i n e °HCr0 4" of unit slope and an intercept at the pKa. The r a t i o CH 2CrO, p—= =— was determined, following the method of °HCr0 4" Symons et a l (80) by measuring the extinction co-e f f i c i e n t of C r ^ 1 at one or more wave lengths over a 68 wide range of a c i d i t i e s . Then i f 6^ i s the 2 4 extinction c o e f f i c i e n t of a sol u t i o n of H^CrO^ at a pa r t i c u l a r wave length and S H C rQ - i s the extinction 4 c o e f f i c i e n t of a solution of HCrO^~ at the same wave length C H 2 C r 0 4 = ^ H C r 0 4 H C r 0 4 H 2 C r 0 4 where 6 i s the extinction c o e f f i c i e n t of a solution containing both F^CrO^ and HCrO^" (91). The p a r t i c u l a r wavelength at which measurements were made was chosen from a comparison of the spectra of the two species such that a maximum change i n the extinction c o e f f i c i e n t was obtained. For perchloric a c i d , s u l f u r i c acid and hydrochloric acid solutions a plot of ^HCrO " R !og 4 against H was then prepared and the pKa was obtained from the zero intercept. (See Figs. 3, 4 and 5 . ) For solutions of phosphoric acid and n i t r i c acid where no H = functions are a v a i l a b l e a plot of 6HCrO " - e log 4 6 " V r 0 4 against H Q was used f o r the pKa determination. (See Figs. 6 and 7). I t i s to be expected that the use of H Q instead of H would not introduce too great an error i n the 69 1 u I u n 1 CL) +1 Fig. 3 p K a of HfxOA r in H 2 S 0 4 / v 7 0 / / Slopes / H_o =1.38 -1 + / 0 A c i d i t y -) U) o F ig .4 p K a of h ^ C r O ^ in H C I 0 4 0 / ^ / S l o p e = 1.04 Ho# H_ o _-1 / / i -1 Ac id i t y - 2 70 'cf L_ O u I ICL) +.1 F ig 5 p K a Of / / H 2 C r 0 4 in HCI / / / / o/ f / / 0 f _0 7 / s l o p e = 1.68 / / Ho • 7 / / f / H_o /> / ./ / / /A torn/ | A c i d i t y 9 a) o £1 F i g . 6 p K a of H 2 C r 0 4 in H N Q 3 0 S lope = 0 .83 71 72 value obtained, since i n the cases where both the and H Q functions are known f o r a p a r t i c u l a r medium they do not deviate greatly i n reasonably d i l u t e solutions. This i s evident from a consideration of Figs. 3, 4 and 5 where both the H Q and H_ scales have been used. The H Q functions used were those reported by Paul and Long (91) while the H_ functions were obtained from a report by Boyd (92, 93). The pKa was also determined i n phosphoric acid containing a 1:1 molar r a t i o of sodium perchlorate (Fig„ 7 ) . Since no a c i d i t y function was a v a i l a b l e f o r t h i s medium i t was determined i n these laboratories. Its determination i s presented i n the next section. It should be noted that although theory predicts the plots i n F i g s . 3-7 should have a slope of unity, most of the plots d i f f e r considerably from t h i s . The reason f o r t h i s i s thought to be due to the same phenomenon that causes the values of the pKa to change from medium to medium and which w i l l be discussed l a t e r . 6. ACIDITY FUNCTION FOR THE PHOSPHORIC ACID-SODIUM PER-CHLORATE SYSTEM While glass electrode pH measurements may be 73 used to determine the a c i d i t y of d i l u t e acid solutions the a c i d i t y of more concentrated solutions i s generally defined i n terms of the H o function o r i g i n a l l y con-ceived by Hammett and Deyrup (90). The determination of t h i s function makes use of uncharged bases as i n -dicators and i s defined by the equation; Ho = P KBH + - l 0 S ^BHl where pKgfj + i s the equilibrium constant f o r the reaction B + H + F = ± BH + This concept, which has now become quite general (91), involves the use of weaker and weaker bases as i n -dicators as the a c i d i t y of the solution increases. In the present determination of the a c i d i t y of solutions containing a 1:1 molar r a t i o of phosphoric acid and sodium perchlorate the indicators used were p_-nitro a n i l i n e o = n i t r o a n i l i n e and 4-chloro=2-n i t r o a n i l i n e (Fig. 9). The r a t i o Cg^ -f- was determined spectroscopically using an absorption peak due to the unprotinated base. A stock solution containing 4.57 M phosphoric acid and 4.57 M sodium perchlorate was prepared by d i s s o l v i n g 280.0 gm. NaClO^ i n 3$1 ml. of 6 M H-jPO^ and d i l u t i n g the r e s u l t i n g mixture with water to exactly 500 ml. This solution was then d i l u t e d to give a series of solutions of varying 74 TABLE III ACIDITY OF SOLUTIONS CONTAINING A 1:1 MOLAR RATIO OF H3PO^ to NaCIO Molarity O.D. CBH +/ CB Ho (a) p_-Nitroaniline, pK = 0.99, Xmax = 380 0.0 .766 0.0457 .660 0.161 1.78 0.0914 .596 0.286 1.53 0.183 .503 0.525 1.27 0.457 .317 1.415 0.34 0.914 .175 3.39 0.46 1.33 .060 11.82 -0.03 (b) o-Nitroaniline, pK = -0.20, v m a v = 412 0.0 .660 0.457 .610 0.0823 0.34 0.730 .570 0.175 0.57 0.914 .542 0.216 0.46 1.33 .373 0.302 -0.08 2.74 .185 2.57 -0.60 3.66 .064 9.30 -1.17 (c) 4-Chloro-3-nitoaniline, pK = -0.95, X m Q V = 425 IucL JC 0.0 .648 1.83 .574 0.130 -0.08 2.74 .438 0.479 -0.60 3 . 6 6 4 .57 .246 .105 1.630 -1.17 5.17 =1.66 TABLE IV ACIDITY OF SOLUTIONS CONTAINING A 1:1 MOLAR RATIO OF HN03 to NaClO^ Molarity O.D. CBH +/ CB Ho (a) E - N i t r o a n i l i n e , pKa = 1.00, X m a x = 3^3 0.0 .741 0.0139 .652 0.136 1.89 0.0232 .618 0.198 1.68 0.0464 .529 0.401 1.38 0.0928 .396 0.870 1.06 0.186 .249 1.97 0.705 0.371 .124 5.05 Q.-297 0.464 .095 6.81 0.167 (b) o - N i t r o a n i l i n e ? pKa = -0.28, X m Q V = 383 0.0 .675 0.0928 .645 0.0460 1.060 0.1856 .609 0.109 0.705 0.464 .497 0.358 0.167 0.812 .373 0.809 -0.20 1.218 .246 1.74 -0.53 2.030 .109 5.17 -1.00 76 2.676 .044 14.40 -1.45 ) 4-Chloro-2. - n i t r o a n i l i n e , pKa - -0.98, X max ~ 0.0 .680 0.464 .641 0.0605 -1.-218 0.812 .591 0.150 -0.824 1.218 .511 0.330 -0.481 1.338 .481 0.415 -0.382 1.784 .381 0.786 -0.105 2.03 .321 1.12 -1.00 2,23 .283 1.40 -1.16 2.44 .208 2.27 -1.29 2.68 .192 2.55 -1.45 3.12 .125 4.44 -1.60 3.57 .072 8.42 -1.85 4.06 .030 21.7 -2.25 * For t h i s i n d i c a t o r the o p t i c a l density at t h i s wave length increased with time i n the more concentrated acid solutions. Hence a l l readings were extrapolated to zero time. F ig .8 Ionization of ind ica tors in H 3 P 0 4 " N a C I 0 4 s o l u t i o n s , o.p-ni troani l ine. • J o-n i t roan ihne. ©, 4-chloro-2-n i t roani l ine. Fig.9 Ionization of ind ica to rs in ' H N 0 3 - N a C I 0 4 ' so lu t ions , o ^ - n i t r o -anil ine. • .o-n i t roani l ine. o ^ - c h l o r o - ^ n i t r o a n i l i n e . 79 concentrations. F i f t y m i c r o l i t e r s of an a l c o h o l i c solution of indicator was then added to each of one dozen f i v e ml. volumetric f l a s k s . The f l a s k s were then f i l l e d to the mark with solutions of concentra-tions varying from pure water to 5.47 M a c i d . In addition one f l a s k was f i l l e d to the mark with 70$ s u l f u r i c a c i d . The v i s i b l e and u l t r a v i o l e t spectrum of the ind i c a t o r i n water was then determined to select the best wave length at which to work and the o p t i c a l density of each solution at t h i s p a r t i c u l a r wave length was measured on a Beckman D.U. Spectro-photometer. The data obtained hasebeen tabulated i n Table II I . In d i l u t e solutions the H determined from o p_-nitroaniline solutions was found to agree to with-i n 0.03 units with pH measurements made with a glass electrode on a Beckman Model G. pH meter. 7. ACIDITY FUNCTION FOR THE NITRIC ACID-SODIUM PERCHLORATE SYSTEM The H Q a c i d i t y function f o r a medium containing a 1:1 molar r a t i o of n i t r i c acid and sodium perchlorate was also determined by the same methods just described f o r the phosphoric acid-sodium perchlorate system. The indicators used f o r t h i s function were also £-nitro-a n i l i n e , o - n i t r o a n i l i n e and 4-chloro - 3-nitroaniline 80 (See F i g . 9). The data obtained i s presented i n Table IV. The H Q values f o r very d i l u t e solutions were found to agree to within 0.01 units with pH measure-ments made with a glass electrode on a Beckmann Model G pH meter. 8. DETERMINATION OF THE CHROMATE ESTER EQUILIBRIUM CONSTANT FOR 2-METHOXYETHANOL AND 2,2,2-TRIFLUOROETHANOL E q u i l i b r i a of the type, ROH + HCrO^" ^ ^ROCrO^" + H^ O have previously been studied spectroscopically by Klaning and Symons (43, 52). Their experimental procedure involved measuring the o p t i c a l density of chromium VI i n aqueous a c i d i c solutions containing varying amounts of an alc o h o l . I f a c t i v i t y c o e f f i c i e n t s can be neglected the following r e l a t i o n s h i p between the difference i n o p t i c a l density (Ad) and the equilibrium constant (Ke) can be derived: Ad [H20] / [ROH] = 4eKel [ c r V I ] - Ke Ad (I) where A€ i s the difference i n extinction c o e f f i c i e n t between HCrO^" and ROCrO^" at 375 m u. and 1 i s the length of the c e l l used. Accordingly the slope of a plot of A d p 20] / [ROH] against Ad should equal -Ke. While Klaning and Symons have determined the equilibrium constant f o r a large number of alcohols i t was found necessary to determine t h i s constant f o r two a d d i t i o n a l alcohols, 2-methoxyethanol and 2,2,2-81 trifluoroethanolo For comparison purposes the equilibrium constant f o r isopropyl alcohol was also determined,, The r e s u l t s obtained are presented i n graphical form i n Fig„ 10. The value of 8.2 f o r isopropyl alcohol obtained i n t h i s investigation agrees very well with the value of 8.2 reported by Klaning and Symons. By use of i d e n t i c a l methods the equilibrium constant f o r 2=methoxyethanol was found to be 8.3. The determination of Re f o r 2,2,2-trifluoro-ethanol was complicated to some extent by two factors; ( i ) the equilibrium constant i s only about one-sixth as large as that f o r the two other alcohols and ( i i ) the difference i n extinction c o e f f i c i e n t s f o r the ester of t h i s alcohol and HCrO^" i s apparently some-what smaller than i n the other cases. As can be seen from F i g . 10, the changes i n o p t i c a l density were not large enough to permit an accurate slope to be drawn. Hence, a d i f f e r e n t approach involving the use of the intercept of the l i n e was employed i n t h i s case. From equation (I) i t can be seen that the value of the intercept should be equal to AeKel [ c r ^ ^ ]. The only unknown i n t h i s term i s A e . Thus an attempt to measure the extinction c o e f f i c i e n t of the ester was made following the procedure of Westheimer et a l $2 (42). An aqueous a c i d i c solution of 2,2,2=>trifluoro-ethanol was extracted with benzene. The organic layer containing extracted ester and alcohol was dried b r i e f l y over CaCl2 and the o p t i c a l density was determined at 375 m jx» The amount of ester present was determined from the decrease of chromium VI concentration i n the aqueous la y e r . This decrease i n concentration was calculated from measurements of the o p t i c a l density of the aqueous solution before and a f t e r extraction with benzene and a consideration of the volume change occurring during the extraction. By t h i s method, the molar extinc t i o n c o e f f i c i e n t of the ester was deter-mined to be approximately 1150 at 375 m ju. From t h i s Ae was calculated to be 310 and the Ke value was determined to be 1.29. While i t i s apparent that t h i s number may contain considerable error i t i s fortunate that the conclusions derived from i t are not very s e n s i t i v e to the accuracy of t h i s p a r t i c u l a r r e s u l t . Thus i t can be seen from a consideration of the calculations to be presented i n Sec. IV. that the same general conclusions would be obtained even i f t h i s r e s u l t contained an error of several hundred percent. Indeed the same conclusions would be derived i f t h i s r e s u l t were disregarded altogether. 8 Fig.10 C h r o m a t e e s t e r equ i l ib r ium C o n s t a n t s 84 TABLE V DATA FOR FIGURE 10 O r V I - 1.00 x 10" 4 M HC10, 4 = 1.00 x 10"3 M [Alcohol] [H20] O.D. Ad Ad x [H (a) Isopopyl Alcohol 0.0 .820 0.785 52.0 .920 .100 6.64 1.31 50.0 .980 .160 6.10 2.62 44.9 1.116 .296 5.09 3.92 39.9 1.200 .380 3.77 5.24 34.7 1.320 .500 3.31 (b) 2-Methoxyethanol 0.0 .820 1.27 50.2 .930 .110 4.35 2.54 44.6 1.026 .206 3.62 3.81 39.3 1.119 .299 3.08 5.08 34.2 1.182 .362 2.44 (c) 2,2,2 -Trifluoroethanol 0.0 .820 1.37 50.2 .830 0.010 .36 4.12 40.1 .843 0.023 .22 6.87 29.0 .855 0.035 .15 85 IV RESULTS AND DISCUSSION 1. OXIDATION OF FLUORO ALCOHOLS IN 77.2$ ACETIC ACID The rates of oxidation of eight a r y l t r i f l u o r o -methylcarbinols were determined at 25°C i n a 77.2$ acetic acid solution containing 3.20 M perchloric ac i d . The deuterium isotope e f f e c t s f o r f i v e of these alcohols were also determined. In addition the rates of oxidation of f i v e substituted 2-propanols along with the deuterium isotope effects f o r three of these were determined f o r comparison purposes. The re s u l t s are presented i n Table VI. A Hammett plot using modified 6^  values f o r both the protio and deuterio compounds i s shown i n F i g . 11, and the same plot using the o r i g i n a l Hammett^ values i s shown i n F i g . 12. Although the co r r e l a t i o n c o e f f i c i e n t s are approximately the same f o r both plots (0.985 and 0.990 respectively) an examination of Figures 11 and 12 reveals that there i s a d e f i n i t e break i n the slope of the plot obtained when the o r i g i n a l Hammett values are used. Such a break i s not apparent f o r the plot i n F i g . 11 where the modified cr*+ values are used. As a further a i d i n deciding whether the rates of oxidation were better correlated with ff^or ^ values attempts were made to determine the rates of oxidation of p_°methoxyphenyltrifluoromethylcarbinol 86 TABLE VI OXIDATION OF FLUORO ALCOHOLS IN 77.2$ ACETIC ACID, 3.20 M HCIO^ (H Q = -2.82), T - 25 C Alcohol No . [ c r V I ] [ a i c ] (MxlO 3) (Mxl03) (1 moles" 1 sec'l) £-t-Butylphenyl- 56 1.79 « 2.68 trif l u o r o m e t h y l -carbinol 57 " .574 .544 k H / k D £-Methylphenyl- 44 tri f l u o r o m e t h y l -carbinol 55 .105 111 534 537 564 540 £-Methylphenyl- 107 trifluoromethyl-carbinol-oC-d 109 110 115 it n w n .0727 .0740 .0747 .0706 7.40-.20 m°Methylphenyl- 64 trif l u o r o m e t h y l -carbinol 68 .350 .340 Phe n y l t r i f l u o r o - 82 ..methylcarbinol 89 41 15.3 .259 .250 .258 Phe n y l t r i f l u o r o - 39 methylcarbinol-oC-d 90 92 16.1 2.68 .0310 .0294 .0297 8.531.15 87 Pj=Chlorophenyl- 80 trif l u o r o m e t h y l -carbinol 81 m-Bromophenyl- 42 trifluoromethyl-carbinol 60 61 91 m-Bromophenyl- 66 trifluoromethyl-carbinol'-oc^d 72 75 m-Nitrophenyl- 47 trif l u o r o m e t h y l -carbinol 62 101 78 119 113 m-Nitrophenyl- 117 trifluoromethyl-carbinol-oOd 118 3,5-Dinitro- 138 phe n y l t r i f l u o r o -methylcarbinol 143 3,5-Dinitro- 132 phe n y l t r i f l u o r o -methylcarbinol- 139 oc-d 1»79 2.68 .174 n n .174 n n .119 n n .106 n n .110 ft n .109 » n .0117 9.80*. 20 n .0111 t» 20.8 .0113 t» 2.68 .0494 n tt .0514 n n .0484 n 14.06 .0453 17»22 .0444 » 15.96 .0474 17.28 .00360 12.20±.2( n 18.69 .00367 n 17.65 .0116 n 12.81 .0118 n 55.10 .00100 12.93±.7* n 10.92 .000834 Isopropyl alcohol l-Fluoro-2-propanol 1 , 3-Difluoro-2-propanol 1 , 3-Difluoro-2-propanol-2-d 1 , 1 , 1 - T r i f l u o r o -2-propanol 1 , 1 , 1 - T r i f l u o r o -2-propanol-2-d l , l , l - , 3 , 3 , 3 -Hexafluoro - 2 -propanol 1 ,1 ,1 ,3,3 ' , 3 -Hexafluoro - 2 -propanol-2-d 175 1.29 1.93 176 » " 187 1.37 2.06 188 '» » 178 " 30.2 179 1 1 30.2 180 " 3 0 . 0 181 *» 2 6 . 0 182 » 37.2 183 "' 25 .8 189 0.685 31 .6 190 0.685 25.3 194 0'.633 63.3 193 0.818 158 576 1.03 190 577 1.08 238 578 1.47 79.5 579 1.14 249 88 4730* _> * 5380 15.4 15.2 .429 .425 .0407 1 0 . 5 0 ! . 1 0 .0292 .0302 .00263 11 .29* .20 -. 0000245 .0000220 .0000285 .0000169 .0000160 .0000158 .00000322 7.8+.5 '.00000210 * Approximate rate only 89 Fig.11 H a m m e t t plot f o r t h e c h r o m i c ac id ox idat ion of s u b s t i t u t e d p h e n y l t r i -f l uo rome thy l ca rb ino l s in 7 7 % a c e t i c a c i d . O p r o t i o c o m p o u n d s 0 0.5 1.0 90 Fig.12 H a m m e t t plot f o r t h e c h r o m i c ac id 2Q o x i d a t i o n of s u b s t i t u t e d p h e n y l t r i -— ' f l u o r o m e t h y i c a r b i n o l s in 7 7 % a c e t i c ac id . p - M e 1.0 1.5 2 0 91 and 0-naphthyltrifluoromethylcarbinol, Since the difference between cr* and values f o r the £-methoxy and the <#-naphthyl group are greater than f o r any of the other groups they are most useful i n d i f f e r e n t i a t i n g between reactions which follow o> values and those which follow ^ values. Unfortunately both of these alcohols reduced chromium VI at a very f a s t rate i n the medium used. This can be attributed to the oxidation of these compounds at s i t e s other than the hydroxylated carbon since i t was l a t e r found that both anisole and naphthalene were very r a p i d l y oxidized under these conditions. A modified w r h o w value of -1.01 was calculated from the slope of the plot i n F i g . 11. While other interpretations (22,27) have been suggested for the negative rtrhow values observed i n the oxidation of alcohols by chromic acid the most d i r e c t explanation f o r the f a c t that any reaction exhibits a negative c value i s that considerable pos i t i v e charge i s located on the oc-carbon atom i n the t r a n s i t i o n state. Further, the observation o f very pronounced isotope e f f e c t s demands that any proposed mechanism include as a r a t e -determining step, f i s s i o n of the oC-carbon=hydrogen bond. Taken together these two observations suggest that hydride abstraction i s taking place. However, 92 the r e s u l t s reported by Westheimer et a l (22) on the oxidation of isopropyl alcohol i n indicate that the o r i g i n a l hydride ion mechanism proposed by Rocek and Krupi£ka (31) cannot be correct. Also the evidence avail a b l e i n the l i t e r a t u r e , p a r t i c u l a r l y the evidence obtained from studies of the oxidation of d i o l s (24, 25, 26) strongly indicates that the chromate esters play the r o l e of intermediates i n t h i s reaction. A l l of these considerations are s a t i s f a c t o r i l y met by a mechanism which employs as a rate-determining step decomposition of the chromate ester by i n t e r n a l hydride ion transfer. \ C ACT— OH R H* 0 R .0, C 0 . 1 1 rCr—OH OH R - ;0 • \ // -V C=0 + Cr / / \ R H0 OH (I) It can be argued that since mechanisms such as the one above involve e s s e n t i a l l y the movement of electrons i n a cycle the eff e c t of substituents should be the same regardless of the d i r e c t i o n of electron flow. However, the very pronounced isotope e f f e c t observed d e f i n i t e l y indicates that f i s s i o n of the oC-carbon-hydrogen bond i s rate-determining. Hence stretching of t h i s bond probably i n i t i a t e s the reaction 93 and the t r a n s i t i o n state must be s i m i l a r to (I)„ I f the mechanism involves simultaneous f i s s i o n of the oxygen-chromium and the carbon-hydrogen bonds the ef f e c t of substituents would s t i l l be more pronounced on the carbon-hydrogen bond since i t i s p h y s i c a l l y closer to the substituents. Such an e f f e c t has been elegantly i l l u s t r a t e d by Buncel and Bourns (95) i n t h e i r work on the de-composition of benzyl n i t r a t e by ethoxide ions. This reaction shows a deuterium isotope e f f e c t of 5.04 at 60°C and an N 1 5 isotope e f f e c t of 1.0196 at 30°C in d i c a t i n g that the carbon-hydrogen and the nitrogen-oxygen bonds are broken simultaneously i n the course of the reaction. In t h i s case i t i s evident that the hydrogen i s being abstracted as a proton and the large p o s i t i v e w r h o " value of +3«40 indicates that the e f f e c t of substituents i s more intimately concerned with cleavage of the oC=carbon-hydrogen bond than the more remote oxygen-nitrogen bond. H C6H5-CH2-0-N02+ EtO" ^ C^H.CHO * NO," + EtOH 94 In the case of a c y c l i c chromate ester de° composition movement of electrons i n the reverse d i -r e c t i o n to that shown i n the reaction sequence involving (I) would r e s u l t i n a proton transfer mechanism iden-t i c a l with the one proposed by Kwart and Francis (50). However, the negative n r h o w values observed i n t h i s and other investigations (27, 32) f o r the chromic acid oxidation of alcohols are not, i n l i g h t of the benzyl n i t r a t e r e s u l t s , compatible with a rate-determining proton transfer. Moreover, as Roc'ek (32) has indicated, one would not expect that protonation of the chromium part of the ester would lead to rate enhancement i f the rate-determining step also involved i n t e r n a l proton transfer to t h i s part of the molecule. The hydride transfer mechanism, on the other hand, r e a d i l y accommo-dates t h i s observed feature of the reaction. In addition i t w i l l be shown that the other r e s u l t s obtained i n t h i s i n v e s t i g a t i o n can also be e a s i l y ex-plained on the basis of such a mechanism. With the exception of 1 ,1 ,1 ,3 ,3 ,3-hexafluoro-2-propanol an approximately l i n e a r r e l a t i o n s h i p exists 95 between the rates of oxidation of trifluoromethyl-carbinols and the observed isotope e f f e c t s . This r e l a t i o n s h i p , which i s i l l u s t r a t e d i n F i g . 13 f i t s i n well with the suggestion f i r s t made by Wiberg (59) that there should be a general c o r r e l a t i o n between rates and isotope e f f e c t s within a group of analogous reactions, slow rates being accompanied by large isotope e f f e c t s . Wiberg based h i s predictions on the experimental observation that the isotope e f f e c t s i n the side chain halogenation of para substituted toluenes varied from 3.22 to 5.11 at 77°C. Since the larges t of these ..isotope e f f e c t s i s somewhat smaller than the maximum e f f e c t predicted from a consideration of zero-point energies of the carbon1-hydro gen and carbon-deuterium bonds, the v a r i a t i o n with rate was explained i n terms of factors which could be related to C-H stretching frequencies (59, 67). In the present case, however, even the smallest isotope e f f e c t observed i s larger than the maximal ef f e c t predicted from a consideration of carbon-hydrogen stretching frequencies. While i t i s not clear from the present r e s u l t s whether the large isotope e f f e c t s are due to loss of bending frequencies i n the t r a n s i t i o n state or to a tunnel e f f e c t i t i s in t e r e s t i n g to note that a co r r e l a t i o n between rate ^ p - C H 3 C 6 H 4 -m - B r C 6 H 4 -\ v o m - N 0 2 C 6 H 4 -C H 3 9 \ . _z2 3 J 5 - d i N 0 2 < ^ H 4 -•p l o g ( K H / K D ) 1,1 1.2 I I Fig.13 Relat ionship b e t w e e n t he r a t e of o x i d a t i o n of t r i f l o ro -methy lcarb ino ls in 7 7 2 % acet ic acid and t h e p r i m a r y d e u t e r i u m iso tope e f fec ts . 97 and isotope e f f e c t s t i l l e x i s t s . Such a c o r r e l a t i o n was not observed by Stewart and Van der Linden (71) i n the permanganate oxidation of these alcohols since the "rho" value for that reaction was almost zero. It does appear however that large isotope e f f e c t s are i n t i -mately connected with the trifluoromethyl group or at l e a s t with the presence of strongly electron with-drawing groups adjacent to the s i t e of oxidation. The isotope e f f e c t observed f o r the chromic acid oxidation of 1,,1,1,3 ?3 ,3-hexaf luoro-2-propanol i s anomalous i n that i t i s only 7.8 while the rate of oxidation of t h i s compound i s the slowest i n the s e r i e s . This low isotope e f f e c t i s not l i k e l y to be due to impurities i n the deuterio compound since i t was c a r e f u l l y p u r i f i e d by vapor phase chromatography before use. Further i t was found that the properties of t h i s compound prepared independently by two d i f f e r e n t workers i n t h i s laboratory agreed well and that the same isotope e f f e c t was observed for each sample. The deuterium content of these samples was e a s i l y checked by nuclear magnetic resonance spectroscopy and found to be above 97$. The small amount of protio compound present gave the rate plot a steep i n i t i a l slope. However, over an extended period of time good straight l i n e k i n e t i c s were observed. Also i t was noted that 98 the rate of reaction of t h i s compound depended to some extent on the concentration of the alcohol present, being 0.025 L moles" 1 sec." 1 at an alcohol concentra-t i o n «6M. and only 0.016 1. moles" 1 sec." 1 at an alcohol concentration of 1.0M. This may be due to an e f f e c t of the solvent that i s observed f o r alcohol concentrations as large as t h i s . The lower than expected isotope e f f e c t may be a manifestation of experimental conditions since i t i s possible i n t h i s case that the intermediate chromium V and chromium IV species are reacting to a considerable extent with the solvent i . e . the reducing p o t e n t i a l of 1,1,1,3 ,3 ,3'=hexafluoro-2-propanol may not be much greater than that of a c e t i c acid, and since the con-centration of the solvent i s so much greater than that of the alcohol a considerable amount of side reaction may be occurring. Kaplan's work has shown that an isotope e f f e c t i s also observed f o r these intermediate species i n the oxidation of 2=propanol= 2-t „ Hence the o v e r a l l r e s u l t of the reduction of these species by solvent would be a decrease i n the magnitude of the observed isotope e f f e c t . The rate of oxidation of phenyltrifluoromethyl-carbinol was found to be l i n e a r with h as indicated by o ' the plo t i n F i g . 14. The a c i d i t y data, reported i n 99 Table VII was obtained by a comparative method using p_=nitrodiphenylamine and 4~chloro»2-nitroaniline as in d i c a t o r s . (The pKa's of these indicators were assumed to be =2.48 and -1.03 respectively (91).) This observed l i n e a r i t y with h i s i n accord with the rate law observed f o r the oxidation of isopropyl alcohol (34) and benzaldehyde (12) i n acetic a c i d . The "rho n value observed f o r the oxidation of the ar y l t r i f l u o r o m e t h y l c a r b i n o l s i s also i d e n t i c a l with the value observed f o r the corresponding oxidation of arylmethylcarbinols i n acetic acid (27). Hence i t i s probable that the trifluoromethyl alcohols are oxidized v i a a reaction path s i m i l a r to the one followed by t h e i r hydrocarbon analogs. 100 J .2 _A / _ 0 V S I o p e = 1.12 zA / - 8 / i i i Ho | | 1.5 2.0 2.5 3.0 F i g . 14 L i n e a r re la t i onsh ip b e t w e e n r a t e and ac id i t y f o r t h e o x i d a t i o n of p h e n y l -t r i f l u o r o m e t h y l c a r b i n o l in 7 7 . 2 % a c e t i c a c i d . 101 TABLE VII VARIATION IN RATE OF OXIDATION OF PHENYLTRIFLUOROMETHYLCARBINOL IN 77.2$ ACETIC ACID WITH H^ o No H C 1 0 L ! H 0 k2 ( M f 82 3.20 -2.82 .259 89 " 9 .250 41 * " .258 155 2.47 -2.12 .0433 159 " " .0440 156 1.83 -1.40 .0059 157 w " .0066 2. OXIDATION OF SOME HALOGENATED 2-PROPANOLS IN 50.1$ SULFURIC ACID The oxidation of isopropyl alcohol by chromic ac i d i n s u l f u r i c acid medium bears a resemblance to the corresponding reaction i n acetic acid solutions i n so f a r as the same rate law i s obtained and an isotope e f f e c t of approximately the same magnitude i s observed (34, 37). (One notable difference i s the slower rate of oxidation i n s u l f u r i c a c i d media of the same acidity») 102 It was of i n t e r e s t , therefore, to determine the rates of oxidation and the isotope e f f e c t s f o r some of the fluori n a t e d alcohols i n s u l f u r i c acid media. A solution of 50.1$ ^SO^ by weight was chosen because t h i s was the highest a c i d i t y at which the isotope e f f e c t f o r the oxidation of isopropyl alcohol could be accurately determined. A preliminary investigation indicated, however, that a study of the oxidation of a r y l t r i f l u o r o m e t h y l c a r b i n o l s i n s u l f u r i c acid solutions was not f e a s i b l e because of the low s o l u b i l i t y of these alcohols i n aqueous a c i d i c media. The halogenated=2-propanols, on the other hand possessed suitable s o l u b i l i t y c h a r a c t e r i s t i c s and the rates of oxidation of seven such alcohols were determined. In addition 1 the deuterium isotope e f f e c t s were determined f o r f i v e of.these compounds. It was expected, i n analogy with the r e s u l t s obtained i n 77.2$ ac e t i c acid, that the isotope e f f e c t s i n s u l f u r i c acid would increase as the rate of oxidation decreased. However, as the r e s u l t s i n Table VIII indicate, the isotope e f f e c t f o r 1,3-dichloro-2-propanol and l,3-difluoro-2-propanol were even lower than the e f f e c t observed for 2-propanol. The reason f o r the lower isotope e f f e c t s observed f o r these two compounds may, however, be due to the 103 occurrence of a small amount of halogen hydrolysis i n s u l f u r i c a c i d media. The hydrolysis product, a primary alcohol, would be oxidized immediately under these conditions with no isotope e f f e c t . This would lead to the observation of an isotope e f f e c t somewhat diminished i n magnitude. Hydrolysis was probably not observed i n 77.2$ a c e t i c acid solutions since the rate of oxidation was much f a s t e r i n that medium and a slow rate of hydrolysis would not have an observable e f f e c t on the r e s u l t s . i In any event i t i s apparent that carbon-hydrogen bond f i s s i o n constitutes the rate-determining step f o r the chromic acid oxidation of a l l alcohols studied i n both 77.2$ a c e t i c a c i d and i n 50.1$ s u l f u r i c acid solutions. A Taft <r - £ plot shown i n F i g . 15 gives a 6* value of =0.94 f o r the reactions i n 50.1$ s u l f u r i c a c i d . This i s comparable i n magnitude with the value of -1.01 obtained f o r i n 77.2$ acetic a c i d and i n -dicates that the reaction paths are probably s i m i l a r , at l e a s t i n some respects, i n both media. Results which w i l l be presented i n part IV of t h i s section indicate that the heats and entropies of a c t i v a t i o n f o r the oxidation of l , l , l - t r i f l u o r o - 2 = p r o p a n o l are almost i d e n t i c a l i n both s u l f u r i c and acetic a c i d mediae This also lends weight to the idea that the 104 TABLE VIII RATES OF OXIDATION OF SUBSTITUTED 2-PROPANOLS IN 50.1$ SULFURIC ACID, 1.03 x 1 0 ° 3 M, T -Alcohol No. [ A l e ] k l k 2 (MxlO 2) ( s e c ^ x (1 molei 10 3 ) sec" 1) Isopropyl 393 1.31 55.2 4.21 alcohol 394 t» 55.8 4.26 395 ro 53.1 4.06 Isopropyl 486 3.43 23.4 0.683 alcohol - 2-d 487 2.14 14.4 0.673 488 1.29 8.09 0.626 l~Chloro~ 2 - 482 1.19 7.59 0.636 propanol 1.87 0.620 483 11.6 485 3.30 20.4 0.619 l=Fluoro-2= 542 2.60 9.65 0.371 propanol 0.388 543 1.30 5.05 544 2.17 8.46 0.390 l~Fluoro - 2 = 545 2.14 1.05 0.0491 propanol-2-d 546 0.676 1.29 0.0524 1 , 3-Dichloro-2-propanol 491 1.76 0.925 0.0525 489 1.41 0.719 0.0510 490 1.06 0.580 0.0545 k H / k D 6.44 7.56 105 lB3°Dichloro=2~ 496 3.43 0.306 0.00893 propanol-2-d 0.161 497 1.75 0.00920 l~Chloro=3~ 504 1.93 0.800 0.0414 Fluoro=2-propanol 505 1.15 0.439 0.0382 l j 3=Difluoro-2-propanol 478 1.29 0.256 0.0198 479 2.13 0.411 0.0193 4$0 4.30 0.839 0.0195 l 9 3=Difluoro-2~propanol-2-d 510 2.14 0.0599 0.00280 4.28 511 0.122 0.00285 1 ,1 ,1-Tr.lfluoro- 465 3.74 0.0436 0.00116 2=propanol 494 3.72 0.0421 0.00113 495 7.44 0.0907 0.00122 1 , 1 9 1-Trifluoro- 469 7.74 0.00865 0.000112 2-=propanol-2-d 485 3.69 0.00370 9.060110 5.85 6.92 106 Fig.15 Taft a>*-Q*"plOt f o r t he ox ida t ion of subs t i t u ted 2-propano ls in 50 .1% H SO . 107 reaction mechanisms are s i m i l a r but not i d e n t i c a l i n both media. A complete explanation f o r the marked i n -crease of rate i n a c e t i c a c i d solutions remains un-obtainable at the present time. 3 0 VARIATION IN THE MAGNITUDE OF ISOTOPE EFFECTS WITH CHANGES IN ACIDITY The magnitude of the isotope e f f e c t was found to correlate with changes i n rates of oxidation caused by introduction of substituents into the structure of the alcohol. It therefore seemed p o s s i b l e that changes i n the magnitude of the isotope e f f e c t might also be observed when the rate of reaction was varied by changing the a c i d i t y of the medium. Hence the re-l a t i v e rates 6°f oxidation of 2-propanol and 2-propanol-2-d were determined i n perchloric acid solutions of varying concentrations. The r e s u l t s , which are pre-sented i n Table IX indicate that within experimental error no change i n isotope e f f e c t i s observed when the rate of the reaction i s changed by a l t e r a t i o n s i n the a c i d i t y of the medium.(in d i l u t e HCIO^). From these r e s u l t s i t would seem that the introduction of substituents into the substrate can cause a change i n the shape of the p o t e n t i a l energy b a r r i e r while a change i n a c i d i t y of the medium does not change the shape of the b a r r i e r ; i t merely i n -108 TABLE IX DEUTERIUM ISOTOPE EFFECTS FOR THE OXIDATION OF ISOPROPYL ALCOHOL IN PERCHLORIC ACID SOLUTION OF VARYING ACIDITIES (T = 25°C, [ c r V 1 ] - 3.4 x 10~ 3 M) [ H C I O 4 ] H O k H k D k H / k D ( M ) (1 moles" 1sec° 1xl0 2) 0.463 +0.20 .159 0.0259 6.2 0.93 -0.18 .713 0.1037 6.9 2.08 -0.82 5.76 0.943 6.1 2.63 -1.05 14.1 2.25 6.3 3.82 ^1.62 65.6 10.4 5.9 4.52 -1.99 171 26.2 6.5 109 creases the equilibrium concentration of the reacting species. 4. THERMODYNAMIC PARAMETERS The v a r i a t i o n of the reaction rate with tern-perature can be u t i l i z e d to obtain the heat and entropy of a c t i v a t i o n , respectively AH^ and &S^. The equation -AH# . AS#  k = kT RT + R *2 h x e as derived from t r a n s i t i o n state theory would predict that i f log k 2/T°K was plotted against 1000/T°K (over a l i m i t e d temperature range), the slope of the str a i g h t m l i n e and the intercept would give respectively A H" and AS# (65b) . The rate constants f o r the oxidation of 1,1,1-trifluoro - 2-propanol and isopropyl alcohol under various conditions were determined at several tem-peratures. As indicated i n F i g . 16 plots of log k 2/T°K against 1000/T°K gave good straight l i n e s from which the heat and entropy of a c t i v a t i o n f o r each of these reactions was calculated. The values determined f o r these two thermodynamic parameters are recorded i n Table X, while the rates f o r the reactions at various 110 TABLE X HEATS AND ENTROPIES OF ACTIVATION FOR THE OXIDATION OF SUBSTITUTED 2-PR0PAN0LS UNDER VARIOUS CONDITIONS Compound Conditions AH AS  (kcal/mole) (cal/mole deg.) 1,1,1-Trlfluoro- 50.1$ HgSO, 8.93 =42.0 2-propanol 1,1,1-Trifluoro- 77.2$ Acetic 8.15 -38.1 2-propanol Acid 3.20 M HCIO^ Isopropyl alcohol 5.29$ H 2S0 10.38 -35.7 I l l TABLE XI RATES USED IN CALCULATION OF THERMODYNAMIC PARAMETERS • No. T(°C) [cr V I J [alc7|' k x k 2 (MxlO 3) (MxlO 2) (sec~^x (1 mole I IO 3) s e c - 1 ) Isopropyl 344 50.0 3.40 i 42.1 3.94 0.00935 alcohol i n 5.29$ H 2S0 4 349 35.8 3.90 42.1 1.84 0.00437 334 25.0 4.10 12.9 0.296 0.00230 l, l , l - T r i f l u o r o - 5 3 7 2-propanol i n 50.1$ H 2S0^ 465 16 .4 25.0 1.03 tt 3.73 3.741 0.0248 0.0436 0.000667 0.00116 494 it 3.72 0.0421 0.00113 495 it 7.44 0.0907 0.00122 536 39.5 ft 3.72 | i 0.0874 0.00232 l, l , l - T r i f l u o r o - 1 8 6 2-propanol i n 77.2$ Acetic 181 Acid, 3.20 M HCIO^ 182 16.4 25.0 « 1.37 tt 3.72 ] 2.60 I 3.72 .677 .760 1.12 0.0182 0.0292 0.0302 184 39.5 » 3.72 2.12 0.0570 112 Fig.16 The rmodynam ic pa rame te rs . 113 temperatures are given i n Table XI. While i t i s possible to explain the large negative entropy of a c t i v a t i o n observed f o r these reactions i n a number of ways i t i s c e r t a i n l y con-s i s t e n t with the c y c l i c t r a n s i t i o n state that has been postulated f o r these reactions. I t i s also of in t e r e s t to note that the thermodynamic parameters f o r chromic acid oxidations i n ac i d medium are very s i m i l a r to those reported f o r acid permanganate oxidations i n s u l f u r i c acid media (72). 5. OXIDATIONS IN CONCENTRATED SULFURIC ACID SOLUTIONS It i s known that the rate of oxidation of many organic compounds by chromic acid i n s u l f u r i c acid media i s maximal i n solutions containing 60-75$ H 2 ^ 4 by weight (21, 30, 31). Pungor and Trompler (30) claim that an explanation for t h i s maximal rate of oxidation of methanol, ethanol and formic acid i n 73$ H 2 S 0 ^ may l i e i n a consideration of the s t a b i l i t y of chromium IV i n various solutions of s u l f u r i c acid. Rocek and Krupicka on the other hand have suggested that the rate versus a c i d i t y plot for the oxidation of isopropyl alcohol may change from a slope of about unity to one of zero at about 55$ ^SO^ due to a protonation of the alcohol at t h i s point. The former explanation requires that rates of oxidation of a l l 114 compounds be maximal at the same point, while the l a t t e r demands that the point of slope change vary with the pKg^+ of the compounds. Rocek has also reported that the rates of oxidation of several carboxylic acids i n acetic a c i d - s u l f u r i c acid mixtures show a break i n the rate versus a c i d i t y plot at an H Q of about -3.6 (20). This i s only 0.3 units l e s s than the H Q of 55% s u l f u r i c a c i d . He has also presented r e s u l t s which he claims indicates that the hydrocarbon n-heptane does not show a rate maximal at t h i s point. However the ex-perimental evidence presented f o r t h i s compound i s very weak and could be interpreted either way. While i t i s not probable that compounds of as widely d i f f e r e n t structures as 2-propanol and p_-toluic acid would have s i m i l a r pKg^+'s, i t i s possible since they are both oxygenated hydrocarbons. Another possible explanation, suggested by Rocek for the occurrence of t h i s break i n the rate versus a c i d i t y p l o t , involves the protonation of chromic acid at t h i s point. While t h i s p o s s i b i l i t y .was re-jected on the basis of the aforementioned r e s u l t s obtained f o r the oxidation of n-heptane, a spectro-scopic study conducted i n the present i n v e s t i g a t i o n showed that there was indeed a change i n the nature 115 of the chromium species at an H Q of about -fc.4. F i g . 17 shows a p l o t of the extinction coefficients of chromium VI at various wave lengths against the a c i d i t y of a s e r i e s of s u l f u r i c acid solutions. The plot shows a well defined sigmoidal curve with a mid-point at H Q = +0,3. This has previously been ascribed to an equilibrium of the type HCr0j^"+ H + ===== H 2 C r 0 4 • Another change i n the spectra i s also observed at about H Q = -ft.4. This change i s not as pronounced since s t i l l another change appears at only a s l i g h t l y higher a c i d i t y . One possible explanation may be that protonation of I^CrO^ i s followed at only s l i g h t l y higher a c i d i t i e s by a sulphonation step. The v i s i b l e and u l t r a v i o l e t spectra of chromium VI i n s u l f u r i c acid solutions of various strengths i s shown i n F i g . 18. Further, Roc'ek's suggestion that the observed crest i n the r a t e - a c i d i t y p l o t was due to protonation •4-of the organic substrate required that the pKg^ of p_-toluic acid be s h i f t e d from -6.92 i n s u l f u r i c a c i d -water media to about -3.8 i n s u l f u r i c a cid-acetic acid media. Such a s h i f t seems e n t i r e l y u n l i k e l y and an attempt was made to determine the pKgjj+ of p_-toluic acid i n a c e t i c a c i d - s u l f u r i c acid media using the a c i d i t y function developed by Rocek. Such an i n v e s t i -_3000 Fig 17 Ext inct ion coef f ic ients of C r ^ in sulfuric acid solut ions € I /o o -.6 H . -7 117 TABLE XII EXTINCTION COEFFICIENTS OF G r V I IN SULFURIC ACID SOLUTIONS $H 2S0 4 H- 1445 |408 $313 J347 J273 2.95 220 260 709 1540 1750 2 , 66 214, 257 704 1550 1750 1.81 213 253 702 1550 1760 1 . 0 1.00 209 236 671 1460 1810 1,75 0.82 205 226 650 1410 1830 2«40 0.74 199 208 627 1370 1870 4.80 0.44 193 173 581 1250 1920 5.29 0.40 193 173 573 1210 1930 10,6 -0,14 194 155 538 1090 1990 20«2 -1.02 168 129 496 1020 2030 25<;8 -1.56 192 140 510 1000 2030 36^4 -2.58 174 118 481 961 1970 41.4 -3.08 178 121 477 938 I960 46,2 -3.54 173 122 491 922 1980 50^4 -3.96 168 118 477 879 1900 54.4 -4.39 173. 130 507 890 1900 59.4 -4.84 218 140 521 867 1900 64.6 -5.52 183 229 858 998 2180 67.9 -5.96 180 292 1160 1010 2310 74<iO -6.68 254 613 2580 1790 4380 78,8 -7.16 241 662 3160 2110 5170 84.0 (-8,00) 192 482 2790 1820 5550 88.5 (-8.60) 216 391 2260 1380 4830 95.0 (-9.40) 206 320 2000 1190 4670 \ Fig.18 U l t r a v i o l e t s p e c t r a of C r 2 1 in H^SO^ s o l u t i o n s . 1000 3 0 0 v , , 3 5 0 J X(mjj) 119 gation was a rather d i f f i c u l t undertaking since the u l t r a v i o l e t peaks usually used f o r such determinations were obscured because of absorption by the a c e t i c acid medium. However, the spectral s h i f t s were followed at two wavelengths and as the p l o t i n F i g , 19 indicates i t i s quite probable that the pKgH+ does not s h i f t appreciably i n acetic a c i d - s u l f u r i c acid media. In a f i n a l attempt to resolve t h i s problem i t was decided to study the rates of oxidation of three alcohols 2-propanol, 1,1,1-trifluoro-2-propanol and.1,1,1,3,3,3-hexafluoro-2-propanol i n solutions of varying s u l f u r i c a c i d concentrations. While the pKg^+ values f o r these alcohols are not known i t i s to be expected that they w i l l vary greatly since i t has been found by Haszeldine (97) that the pKa values for these alcohols vary by fa c t o r s of greater than four u n i t s . The f a c t that i t has recently been found that a good c o r r e l a t i o n exists between the pKa and the pKgg+ values of many organic compounds lends weight to t h i s suggestion (9#)« A r a t e - a c i d i t y p l o t f o r these alcohols i s shown i n F i g . 20. (The r e s u l t s obtained by RoKek (31) from a polarographic i n v e s t i g a t i o n of the rates of oxidation of 2-propanol i n s u l f u r i c acid solutions of varying strengths are included as closed c i r c l e s f o r comparison purposes.) It can be seen that the plots for the 120 Fig. 19 A t t e m p t e d de te rm ina t i on of t h e pKa f o r _p-to1uic acid in acet ic acid. 1.8 1.4 Opt ica l D e n s i t y / / 3 4 5 6 e x p e r i m e n t a l c u r v e 7 8 9 _ _emper ica l c u r v e assuming a pKa of - 6 .9 . 121 d i f f e r e n t alcohols do indeed crest at d i f f e r e n t a c i d i t i e s . However, the distance from crest to crest i s not nearly as great as was predicted from the pKa values of these compounds. In l i g h t of a l l the r e s u l t s obtained i t would seem that the best explanation f o r the observed crest i n the rate versus a c i d i t y plots might l i e i n the prediction of a chromate ester protonation step. The pKg^+ of these esters would be close to that of free chromic acid and i t would not vary greatly with changes i n the structure of the organic substrate. Such a suggestion i s r e a d i l y compatible with the previously proposed c y c l i c mechanism fo r t h i s reaction. Thus the rate law i n 5-70$ s u l f u r i c a c i d i s i n d i c a t i n g that the gross composition of the t r a n s i t i o n state must include one molecule of alcohol, one mole-cule of chromic a c i d and a proton. I f i t i s assumed that a chromate ester i s the intermediate, then the species p a r t i c i p a t i n g i n the rate-determining step R0Cr0 3H + H + ===== R0Cr0 3H 2 + Fig.20 Re la t ionsh ip b e t w e e n Ho and t h e r a t e of o x i d a t i o n of t h r e e 2 - p r o p a n o l s in su lphu r i c ac id s o l u t i o n s . P r i m a r y d e u t e r i u m i s o t o p e e f f e c t s i nd i ca ted in p a r e n t h e s i s . 123 TABLE XIII OXIDATION OF CFoCR0HCHo BY C r V A IN SULFURIC ACID SOLUTIONS AT 25.0°C No. R *H 2S0 4 [or"] (xl0 4M) [ale.] (xlO^M) k l ( x l 0 5 s e c a l ) (1 Moles"* 1sec = 1) k H A D 475 H 29.4 10,3 36.5 0.0766 .000021 474 H 45.7 12 .4 22.0 0.738 .000335 495 H 50,1 10.3 74.4 9.07 .00122 494 H 50.1 10.3 37.2 4.21 .00113 465 H 50.1 10.3 37.4 4.36 .00116 435 D 50.1 10.3 36.9 0.370 .000110 10,54 469 D 50.1 10.3 78.0 0.865 .000112 473 H 57,6 10.3 11.0 5.68 .00515 471 H 63.6 6.18 22.8 141.0 .0619 461 H 68.5 6.18 34.4 671.0 ,195 462 H 68.5 6.18 10.8 136.0 ,126 463 H 68.5 6.18 23.2 422.0 .182 466 D 68.5 6.18 33,5 91.0 .0281 5.99 467 H 76.4 6.18 36.2 1060 .293 468 D 76.4 6.18 38,3 198.0 .0516 5.71 472 H 85.7 6.18 22.4 116.0 .0518 455 H 95.3 6.15 17.5 11.1 .00635 457 H 95.3 6.15 18.2 9.40 .00515 458 D 95.3 6.15 70.9 38.3 .00540 1.32 459 D 95.3 6.15 37.7 12.5 .00332 124 TABLE XIV OXIDATION OF CH3GR0HCH3 IN SULFURIC ACID SOLUTIONS AT 25°C No. R $H 2S0 4 [ c r V I ] [ale.] k± k2 kE^k (xl0 3M) (M), ( x l G ^ e c " 1 ) (1 moles^see" 1) 334 H 5.29 4.10 0.129 0.296 0.0023 338 H 10.6 4.23 0.129 0.821 0.00636 333 H 16.0 4.13 0.129 1.87 0.0145 339 H 20.2 4.13 0,129 3.37 0.0261 332 H 25.8 4.10 0.129 7.66 0.0595 379 H 33.4 3.42 0.0409 7.35 0.179 393 H 50.1 1.03 0,0131 55.2 4.21 394 H 50.1 1.03 0.0131 55.8 4.26 395 H 50.1 1.03 0.0131 53.1 4.06 486 D 50.1 1.03 0.0343 23.4 p. 683 487 D 50.1 1.03 0.0214 14.4 0.673 488 D 50.1 1.03 0.0103 8.09 0.626 426 H 97.3 0.0515 0.00964 0.0230 2.39 428 H 97.3 0.0515 0.01326 0.0380 2.86 429 H 97.3 0.0515 0.01820 0.0484 2.63 427 D 97.3 0.0515 0.00888 0.0168 1.89 430 D 97.3 0.0515 0.0169 0.0345 2.04 i 125 TABLE XV OXIDATION OF 1,1,1,3,393=HEXAFLU0R0~2-PR0PAN0L IN SULFURIC ACID SOLUTIONS AT 25°C No. $H 2S0 4 (xlO4!*) [alc.~] (xl02M) k i (xl0 6sec° 1) k 2 = (1 moles" 507 76»4 6.18 3.18 4.46 0.000140 481 79o4 6.18 3.38 4.52 0.000134 493 85.7 4.12 2.96 9.30 0.000314 508 95.3 6.18 3.22 2.02 0.0000628 126 would be I yO. OH + -OH As long as the reaction was carried out below the pKgjj+ of the ester the rate would be f i r s t order i n hQ, However at a c i d i t i e s above the pKgjj+ the rate would be zero order i n h Q . Further protonation would lead to formation of a diprotonated species, R 0^ OH C C r — O H R H OH This ion would not contribute to the rate since there are no unprotonated oxygen atoms avail a b l e f o r p a r t i c i -pation i n a hydrogen transfer step. Hence one would observe a l e v e l l i n g i n the plot of log k v.s. H Q at the pKgjj+ of the ester. Further, the R groups would be expected to have much l e s s e f f e c t on the pKgH+ value of t h i s ester than they would have on the free alcohol. This i s i n accord with the observation that positions of the change i n slope of the plots i n Fig, 20 do not' d i f f e r by as great a magnitude as expected from a consideration' of the probable pKg^+ values f o r the free alcohols. i 127 The s i t u a t i o n i s complicated by the f a c t that an-other species of chromium VI appears i n about 75$ s u l -f u r i c acido Mishra and Symons (99) have suggested that t h i s new species i s a d i s u l f a t e 9 GrOg(OSO^H^o Results which w i l l be presented i n section 9 of t h i s thesis indicate that the chromium VI species e x i s t i n g i n 20 to 55$ H 2 S £ \ * s a monosuirat^e9 H£rS0y 0 I f t h i s species i s f i r s t protonated, as the curve i n Fig» 17 suggests 9 and then further sulfonated to form a new species i n very concentrated acid, the f i n a l species would be a protonated d i s u l f a t e , HOCrO(OS0 3H) 2„ Also at approximately the same place that the new species of chromium VI i s observed to appear the magnitude of the isotope e f f e c t obtained i n the oxidation of 1 ?1 81-trifluoro»2-propanol begins to decrease,. It i s possible, therefore, that t h i s species of chromium VI oxidizes v i a a d i f f e r e n t mechanism than HgCrO^ does* Or i t may be that a d i f f e r e n t step, such as ester formation, becomes rate-determining 0 In t h i s connection i t i s i n t e r e s t i n g to note that while oxidations i n s u l f u r i c a c i d solutions of less than 63o6$ give good straight l i n e plots from the i n i t i a l point on, above 63«6$ s u l f u r i c t h i s i s not so as Figures 21 to 26 i l l u s t r a t e , . In each of these runs chromium VI of a d e f i n i t e concentration i n the various s u l f u r i c a c i d solutions was f i r s t thermostated at 25°C 128 i n a 1 cm. c e l l , alcohol was then added and the o p t i c a l density at 349 m ji was observed at i n t e r v a l s . It can be seen from the plots that an intermediate (or a species i n equilibrium with the intermediate) forms i n i t i a l l y and that t h i s entity then decomposes by a f i r s t order rate law. In 96.3$ s u l f u r i c a c i d a good stra i g h t l i n e from the i n i t i a l point on i s again obtained. Also i n 96.3$ ^2^4 ^ e ^ s o t o P e e f f e c t i s almost unity. (The difference from unity could be due to a secondary isotope effect.) These two f a c t s taken together seem to indicate that the formation of t h i s intermediate i s rate-determining i n very con-centrated ac i d . i . e . If the intermediate i s an ester, decomposition must occur simultaneously with ester formation. The d i s t i n c t i o n between the proposed mechanisms i n 20 to 60$ and i n 96$ s u l f u r i c acid could then be i l l u s t r a t e d as follows. In 20 - 60$ H«S0, : R2CH0H + H2CrO K + H20 (1) H 0Cr0 20H k2 IV •> R2C0 + Cr (2) H Fig.21 Oxidat ion of 1,1,1-trif l uo ro -2 -p ropano l in 5 7 6°/o 1.8 2 + logO.D. 1.4 2 0 0 3 0 0 I Time(min.) I 4 0 0 Fig.22 O x i d a t i o n of 1,1,1-t r i f luoro-2-propanol J . 6 o< in 6 3 . 9 % H 2 S 0 4 . .1.5 2+log O.D. J . 2 J.1 J.O I T ime (mm.) i i i ^ T i m e (min.) J . 8 F i g . 2 4 O x i d a t i o n of 1,1,1-tnf l u o r o - 2 - p r o p a n o l in 7 6 . 4 % H 2 5 0 4 . J 7 J . 6 2 + log O.D. J . 4 J -3 J .2 i T ime (mm.) | i I o< _2.1 F ig .25 O x i d a t i o n of l ^ l - t r i f l u o r o - 2 - p r o p a n o l 85.7 % H 2 S 0 4 . L2.0 ^ J . 9 2+log O.D. 1.8 1.7 1.6 1.5 i T i m e ( m i n . ) I 2 0 F ig 26 O x i d a t i o n of 1 ^ 1 - t n f l u o r o - 2 - p r o p a n o l in 5 0 10CL.. , . J 5 0 I T ime (mm.) i 2 0 0 135 In t h i s case step (1) would be a rapid equilibrium with step (2) being the rate-determining reaction. The rate law would be V - Kk 2 [RgCHOHJ [HgCrOj h Q and the reaction would show a primary deuterium isotope e f f e c t . Since the extinction c o e f f i c i e n t of the chromate ester i s la r g e r than that of chromic a c i d at 349 m p. (52) the equilibrium concentration of the ester must remain low during the course of the reaction. If the ester concentration did not remain low one would observe a sharp increase i n the o p t i c a l density a f t e r introduction of the alcohol into the chromic a c i d solution and the st r a i g h t l i n e obtained by p l o t t i n g log [ c r V I J against time would not extrapolate back through l o g n i t i a l at time zero. This i s i n accord with the observation made by Kwart and Francis that addition of a t e r t i a r y alcohol did not greatly i n h i b i t the rate of oxidation of oC-(£-nitrophenyl)«=ethanol i n a c e t i c acid solutions. If the chromate ester concentration were high much of the t o t a l chromium present would have been t i e d up i n the form of an unreactive t e r t i a r y ester and a large decrease i n the reaction rate would have been observed. 136 In 96$ H 2S0^: + R2CH0H + HOCrO(OS03H) + H +H20 (3) OCrO(OS0 3H) 2 IV R2C0 + Cr (4) H In t h i s case step (3) would be a slowly es-tablished rate c o n t r o l l i n g equilibrium, step (4) would primary deuterium isotope e f f e c t when the oC-hydrogen atom of the alcohol was substituted with deuterium, but could however, show a secondary deuterium isotope e f f e c t associated with the ester forming step e From a consideration of Figures 21 to 26 i t i s apparent that i f t h i s i s a correct picture of the mechanism the equilibrium concentration of ester must become greater i n very concentrated acid solutions and the rate at which t h i s intermediate i s formed must become slower and slower u n t i l i n 96$ H 2 S 0 ^ the rate of decomposition i s equal to the rate of ester formation,, 6. DETERMINATION OF THE VALUE FOR THE CF 3 GROUP Taft (100) has indicated that the rates of a be f a s t and the rate law would be V «= k^ This mechanism would not lead to a 137 wide var i e t y of reactions of a l i p h a t i c compounds of the type RY f i t with good p r e c i s i o n the equation l o g k/k Q = e* where o i s the polar substituent constant f o r the group R, e* i s a constant g i v i n g the s u s c e p t i b i l i t y of a given reaction series to polar substituents, k i s the rate constant f o r the reaction of RY and k i s the o rate constant for the reaction of CH^Y. Since i t was of some i n t e r e s t i n the present i n v e s t i g a t i o n to know the value of the polar substituent constant f o r the CF^ group, the rates of oxidation of a series of primary alcohols of the type RCHgOH were determined at 25°C i n 3.82 M HCIO^. (Roc'ek (32) had previously observed that the rates of oxidation of a number of alcohols of t h i s type i n 1.0M HCIO^ at 20°C could be f i t t e d to Taft's equation.) The r e s u l t s obtained i n the present work f o r the oxidation of n-propanol, ethanol, methanol, 2-methoxyethanol, 2-chloroethanol, 2,2,2-trifluoro-ethanol and isopropyl alcohol are l i s t e d i n Table XVI 0 A p l o t of the log of the rate constant against 6s values f o r the f i r s t f i v e alcohols mentioned was then made and the best straight l i n e was extrapolated to log k f o r 2,2,2-trifluoroethanol. As the plot i n 13$ i f w Figure 27 shows t h i s procedure gives a 6> value f o r the CF^ group of about +3o9. I t i s of i n t e r e s t to note that a good c o r r e l a t i o n with the oxidation rates of substituted 2-propanols in 50.1$ s u l f u r i c acid i s also observed when t h i s value f o r the CF^ group i s used (See Figure 1 5 ) . Figure 27 indicates that the oxidation of methanol i s anomonously slow (or that the <^  value f o r the H substituent i s too low). This has also been observed by Roc'ek i n his work (32). He has suggested that t h i s anomaly may be due to the f a c t that i t i s a methyl group which i s oxidized i n methanol while i n the other primary alcohols a methylene group forms the point of attack. 7. DETERMINATION OF THE "TRUE RATE CONSTANT", k 2 -k . /K FOR THE CHROMIC ACID OXIDATION OF SOME PRIMARY ALCOHOLS If a chromate ester i s assumed to be an i n t e r -mediate i n the chromic a c i d oxidation of alcohols the complete rate law for the reaction i n moderately concentrated a c i d i s , V = Kk 2 [H 2 C r 0 4 j |R0HJ h Q where K i s the ester equilibrium constant and k 2 i s the rate constant f o r ester decomposition. Thus the 139 TABLE XVI OXIDATION OF SUBSTITUTED METHANOLS IN 3.82 M HCIO, (H -4 o =-1.62), 1.03xl0°3M, T=25.0°C. Alcohol No. [ale.] (xlO 2 ) k ^ s e c " 1 ) kg ( l moles" 1 n-Propanol 551 1.34 0.0131 97.7 552 3.12 0.0285 - 91.3 Ethanol 554 5.72 0.0435 76.0 555 4.56 0.0329 72.2 2-Methoxy- 559 4.24 0.00623 14.7 ethanol 560 2.12 ,0.00322 15.2 2-Chloro- 547 5.03 0,00311 6.20 ethanol 549 4.05 0.00232 5.74 Methanol 556 4.15 0.00175 4.22 557 8.30 0.00366 4.41 T r i f l u o r o - 558 68.8 0.000023 0.0033 ethanol 561 45.9 0.000018 0.0039 Isopropyl 367 4.09 0.0268 61.7 alcohol 140 bimolecular rate constants ( k Q b s ) which are usually-reported f o r t h i s reaction are a c t u a l l y a product of two constants. The "true rate constant" f o r the rate-determining ester decomposition would be given by, kg = k Q D S /K a n ( * could only be evaluated i f an i n -dependent determination of the ester equilibrium could be made. Fortunately Klaning (52) has developed a method f o r the determination of t h i s equilibrium constant i n d i l u t e acid solutions and Klaning and Symons (43) have l i s t e d the values obtained f o r a large number of alcohols. In the same manner (as described i n the experimental section) we have de-termined the equilibrium constant f o r 2-methoxy-ethanol and 2 S )2 i )2-trifluoroethanol, By use of these equilibrium constants i t was possible to determine the "true rate constants" (kg) f o r the oxidation of si x primary alcohols. The pertinent data i s l i s t e d i n Table XVII. A plot of log kg against Taft's <r values re-vealed that the € value f o r the rate-determining step f o r t h i s reaction i s negative (Figure 27). This i s i n accord with the suggestion previously made that the rate-determining step involved hydride transfer and casts serious doubts on the explanation o r i g i n a l l y given by Kwart and Francis (27) f o r the 141 TABLE XVII CALCULATION OF THE "TRUE RATE CONSTANTS", k2 " k o b s / K F 0 R S 0 M E SUBSTITUTED METHANOLS IN AQUEOUS PERCHLORIC ACID SOLUTIONS Alcohol k 2 x 102 (1 moles" sec" ) K k 2/K n~Propanol 94.5 9.7 9.74 -0.100 Ethanol 74.1 6.9 10.75 0.000 2-Methoxy-ethanol 15.0 8 .3 1.81 0.520 2~Chloro= ethanol 5.97 5.9 1.01 1.050 Methanol 4.32 4.7 0.92 0.490 2 , 2 , 2-Trifluoro - 3 . 6 x l O ~ 3 ethanol 1 .3 2.77xlO" 3 3.9 Isopropyl alcohol 61.7 142 - E t -_1 \ M e O C H 2 -_ 0 o X c i C H o -H - \ 2+log k , \ 9 = K o b s . \ -1.11 _=} _z2 0 1 * 2 1 I Of I 3 \ C F 3 ~ 5 I ^ i F ig .27 R e l a t i o n s h i p b e t w e e n k o t o s and or* f o r t h e o x i d a t i o n of s u b s t i t u t e d methano ls . F i g . 2 8 Re la t i onsh ip b e t w e e n t h e " t r u e r a t e c o n s t a n t " and o * . 143 negative nrho r t value observed i n the oxidation of arylmethylcarbinols. Since the ester mechanism had received general acceptance they suggested that the observed negative w r h o n i s a r e s u l t of changes caused by the substituents i n the position of the ester equilibrium. It should be noted, however, that while these r e s u l t s to a large extent i n v a l i d a t e the contentions of Kwart and Francis, the suggestion made by Westheimer (22) and by Wiberg (2) that the negative "rho" i s merely a consequence of the removal of two electrons as well as a proton s t i l l remains as a p o s s i b i l i t y . g. ACIDITY FUNCTIONS FOR NITRIC AND PHOSPHORIC ACID SOLUTIONS WITH ADDED SODIUM PERCHLORATE In the course of the i n v e s t i g a t i o n of the mechanism of chromic a c i d oxidations i t became necessary to de-termine the a c i d i t y functions f o r n i t r i c and phosphoric acid solutions containing added sodium perchlorate. In each case a 1:1 molar r a t i o of acid to s a l t was used. The functions were determined by use of weak bases for indicators i n the manner described by Paul and Long (91). The derived a c i d i t i e s f o r these solutions are presented i n Tables XVIII and XIX and the functions are plotted i n Figures 29 and 30. The a c i d i t y 144 TABLE XVIII ACIDITIES OF SOLUTIONS CONTAINING A 1:1 MOLAR RATIO OF PHOSPHORIC ACID AND SODIUM PERCHLORATE HOP0, - NaCIO, H 3 4 4 o (M) 0.25 1.15 0.50 0.81 1.00 0.48 1.50 0.14 2.00 -0 .17 2.50 -0.47 3.00 -0.76 3.50 -1.05 4.00 -1 .34 4.50 -1 .64 145 TABLE XIX ACIDITIES OF SOLUTIONS CONTAINING A 111 MOLAR RATIO OF NITRIC ACID AND SODIUM PERCHLORATE HNO-a - NaCIO, H 3 4 o (M) 0.25 0.55 0 .50 0.12 1.00 -0.35 1.50 -0.69 2.00 -0.99 2 .50 -1.28 3.00 =1.58 3.50 -1.88 4.00 =2.17 F i g . 2 9 C o m p a r i s o n of t h e Ho func t ion f o r a 1:1 r a t i o of H 3 P 0 4 and N a C I 0 4 (o) w i t h t h a t _-1 f o r H 3 P Q 4 Ho 0 1 [ H 3 P 0 4 - N a C I O j (M) 3 4-4 F i g . 3 0 C o m p a r i s o n of t h e Ho f u n c t i o n f o r a 1:1 r a t i o o f HNO3 a n d N a C ' 0 4 (o) w i t h t h a t f o r 0 HN0 3 - N a C!0 4 3 4 148 functions f o r phosphoric and n i t r i c acid solutions containing no added s a l t s are included i n these figures f o r comparison purposes. While no studies i d e n t i c a l to t h i s one have previously been conducted Harbottle (101) and Paul (102) have reported that the addition of varying amounts of sodium perchlorate to perchloric a c i d and hydrochloric acid solutions r e s u l t s i n an increase i n the a c i d i t y of the medium. Paul and Long (91) have suggested a possible reason f o r t h i s increase. They point out that -H Q can be defined as, f H + -H 0 - l o g C R+ + l o g - y ^ p * log f B and that while studies of s a l t e f f e c t s on nonelectro-l y t e s indicate that the acids hydrochloric, n i t r i c , s u l f u r i c and perchloric have l i t t l e tendency either to s a l t i n or s a l t out, the sodium s a l t s of these acids show d e f i n i t e tendencies to do so. From t h i s they deduce that the more rapid increase i n -H as compared to l o g CJJ+ i n acid solutions containing no f -f added s a l t i s due to an ef f e c t on the log H term. This they suggest i s plausible since studies have indicated that the influence of e l e c t r o l y t e concentra-t i o n on the a c t i v i t y c o e f f i c i e n t of a small cation can 149 become very large r e l a t i v e to that of a large cation. The addition of neutral s a l t s would s a l t out the non-el e c t r o l y t e causing an increase i n the (log fg) term and as i n the previous case increasing the e l e c t r o -l y t e concentration would have an e f f e c t on the log ^H* term. f H B + The functions obtained i n the present i n v e s t i g a t i o n are i d e n t i c a l with the previously determined functions containing no added s a l t s i n d i l u t e regions. This i s reasonable since the addition of small amounts of added s a l t s should have very l i t t l e e f f e c t on the a c i d i t y of d i l u t e aqueous solutions. In more con-centrated solutions.however, the addition of i n -creasing amounts of added s a l t s r e s u l t s i n a continual increase i n a c i d i t y as compared to solutions containing no added sdat. This would seem to be consistent with the suggestions previously advanced by Paul and Long. 9. VARIATION IN THE pKa OF CHROMIC ACID WITH CHANGES IN THE IDENTITY OF THE MINERAL ACID SOLVENT The apparent pKa of chromic a c i d was determined, as has been described, by p l o t t i n g ,, eHCrO " - G log 3 against H Q or H_ f o r a variety of mineral acids. The 150 r e s u l t s , presented i n Table XX, indicate that the value of -1.01 obtained i n perchloric acid i s i n good agreement with the value of -0.98 previously determined by Bailey, Carrington, Lott and Symons (80). The sur p r i s i n g aspect of the r e s u l t s , however, i s that the apparent pKa's are not constant from one acid medium to another. The v a r i a t i o n i n pKa i s very large and coupled with the evidence that w i l l be presented i n the next section completely rules out the p o s s i b i l i t y that the v a r i a t i o n could be due to experimental error. In t h i s connection i t i s also noteworthy that the reported pKa of permanganic acid d i f f e r s from perchloric to s u l f u r i c a c i d media (72). The protonation of organic indicators such as the cyanocarbon anions i s known to occur by means of a conventional equilibrium step involving a cyanocarbon anion, the corresponding cyanocarbon acid and a proton (92). C(CN)~ + H + HC(CN) 3 However, the r e s u l t s presented i n Table XX indicate that i n the protonation of an inorganic indicator such as chromate anion or permanganate anion some consider-ation must also be given to the proton source, i . e . The equilibrium must be established i n such a way that there i s a d i r e c t involvement of the mineral acid which 151 TABLE XX COMPARISON OF THE pKa OF HgCrO^ IN AQUEOUS SOLUTIONS OF VARIOUS MINERAL ACIDS AND THE POINT OF SLOPE CHANGE IN A PLOT OF log K v.s. H Q FOR THESE ACIDS Acid pKa H scale Point of H. scale s l o p e c h i HN03 -1.91 -2.45 HC10 4 -1.10 -0.83 -1.25 H 2S0 4 +0.51 +0.34 +0.02 HCI +0.76 +0.52 +0.59 H 3P0 4 +1.74 +1.10 HoP0,-NaC10, +1.38 +1.35 (1:1 molar r a t i o ) * This value was determined i n these laboratories by-Miss M. M. Wei. 152 furnishes the protons. HCrO^" + 2H + + A" ; * HCrC^A + H20 Further, i t has been observed that at a c i d i t i e s below the apparent pKa of chromic acid the spectral VI character and o x i d i z i n g a b i l i t y of Cr solutions are independent of the proton source. However, at a c i d i t i e s greater than the pKa of chromic acid both spectral character and o x i d i z i n g a b i l i t y are found to depend on the i d e n t i t y of the mineral acid present. This i s i l l u s t r a t e d by the data presented i n Tables XXI, XXII and XXIII and by the spectra reproduced i n Figures 31-34. This evidence would seem to indicate that the ion CrO^A" i s not present i n s i g n i f i c a n t concentra-tions under conditions where the monoanion i s the dominant species.but that HCrO^A i s the major neutral „ VI Cr species. H 20 + CrC^A" H C r ( \ ~ + H A (i) H 20 + HCr0 3H H 2 C r G 4 * H A i . e . Equilibrium (i) would seem to be to the r i g h t but equilibrium ( i i ) to the l e f t . Indications are that the following species are formed i n the protonation, of HCrO,~ by the various 153 TABLE XXI SPECTRAL DATA FOR THE ACID CHROMATE ION IN VARIOUS MINERAL ACIDS Mineral Concentration A max 6 Acid (M) HCIO^ 1.74 x 10' 2 430 230 349 1584 257 2140 H 2S0 4 pH - 2.95 430 242 3*9 1572 258 2170 HCI 1.0 x 10~ 3 430 219 349 1528 257 2070 H 3P0 4 1.20 x 10" 2 435 245 349 1616 257 2200 154 TABLE XXII SPECTRAL DATA FOR CHROMIC ACID IN VARIOUS MINERAL ACIDS Mineral Concentration ^ max G Acid (M) HC10 4 3.50 M 440 193 347 1158 265 1760 H 2S0 4 3.10 M 445 184 349 1028 265 2050 HCI 3.60 M 445 176 354 1240 283 1699 248 2110 H 3P0 4 3.00 M 445 198 348 1090 275 1920 155 TABLE XXIII RATE OF OXIDATION OF ISOPROPYL ALCOHOL IN MINERAL ACIDS OF A GIVEN ACIDITY. (H = -2.5) Mineral k o b s ^ m ° l e s ~ ^ s e c " ^ ) Acid HC1 7.1 x IO" 3 H3PO^ 0.28 H 2S0 4 0.63 HC10 4 5.0 HN03 18 Fig.31 U l t rav io le t s p e c t r u m of C r ^ - in 3.50 H C I 0 4 Fig. 32 U l t r a v i o l e t s p e c t r u m of Cr m in 3.1 O M H 2 S 0 4 Fig.33 U l t rav io le t s p e c t r u m of C r ^ - in 3 . 6 0 M HCI Fig .34 U l t r a v i o l e t s p e c t r u m of Cr m in 3 . 0 0 M h U P O 160 mineral acids. HCr0 4° + H + H 3P0 4 0 ll HOCr-II 0 0P0 3H 2 + H 20 0 HCrO, "* + H "fr HC1 » HOCr-4 II 0 0 . + II HCrO^ + H + H 2S0 4 > HOCr-0 0 HCrO, " + H + HC10, > HOCr-4 4 || 0 0 - + 11 HCrO, + H + HNOo > HOCr 4 3 || 0 CI + H 20 -OS03H + H 20 •0C103 + H20 •0N02 + H20 The postulation of the formation of these species i s not completely without precedent since i n addition to the aforementioned suggested formation of CrO-jCl" (34), Holloway (103) has also found spectroscopic evidence f o r the formation of a complex between chromic and phosphoric acid and Chang (104) has obtained spectro-scopic evidence f o r CrS0y = or i t s protonated analogs. The well known case of complex formation between two acid chromate anions to give a dichromate anion can also 161 be considered as another example of t h i s general re-action. 2HCr0 4~ * CrgOy"" + HgO Further, i f t h i s i s a correct picture one would expect the pKa's of these species to increase with decreasing electron withdrawing power of the A groups. For example, since the b i s u l f a t e group has greater electron withdrawing power than chloride one would expect the pKa f o r HgCrSOy to be more negative than the pKa of HCrO^Cl. i . e . The greater the electron withdrawing power of the A group the l e s s available f o r protonation w i l l be the electrons associated with the oxygen atom attached to chromium. Since the strengths of the mineral acids, HA, also vary i n a s i m i l a r way with electron withdrawing power of the A groups one would expect that a c o r r e l a t i o n should e x i s t between the strengths of the mineral acids and the corresponding HCrO^A species. Such a c o r r e l a t i o n with acid strengths i s seen to exist f o r a l l the mineral acids investigated with the exception of HNO^. i . e . The strengths of these acids increase i n the order H 3P0 4 < HN03 < HCI < HgSO^ < HCIO^ (91) while the-pKa's increase i n the order H3CrPOy < HCr0 3Cl < H 2CrS0 ? < HCrClOy < HCrN0 6. 162 The apparent anomaly of the HCrNO^ species i n t h i s respect i s not as yet completely understood. It i s possible however, that the bonding i n t h i s species i s d i f f e r e n t from the bonding i n the other cases. For example, HCrNO^ may ex i s t as an oxygen bridged compound 0 0 HOCr N - 0. 0 0 I f t h i s i s so i t would be expected VI that Cr would have a considerably d i f f e r e n t u l t r a -v i o l e t spectrum i n HNO^  as compared to other mineral acids. However, n i t r i c a c i d i t s e l f has an absorption band at 350 m ji blanking out the most i n t e r e s t i n g part of t h i s spectrum. The spectra of chromic acid i n H^PO^, HgSO^ and HCIO^ a r e a l l very s i m i l a r , as would be expected i f a l l of these species contained chromium-oxygen bonds. The spectrum of chromic acid i n HC1, on the other hand, shows considerable v a r i a t i o n from the spectra obtained i n H^PO^, H 2S0 4 and HCIO^. (See Figures 31-34). The explanation f o r t h i s v a r i a t i o n being that i n the pa r t i c u l a r case of hydrochloric acid a chromium-chlorine bond i s substituted f o r a chromium-oxygen bond. The bonding i n these species could be considered to be mostly covalent i n analogy with the bonding i n 163 chromium compounds such as chromyl chloride (105), chromyl bromide (105), and chromyl acetate (15). ¥hile covalent perchlorate compounds are not widely known some have recently been prepared (107, 108) and i t i s conceivable that even the bonding i n HCrClOr, i s covalent i n nature to a considerable extent. In t h i s connection i t should be noted that addition of sodium perchlorate to solutions of phosphoric and n i t r i c acid did not greatly a f f e c t the observed pKa of chromic acid i n these solutions. This indicates that perchlorate w i l l not replace either phosphate or n i t r a t e i n these compounds. In addition to the .above considerations i t follows that i f the protonation equilibrium involves a molecule of mineral a c i d a plot of log ejjcr*0 ° - e against eith e r a simple H Q or H = function should not necessarily give a straight l i n e of unit slope. The function that should give «rvise ,1tova l i n e a r p l o t of unit slope can be derived i n the following way: HCrO^" + H + + HA === HCr0 3A + HgO K = ( a H C r Q ^ ) (a H+) ( a H A )  taHCrO,A) ( aH o0) 164 pK = log XL - log aH + fHCrO, " log aHA CHCrO HCr0oA pK - log CHCrCUA + H. + log HA Hence a plot of log 6 HCrO against the function, e 6 a H. + log 2 , should y i e l d a straight l i n e of unit slope. Such a function has been calculated f o r HC1, HgSO^ and HCIO^ using the value f o r the logarithm of the a c t i v i t y of water given i n r e f . (49) and the mean a c t i v i t y c o e f f i c i e n t s f o r the el e c t r o l y t e s l i s t e d i n r e f . (109). The plots f o r HC1 and HgSO^ i l l u s t r a t e d i n Figures 35 and 36 indicate that good straight l i n e plots with unit slopes are obtained when t h i s function i s used. By comparing these plots with those i n Figures 3 and 5 where the simple H Q and H„ functions are used i t i s evident that the slopes obtained when t h i s new function i s used are much closer to unity. This r e s u l t serves to further j u s t i f y the previous assumption that there i s a d i s t i n c t difference between the protonation reaction of organic and inorganic bases of t h i s type. In the l a t t e r case a molecule of the proton supplying aHA 165 0) I L. ( J X CL) Fig. 3 5 " p K a " of H 2 C r S O y 0.6 0.0 S l o p e = 0 . 9 6 x I C L ) +2 +1 H_ + l og a H 2 0 ; a H A O H_+log a H 2 Q / a H A F ig .36 " p K a " of H C r 0 3 C I . / _0.6 _ O 0 O / ^ S l o p e ^ 0 .99 z.0.6 / /O i +1 i 166 mineral a c i d becomes incorporated into the species when protonation occurs. This has not been observed f o r the organic bases. In the case of perchloric acid the slope obtained with the simple H„ function i s almost unity (Fig. 4) and ap p l i c a t i o n of the new function resulted i n no improve-ment. U n a v a i l a b i l i t y of the required data prevented the cal c u l a t i o n of t h i s new function f o r the n i t r i c and phosphoric acid systems. 10. VARIATION IN RATE OF CHROMIC ACID OXIDATIONS WITH CHANGES IN THE IDENTITY OF THE MINERAL ACID SOLVENT If the picture of the chromic acid equilibrium presented i n the preceding section i s correct one would expect that the physical and chemical properties of chromic acid would vary somewhat with the i d e n t i t y of the mineral acid solvent. It has previously been pointed out that while the acid chromate spectrum i s i d e n t i c a l f o r a l l the mineral acids investigated (Table XXI) the spectrum of chromic acid i t s e l f varied somewhat from acid to acid (Figures 31-34). S i m i l a r l y i t was found that the rate of oxidation of isopropyl alcohol by the acid chromate anion depended only on the a c i d i t y of the medium and not on the nature of the proton supplying mineral a c i d . On the other hand i t 167 has been observed that the rate of oxidation by-chromic acid depends not only on the a c i d i t y of the medium, but also on the i d e n t i t y of the mineral a c i d . These observations are i l l u s t r a t e d i n F i g . 37 where the logarithm of the rate constant f o r oxidation of isopropyl alcohol by chromium VI i s plotted against H Q f o r ^ P O ^ , HCI, H 2S0 4, HCIO^ and HNCy In addition, points are plotted f o r solutions of H^PO^ and HNO^  containing a 111 molar r a t i o of sodium perchlorate. It can be seen that the addition of sodium perchlorate to these solutions had no e f f e c t on the rate of oxidation outside of i t s e f f e c t on the a c i d i t y of the medium. It i s i n t e r e s t i n g to note that a sharp change i n slope i s observed f o r each p a r t i c u l a r mineral acid sol u t i o n very near to the pKa of chromic acid i n that medium (Table XX). Apparently the rate of oxidation by the protonated species of chromium VI varies with the a c i d i t y of the solution i n a d i f f e r e n t manner than does the rate of oxidation by the unprotonated acid chromate anion. The observation that a v a r i a t i o n i n oxidation rate does not appear u n t i l a f t e r the pKa VI of the Cr species supports the suggestion that t h i s equilibrium i s such that a molecule of mineral acid does not become incorporated i n the monoanion. 168 F i g . 3 7 C h r o m i c ac i d ox i da t i on of i s o -p ropy l a l coho l in a q u e o u s s o l u t i o n s of m i n e r a l ac ids . 169 TABLE XXIV RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN SULFURIC ACID SOLUTIONS AT 25°C. No. [H 2S0J (M) Ho [ c rv i ; (MxlO 3) [ A l e ] (M) k l l ( s e c - 1 ) " 2 (1 moles sec" 1) 3 4 3 2 . 6 2 4.07 0.129 6 . 4 8 x l 0 " 8 5 . 0 3 x l O " 7 342 0 . 0 1 0 5 1.98 4.10 0.129 6 . 0 9 x l O " 7 4.95xlO" 6 3 3 7 0.0214 1 . 7 7 4 . 2 7 0.129 9 . 4 7 x l O " 7 7 . 3 4 x l O " 6 341 0.0464 4.17 0.129 3.68xlO" 6 2.85x10"5 335 0.116 0.78 4 . 1 0 0.129 3 . 1 4 x 1 0 " 5 2.43xlO" 4 340 0.24 0.44 4 . 0 7 0.129 1.03xlO" 4 7 . 9 8 x l O ~ 4 334 0.54 0.08 4.10 0.129 2.96xlO" 4 2.30xlO" 3 338 1.06 -0.35 4.23 0.129 8.21X10"4 6 . 3 6 x l O " 3 333 1 . 6 2 -0.72 4.13 0.129 1.87xlO" 3 1.45xlO" 2 339 2.26 -1.02 4.13 0.129 3 . 3 7 x l O " 3 2 . 6 l x l 0 " 2 3 3 2 3.04 -1.43 4.10 0.129 7.66xlO" 3 5.95xlO" 2 379 4.19 -1.95 3.46 0.0409 7.35xlO" 3 1.79X10"1 393 7.10 - 3 . 3 7 1.03 0.0131 5.52xlO" 2 4.21 394 7.10 - 3 . 3 7 1 . 0 3 0.0131 5.58xlO" 2 4.26 395 7.10 - 3 . 3 7 1.03 0.0131 5.31xlO" 2 4 . 0 6 1 170 TABLE XXV RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN PERCHLORIC ACID SOLUTIONS AT 25°C. No. [HC10J (M) Ho (MxlO3) [Ale] (M) k l 1 (sec"1) k 2 (1 moles sec"1) 370 0.0301 1.50 3.42 0.129 2.69xlO"6 2 .08xlO° 5 365 0.192 0.70 3.42 0.129 4.29xlO"5 3.32xlO"4 361 0.463 0.20 3.42 0.129 2.13xlO"4 1 .59xlO° 3 364 0.93 -0.18 3.42 0.129 9.20xlO"4 7.13xlO"3 362 2.08 -0.82 3.60 0.129 7.44xlO"3 5.76xlO"2 363 2.63 -1.05 3.42 0.129 1.82xlO"2 1.41X10"1 366 2.98 -1.21 3.42 0.0409 9.8lxlO" 3 2 . 4 0 X 1 0 " 1 367 3.82 -1.62 3.42 0.0409 2.68xlO"2 6.56X10"1 368 4.53 -1.99 3.42 0.0409 6.57xlO'2 1.71 376 5.34 -2.44 0.342 0.00873 3.51xlO"2 4.03 1 TABLE XXVI RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN PHOSPHORIC ACID SOLUTIONS AT 25°C No Ho [ A l e ] , k l -1, (M) (MxlO^ ) (M) (sec L ) (1 moles sec"*-1-) 386 D i l . 2.45 3.42 0.129 1.54xlO° 7 1.19xlO" 6 385 D i l . 2.22 3.42 0.129 2.43xlO" 7 1.88xlO" 6 382 0.504 0.97 3.42 0.129 1.51x10"5 1.17xlO~ 4 392 0.903 0.69 3.42 0.129 2.88xlO~ 5 2.23xlO" 4 1 171 373 3.71 0.01 3.42 0.129 1.39x10" •4 1.08x10" •3 374 3.61 -0.25 3.42 0.129 2.81x10" •4 2.17x10" •3 375 4.51 -0.53 3.42 0.129 5.60x10° •4 4.34x10" =3 371 5.44 -0.82 3.42 0.129 1.16x10° •3 9.04x10° •3 377 6.70 -1.32 3.42 0.129 3.66x10' •3 2.84x10" -2 380 10.00 -2.59 3.42 0.0409 9.96x10" •3 0.244 383 13.16 -3.30 3.42 0.0384 5.82x10" •2 1.52 TABLE XXVII RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN HYDROCHLORIC ACID SOLUTIONS AT 25°C. No. [HCI] Ho & rv i : [ A l e ] k2 _ (M) (MxlO- ) (M) (sec" 1) (1 moles" sec-1) 519 0.009 2.06 3.42 0.129 4.65x10" •7 3.60xlO" 6 529 0.15 0.81 3.42 0.129 2.89x10" •5 2.24x10" 4 513 0.40 0.33 3.42 0.129 9.56x10" -5 7.40xl0" 4 514 2.0 -0.69 3.42 0.129 1.98x10* •4 1.53xlO° 3 512 3.6 -1.26 3.42 0.129 3.58x10" .4 2.97xlO" 3 515 6.1 -2.17 3.42 0.129 1.39x10" •3 1.08xlO" 2 1 172 TABLE XXVIII RATES OF OXIDATION OF ISOPROPYL ALCOHOL " IN NITRIC ACID SOLUTIONS AT 25°C. No. J H N O 3 ] Ho [ c rv i ; (MxlO 3) [ A l e ] (M) k l l (sec 1 ) k 2 (1 moles sec.""1) 523 0.245 0.56 3 . 4 2 0.129 6.67x10" • 5 5 . l 6 x l 0 ~ 4 520 0.724 0.00 3 . 4 2 0.0436 1.56x10" . 4 3 . 8 0 x l O " 3 521 2 . 4 5 -0.82 3 . 4 2 0.129 9.73x10" • 3 7 . 5 5 x l O " 2 518 5 . 7 1 -1.74 3 . 4 2 0.0436, 6.45x10" •2 1.61 522 4.08 -1.34 3.42 0.129 6 . 0 5 x 1 0 " •2 4.69X10"1 525 7 . 5 5 -2.10 3.42 0.00513 6.01 526 9.42 -2.48 3.42 0.00513 17,95 527 11.31 -2.86 3.42 0.00513 27.0 524 14.2 - 3 . 4 5 3.11 0.00467 63.4 1 173 TABLE XXIX RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN SOLUTIONS CONTAINING A 1:1 MOLAR RATIO OF ^PO^ to NaCIO No. [H3P04-NaC104; [ c r V I ] [Alc.J. k l k 2 (M) (MxlO 3) (M) (sec" 1) (1 moles"' sec" 1) 542 0.00914 2 .30 3.42 0.129 2.81x10" 7 2.18xl0" 6 541 0.0457 1.78 3.42 0.129 1.69x10" •6 1.31x10"5 540 0.183 1.27 3.42 0.129 6.12x10' •6 4.74xlO" 5 539 0.457 0.84 3.42 0.129 1.40x10" •5 1.085x10" 535 1.37 0.18 3.42 0.129 8.14x10" •5 6.30x10" 4 534 2.28 -0.35 3.42 0.129 3.20x10" •4 2.48xlO" 3 533 4.11 -1.40 3.42 0.129 6.18x10" •3 4.79xlO" 2 1 4 TABLE XXX RATES OF OXIDATION OF ISOPROPYL ALCOHOL IN SOLUTIONS CONTAINING A 1:1 MOLAR RATIO OF HN03 to NaCIO No. [HNO-^-NaClO^ (M) 2 H o i c rv i ; (MxlO' [Ale] !)(M) (sec**1) K 2 (1 moles sec." 1) 575 0.0417 1.40 3.42 0.129 3.28xlO" 6 2.55xl0" 5 574 0.0417 0.22 3.42 0.129 3.15xlO" 4 2.44xlO° 3 573 1 .54 -0.72 3.42 0.129 7.46xlO" 3 5.78xlO" 2 572 3.08 -1.63 3.42 0.129 5.36xlO" 2 1.23 571 4.02 -2.18 3.42 0.129 8.92 1 174 The o x i d i z i n g a b i l i t y of the protonated species, HCrCXjA, increases i n the order HgCrPOy < HCrClC^ < H 2CrS0y "(HCrClOy ^HCrNO^. This i s the same order i n which the pKa*s of these species vary. Such a corre-l a t i o n i s not unexpected since nkn groups which are electron withdrawing w i l l decrease the ease of protonation of the species, but increase the tendency of the chromium species to accept electrons from a reducing agent. From a consideration of the r e l a t i v e electron withdrawing power of the HPO^-, CI", HSO^-, ClO^- and NO^" groups the p o s i t i o n of the HCrNO^ species i n the above series i s anomalous i n that one would not expect the NO^ " group to be the most strongly withdrawing group i n the s e r i e s . This anomaly was also noted i n the case of the r e l a t i v e pKa's of these species and i s not as yet completely understood. The slopes of the l i n e s i n F i g . 37 are also of some i n t e r e s t . The slope of the plot of H Q against the logarithm of the rate constant f o r oxidation by chromate ion i s about 1.8, while the slopes f o r oxidation i n regions beyond the pKa f o r chromic acid i n the various media are much lower, being i n most cases close to unity. It has previously been suggested by Westheimer (37) that the high slope i n d i l u t e acid region a r i s e s from the f a c t that the chromate ester 175 can decompose either with or without involvement of a proton. R2CHOH + HCrO^" + H + R2CHOCr03H + H20 k a R2CHOCr03H » R2CO + H 2Cr0 3 k b R 2CHOCr0 3H + H + > R2CO + H 3CrC> 3 + V = K ;R 2CHOH] [HCr04-] ( ^ | y ] + ^ ^ 2 ) [H20] In the more a c i d i c regions, beyond the pKa of chromic a c i d , the following mechanism i s consistent with the r e s u l t s obtained i n t h i s as well as other investigations. K R2CHOH + H0Cr0 2A R 2CH0Cr0 2H + H 20 k 2 R 2CH0Cr0 2H + H* > R2CO + H 2CrG 2A + V. CONCLUSION The ultimate goal i n the study of any chemical r e-action i s a complete determination of the mechanism of 176 that reaction. A mechanism simply refers to the path that the molecules follow i n going,from reactants to products. Complete mechanism elucidation involves determination of the movements and i n t e r a c t i o n of a l l p a r t i c l e s (atoms, ions and electrons) during the course of the reaction as well as an estimate of the energy and configuration of every intermediate or t r a n s i t i o n state formed. While such a complete mechanism elucidation has not been obtained f o r the chromic a c i d oxidation reaction, many features of the reaction have been investigated and i t i s possible to estimate on the basis of chemical knowledge what form the unknown features of the reaction w i l l take. The oxidation of a l l alcohols used i n t h i s i n -vestigation showed a very pronounced primary isotope e f f e c t when the oc-hydro gen was substituted with deuterium. The only case i n which such an isotope e f f e c t was not observed was f o r oxidations carried out i n 96$ s u l f u r i c a c i d . These observations i n -dicate that the rate-determining step of t h i s reaction involves carbon-hydrogen bond f i s s i o n except i n very concentrated s u l f u r i c acid solutions. In addition the w r h o n value f o r the rate-determining step of the reaction was found to be negative i n d i c a t i n g that considerable p o s i t i v e charge resides on the oc-carbon 177 atom i n the t r a n s i t i o n state. These f i r s t two obser-vations taken together lead one to the conclusion that the rate-determining step must involve some form of hydride ion transfer, since only i n t h i s way can a carbon-hydrogen bond be broken and leave a re s i d u a l p o s i t i v e charge on the carbon atom. While these considerations are not as straight forward i f one considers the p o s s i b i l i t y of a c y c l i c i n t e r n a l hydrogen transfer as the rate-determining step, one i s s t i l l l e f t with the conclusion that the reaction must occur mainly by hydride transfer. A review of the l i t e r a t u r e reveals that there i s considerable evidence i n d i c a t i n g that chromate esters play the r o l e of intermediates i n the chromic a c i d oxidation of secondary alcohols. For example; (i) d i i s o p r o p y l ether (where no ester formation i s possible) i s oxidized only about 1/1500 as r a p i d l y as isopropyl alcohol under the same experimental conditions (22), ( i i ) c i s - l , 2 - d i o l s are oxidized much more r e a d i l y than t r a n s - l , 2 - d i o l s (24, 25, 26), ( i i i ) chromate esters have been observed to form at l e a s t under some experimental conditions (42) and (iv) i n the oxidation of certain highly hindered alcohols the deuterium isotope e f f e c t becomes unity i n d i c a t i n g that the e s t e r i f i c a t i o n reaction has become rate 178 c o n t r o l l i n g (104). The r e s u l t s of the present i n -vestigation are also consistent with the ester mechanism. A study of the rates of oxidation of three substituted 2-propanols i n solutions of con-centrated s u l f u r i c a c i d has indicated that the rate maxima found around 70$ ^SO^ i s probably associated with the protonation of the chromate ester at t h i s point. It has further been suggested that the disappearance of an appreciable isotope e f f e c t i n very concentrated s u l f u r i c acid solutions i s due to rate-determining ester formation i n t h i s medium. Assuming that the chromate ester i s an i n t e r -mediate i n t h i s reaction another problem concerning the mechanism presents i t s e l f . Does the ester de-compose i n an open form or does i t undergo a c y c l i c decomposition? While there i s at present no d e f i n -i t i v e evidence on which to make a choice between a c y c l i c and a non-cyclic decomposition i t would seem that the c y c l i c mechanism most adequately explains a l l the observed features of the reaction. While Richer, P i l a t o and E l i e l (56) have recently challenged t h i s mechanism on the basis of the rates of oxidation of several alkylated cyclohexanols, Kwart (57) has adequately r e p l i e d to t h i s attack pointing out that the observed rates can be expressed i n such a way 179 that they a c t u a l l y substantiate the c y c l i c decomposi-t i o n mechanism. If a c y c l i c decomposition mechanism i s operative the s i g n i f i c a n c e of a negative nrho t t value i s not r e a d i l y apparent since such a mechanism could f a c i l -i t a t e electron transfer either through the O-Cr bond or through the developing H-0 bond. R 0 V OH R C >Cr or R 0 K OH / N /H R H * 0 The mode of transfer would depend on the amount of o r b i t a l overlap between these atoms i n the t r a n s i t i o n state. Observance of a negative w r h o w value f o r such a mechanism would seem to indicate that the trans-ference of electrons i s mostly through the H-0 bond. Another approach to the nature of the t r a n s i t i o n state would be to consider i t as a resonance hybrid of several possible forms as Roc'ek (4) has recently done f o r the corresponding oxidation of hydrocarbons by chromic a c i d . R ,0 OH \ ! R' C r — A \ HO R 0 OH R HO R 0 OH / C . ;Cr_A R HO II III 180 In t h i s instance a negative n r h o * value would merely indicate that form I contributed more to the t r a n s i t i o n state than did form I I . This again simply means that the reaction proceeds mainly by hydride tr a n s f e r . The present study has also indicated that the det a i l e d mechanism of oxidation of alcohols by chromium VI may vary depending on the a c i d i t y of the medium used and i n some regions upon the i d e n t i t y of the mineral a c i d used as a proton source. In d i l u t e acid the following mechanism i s operative: R2CH0H + HCrO^" + H4" ^  1 R2CHOCr03H + H 20 R 2CHOCr0 3H > R 2C0 + C r 1 7 + H 20 R2CHOCr03H + H + > R 2C0 + C r 1 7 + H 30 + Such a mechanism c o r r e c t l y predicts the observed rate law, i t allows C-H bond f i s s i o n to be the rate-deter-mining step and i t indicates that the reaction rate should be independent of the nature of the proton source i n t h i s region. In more a c i d i c solutions, beyond the pKa of chromic a c i d , t h i s mechanism i s modified i n the following manner: HCrO^" + H + + HA » HOCrOgA + H 20 1 8 1 R2CH0H + HOCrOgA * RgCHOCrOgA + HgO This mechanism f i t s the observed k i n e t i c rate law, i t allows C-H bond f i s s i o n to be the rate-determining step, i t suggests that a negative n r h o n value should be observed, i t incorporates ester rather than alcohol protonation into i t s framework and i t i n -dicates that the rate of reaction should be dependent on the nature of the mineral acid medium i n which the reaction i s ca r r i e d out. F i n a l l y , i n very concentrated s u l f u r i c acid the mechanism changes i n such a way, that the deuterium isotope e f f e c t c l o s e l y approaches unity. Such an observation suggests that ester formation may become rate-determining i n t h i s medium. R \ + + ^CHOH + HOCrO(OS0 3H) 2 ^ RgCHOCrO(OS03H)2 + H 20 (slow) R + + R 2CHOCrO(OS0 3H) 2 > R2CO * HOCr(OS0 3H) 2 182 VI SUGGESTIONS FOR FURTHER RESEARCH During the course of t h i s investigation a very-b r i e f study was made of the use of t r i f l u o r o a c e t i c a c i d as a solvent f o r chromic acid oxidation reactions. This reagent has good solvent properties s i m i l a r to those of a c e t i c acid and i t has the added advantage of possessing no oxidizable carbon hydrogen bonds. By use of t h i s solvent one should then be able to eliminate any p o s s i b i l i t y of interference i n the reaction by solvent oxidation. In order f o r i t to be useful i n mechanistic studies one would f i r s t have to determine an a c i d i t y function f o r the t r i -f l u o r o a c e t i c a c i d - s u l f u r i c a c i d system s i m i l a r to the one that Mares and Roc'ek (6) have developed f o r the a c e t i c a c i d - s u l f u r i c acid system. The b r i e f use made of t r i f l u o r o a c e t i c acid in t h i s i n v e s t i g a t i o n revealed however that i t had one disadvantage - i t i s d i f f i c u l t to pu r i f y . When d i s -t i l l e d from chromic acid a v o l a t i l e red complex, apparently containing chromium VI, d i s t i l l s over at 72°C. It should be i n t e r e s t i n g to investigate further the exact nature of t h i s complex. Another p o s s i b i l i t y f o r further work would be I the use of a series of substituted sulfonic acids as proton sources. It should be possible to tes t 183 the ideas concerning mineral acid incorporation beyond the pKa of HgCrO^ by use of these acids. This should be possible since one i s able to predict with f a i r c ertainty what e f f e c t incorporation of these acids would have on the physical and chemical properties of the chromium species, i . e . One should be able to predict from p r i o r knowledge how the properties of HO-CrOgOSOgC^H^Crij should d i f f e r from those of HOCr0 2OS0 2C 6H 4N0 2. 184 VII APPENDIX - RATES NOT APPEARING IN THE MAIN BODY OF THE THESIS OXIDATION OF ISOPROPYL ALCOHOL IN BUFFERED SOLUTIONS No. [NaH 2P0j [HClOj pH Temp. [Cr I V] [Ale] k 2 (M) (M) (°C) (M) (M) 1 moles sec" 1 3 5 1 1 . 0 0 1 0 . 4 1 6 1 . 7 3 9 7 . 6 0 . 0 0 3 6 1 0 . 0 4 3 6 2.85x10" 3 5 3 1 . 0 0 1 0 . 2 0 8 2 . 3 9 7 . 6 0 . 0 0 3 4 7 0 . 0 6 5 4 1 . 1 9 5 x 1 0 3 5 4 1 . 0 0 1 0 . 1 0 4 2 . 7 0 9 7 . 6 0,00428 0 . 0 6 5 4 5.87x10° 3 5 5 1 . 0 0 1 0 , 0 4 1 6 3 . 1 5 9 7 . 6 0.00338 0 . 0 6 5 4 2 . 9 5 x 1 0 " 356 1 . 0 0 1 0 . 4 1 6 1 . 7 0 . 0 0 . 0 0 3 4 0 0 . 1 3 1 3 . 2 8 x 1 0 " 3 5 7 1 . 0 0 1 0 . 2 0 8 2 . 3 0 . 0 0 . 0 0 3 4 0 0 . 1 3 1 . 1 . 0 0 x 1 0 " 3 5 8 1 . 0 0 1 0 . 1 0 4 2 . 9 0 . 0 0 . 0 0 3 4 0 0 . 1 3 1 3 . 7 5 x 1 0 " ' 3 5 9 1 . 0 0 1 NaOAc (M) 0 . 0 4 1 6 3 . 3 0 . 0 0 . 0 0 3 4 0 0 . 1 3 1 1 . 5 1 x 1 0 " 3 6 0 1 . 0 0 1 . 0 1 * 0 . 0 0 . 0 0 3 4 0 0 . 1 3 1 5 . 3 5 x 1 0 ° * When t h i s acetate buffer was used the pH gradually increased from an i n i t i a l value of 2 . 3 to a value of 3 . 1 a f t e r one quarter of the chromium VI had reacted. 185 VIII BIBLIOGRAPHY 1. F. H* Westheimer and A. Novick, J. 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O'Donnell, Ph.D. Thesis, University of B r i t i s h Columbia (1962). 99. H. C. Mishra and M. C. R. Symons, J. Chem. S o c , 4411 (1962) 100. R. W. Taft, Steric E f f e c t s i n Organic Chemistry (Edited by M. S. Newman) Wiley, New York (1956) pp. 556 101. G. Harbottle, J. Am. Chem. S o c , 22, 4024 (1951) 102. M. A. Paul, J. Am. Chem. S o c , 26, 3236 (1954) 103. Frank Holloway, J. Am. Chem. S o c , r]hL% 224 (1952) 104. F.« H, Westheimer, Private Communication to Ross Stewart 105. W. H. Hartford and M. Darrin, Chem. Rev., £8, 1 (1958) 191 106. V. H. L. Krause and K. Stark, Z. Naturforsch., 176. 1 (1962) 107. B. J. Hathaway and A. E. Underhill, J. Chem. S o c , 3705 (I960) 108. H. C. Clark, Private communication. 109. C. F. Prutton and S. H. Maron, Fundamental P r i n c i p l e s of Physical Chemistry. The Macmillan Company, New York (1951), PP. 478. Fig.1 T y p i c a l r a t e p l o t s f o r t h e o x i d a t i o n o f a r y l t r i f l u o r o -m e t h y l c a r b i n o l s in 7 7 2 % a c e t i c a c i d . -0 5 0 T i m e (min) F ig .2 T y p i c a l r a t e p l o t s f o r t h e o x i d a t i o n of a l c o h o l in 0 . 4 6 3 M H C I 0 4 . 1 5 0 i s o p r o p y l 6 9 1 o u Cxi 2 I U K J X 1 Ci) CJL) 0 1 +1 F , 9 - 3 p K a of H 2 C r Q 4 d * in H 2 S 0 4 / / °/ 7 0 / / S l o p e = 1 . 3 8 / & H o • / H . o -1 +V 9 A c i d i t y -) F i g . 4 p K a o f H 2 C r 0 4 in H C i O 0 S lope=1 .04 H o • H . o -1 -1 A c i d i t y - 2 70 71 Fig.7 p K a O f h ^ C r C ^ in H 3 P 0 4 f - N a C D 4 • + X CD O CQ U D) 3 (5/ I ^ -1 F i g . 8 Ion iza t ion o f i n d i c a t o r s in H3PO4-NaCI0 4 s o l u t i o n s , o.p-n ' i t roani l ine. • , o - n i t r o a n i i i n e . ©, 4-chloro-2-nitroanir:ne. Fig.9 Ion i za t i on o f i n d i c a t o r s in HNC^ -NaC IO^ s o l u t i o n s , o, .p-ni t ro-an i l ine . • . o - n i t r o a n i l i n e . o ^ - c h l o r o - ^ n i t r - o a n i l i n e . S3 _ 8 _ 7 _ 6 a . S l o p e = - 8.14 _ 5 \ C H 3 C H O H C H 3 _ 4 A d C y v c a S l o p e = - 8 . 3 \ . C H 3 C H 2 C H 2 O H \ J C F 3 C H 2 O H ^ .1 .2 .3 .4 . 5 <b | | | A d I I .6 I F ig .10 C h r o m a t e e s t e r e q u i l i b r i u m C o n s t a n t s . 8 9 Fig.11 H a m m e t t p l o t f o r t h e c h r o m i c a c i d o x i d a t i o n of s u b s t i t u t e d p h e n y l t r i -f l u o r o m e t h y l c a r b i n o l s in 7 7 % a c e t i c a c i d . O p r o t i o c o m p o u n d s 0 0 . 5 1.0 90 F i g . 12 _2 .0 H a m m e t t p lot f o r t h e c h r o m i c a c i d o x i d a t i o n o f s u b s t i t u t e d p h e n y l t r i -f l u o r o m e t h y l c a r b i n o l s in 7 7 % a c e t i c a c i d . Q - M e & f i - t - B u V m ~ M e H \ \ p . e - c i J . O \ ^ m - B r 2 + l o g k \ . 6 = - Q 9 5 _0 .5 m - N C b ^ I O) 1 J 3 - d i N 0 2 C I I 1.0 1.5 2.0 •P l o g ( K H / K D ) , ]j1 ™ _ Fig.13 R e l a t i o n s h i p b e t w e e n t h e r a t e of o x i d a t i o n o f t r i f l o r o -m e t h y l c a r b i n o l s in 7 7 2 % a c e t i c ac id a n d t h e p r i m a r y d e u t e r i u m i s o t o p e e f f e c t s . 100 J . 2 3 _A _ 0 < / S l o p e = 1.12 lA 7 x 8 P 1 1 I Ho i 1.0 1.5 2.0 2.5 3.0 F i g . 14 L i n e a r r e l a t i o n s h i p b e t w e e n r a t e a n d a c i d i t y f o r t h e o x i d a t i o n o f p h e n y l -t r i f l u o r o m e t h y l c a r b i n o l in 7 7 . 2 % a c e t i c a c i d . 106 Fig.15 Taft a>*-Q* plat f o r t h e o x i d a t i o n of s u b s t i t u t e d 2 - p r o p a n o l s in 5 0 . 1 % H S Q . Fig.16 T h e r m o d y n a m i c p a r a m e t e r s . .4 .0 C I 3 C H O H C H 3 in 7 7 . 2 % A c O H 3.1 *32 1 Q Q 0 3 3 J . . . . T I 3.4 1 _3000 F ig 17 E x t i n c t i o n c o e f f i c i e n t s o f C r in su l fur ic ac i d so l u t i ons . 2 0 0 0 • — * -1 0 0 0 3 4 7 m j j 313 mj j -O -Q -4 4 4 m j j H» • ~5 -.6 H - - 7 - .8 - .9 M On Fig.18 U l t r a v i o l e t s p e c t r a of C r 1 2 1 in H g S Q j s o l u t i o n s . 1 0 0 0 I X (m j j ) t-1 00-120 be J . 4 O p t i c a l D e n s i t y F ig . 19 A t t e m p t e d d e t e r m i n a t i o n of t h e pKa f o r _p-to!uic a c i d in a c e t i c ac id . ^ / / / / 3 4 5 6 — - e x p e r i m e n t a l c u r v e — c u r v e a s s u m i n g a p K a o f 7 8 9 _ _ j e m p e r i c a l - 6 . 9 . F i g . 2 0 R e l a t i o n s h i p b e t w e e n H o a n d t h e r a t e o f o x i d a t i o n o f t h r e e 2 - p r o p a n o l s in s u l p h u r i c a c i d s o l u t i o n s . P r i m a r y d e u t e r i u m i s o t o p e e f f e c t s i n d i c a t e d in p a r e n t h e s i s . Fig.21 O x i d a t i o n o f l ^ l - t r i f l u o r o - 2 - p r o p a n o l in 5 7 6 ° / © H 2 S 0 4 . 16 2+ logO.D . 2 0 0 3 0 0 T i m e (min.) i J . 6 0( F i g . 2 2 O x i d a t i o n o f 1 ,1 ,1 - t r i f l uo ro -2 -p ropano l in 6 3 . 9 % H 2 S 0 4 . J . 4 .2+log O.D. J . 2 J . O I T i m e (mm.) i i | ^ T i m e (min.) L-2.1 F ig . 2 5 O x i d a t i o n o f 1 ,1 ,1 - t r i f l uo ro -2 -p ropano l in 8 5 . 7 % H 2 S 0 4 . L 2 . 0 J . 9 2+ l o g Q D . 1.8 U . 7 1.6 1.5 5 J 1.0 „ s _ % 1 5 - 2 D F i g . 2 6 O x i d a t i o n o f 1,1,1-trifiu o r o - 2 - p r o p a n o l 9 5 3 % H 2 S 0 4 . in J S J . 6 J . 4 5 0 i 1 0 0 T . , . J 5 0 i T i m e (mm.) ; 2 0 0 i F i g . 2 7 R e l a t i o n s h i p b e t w e e n k 0 k S a n d a>* f o r t h e o x i d a t i o n o f s u b s t i t u t e d m e t h a n o l s . -1 F i g . 2 9 C o m p a r i s o n of t h e Ho f u n c t i o n f o r o f H 3 P 0 4 a n d N a C I 0 4 (o) w i t h t h a t f o r H 3 P G 4 (•). 1:1 r a t i o LJ rio o [ H 3 P 0 4 - N a C I 0 4 ] ( M ) 3 4 i— •v ON F i g . 3 0 C o m p a r i s o n o f t h e H o f u n c t i o n f o r 1:1 r a t i o o f HNO3 a n d N a C I 0 4 (o ) w i t h t h a t f o r HNO3 (•). H N O g - N a C I O ^ 3 _L F i g . 31 U l t r a v i o l e t s p e c t r u m o f Crm in 3 . 5 0 H C I 0 4 . . 2 0 0 0 € J 0 0 0 2 5 0 i 3 ? ° X(mj j . ) 3 5 0 i 4 0 0 ^ - " ' i ., i,..... V) i F i g . 3 2 U l t r a v i o l e t s p e c t r u m o f Crm in 3 .10M H 2 S 0 4 . F i g 3 3 U l t r a v i o l e t s p e c t r u m o f C r 5 3 in 3 . 6 0 M HCI. F i g . 3 4 U l t r a v i o l e t s p e c t r u m o f C r * 1 in 3 . 0 0 M h k P Q L2000 165 I a u I U) o I CL) I 0) F ig . 3 5 " p K a " o f H 2 C r S O ? . _ 0 . 6 _0 .0 / p S l o p e = 0 . 9 6 - 0 J 6 X 1 +1 I H.+ l o g a H 2 Q / a H A F ig . 3 6 " p K a " o f H C r 0 3 C I . • _0 .6 _0 .0 O / ' S l o p e = 0 . 9 9 z .06 +2 I +1 I H-+ log a H 2 o / a H A 168 F i g . 3 7 C h r o m i c a c i d o x i d a t i o n o f i s o -p r o p y l a l c o h o l in a q u e o u s s o l u t i o n s o f m i n e r a l a c i d s . 

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