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Kinetics of the cis-trans isomerization of azobenzene Ciccone, Stefania 1959

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KINETICS OP THE CIS-TRANS ISOMERIZATION OF AZOBENZENE  by  STEFANIA CICCONE  A THESIS SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE i n the Department of CHEMISTRY at the UNIVERSITY OF BRITISH COLUMBIA  We accept this thesis as conforming to the required standard  Members of the Department of Chemistry THE UNIVERSITY OF BRITISH COLUMBIA September, 1959  ii  ABSTRACT  The catalytic effects of several acids and metal salts on the cis-trans isomerization of azobenzene in aqueous ethanol were examined kinetically. The effect of perchloric acid i s apparently due to H ions: a catalytic +  mechanism involving the formation of the conjugate acid of azobenzene has been postulated to interpret these results* To account for the much higher catalytic activity found for hydrochloric acid, an additional path, involving catalysis by undissociated HC1 molecules has been proposed. Acetic acid was found to be inactive. Of the metal salts examined only those of Cu shoved pronounced catalytic activity, which i s interpreted in terms of a catalytic mechanism involving coordination of cupric ions with the azo group* ++ Simultaneous., coordination of Cu  + and H has been proposed to account for the  high catalytic activity of cupric salts in the presence of acids.  In presenting  t h i s thesis i n p a r t i a l fulfilment of  the requirements for an advanced degree at the  University  of B r i t i s h Columbia, I agree that the Library s h a l l make i t f r e e l y available for reference  and study.  I further  agree that permission for extensive copying of t h i s thesis for scholarly purposes may  be granted by the Head of my  Department or by his representatives.  It i s understood  that copying or publication of this thesis for f i n a n c i a l gain s h a l l not be allowed without my written permission.  Department The University of B r i t i s h Columbia, Vancouver Canada.  ACKNOTOEDGEMENT  The author is grateful to Dr. J. Halpern for his support and inspiring direction of the research reported in this thesis and for his constructive criticism of the manuscript during i t s preparation. Thanks are also due to Dr. Boss Stewart for his kind collaboration* The author is also indebted to the National Research Council of Canada for financial assistance.  TABLE OP CONTENTS Page INTRODUCTION  1  Cis-trans isomerization of azobenzene.. . . . . . Solvent effects  .....  ••  1  . . . . . . . . .  2  Substituent effects . . . . . . . . . . . . . . . . . . . . . . . .  2  Catalytic effects  3  . . . . . . . . . . . . . . . . .  General consideration of acid catalysis  4  Structure of the conjugate acid of azobenzene  14  Scope of present investigation  15  EXPERIMENTAL  16  Materials  16  Other reagents  16  Spectrophotometry and kinetic measurements . .  16  RESULTS AND DISCUSSION  21  Effect of perchlorid acid  21  Effect of hydrochloric acid  27  Effect of acetic acid  38  Effect of cupric perchl orate  39  Effect of cupric chloride  50  Effect of cupric acetate  50  Results i n benzene  55  Miscellaneous effects SUMMARY AND CONCLUSIONS  . . . . .  55 58"  SUGGESTIONS FOR FURTHER WORK  60  REFERENCES . . . . .  62  LIST OF TABLES  Table No.  v  P a  ge  I.  Active and inactive substances (7)  5  II.  Effect of perchloric acid at 60° C  24  Effect of Hydrochloric acid at 60° C  30  Effect of acetic acid at 60° C  33  Effect of cupric perchlorate  41  VI.  Effect of cupric chloride  53  VII.  Effect of cupric acetate  54  Miscellaneous results  56  III. IV. V.  VIII.  LIST OF FIGURES  Figure No. 1.  h 0  2.  Page  f o r aqueous solutions of HC1 and HCU)  (8) . . . . . . . . 4 i n mixed solvents containing 0.1 M HC1. (13) • • • • • • • •  10  H 12 o 3. Dependence on solvent composition of rate of reactions following an Ho- dependent mechanism (12, 13) (Rearrangement of Ethynylpropenyl-Carbinol i n solution of 1 M HCl) • • • • • • • • 13 4.  Spectrum of trans-azobenzene i n 50 mole % ethanol i n the presence of HC10 , HCl and AgClO^ . . 18 4  5.  Spectrum of cis-azobenzene i n 50 mole % ethanol i n the presence of HC10„, HCl, and AgCIO 4  6.  7.  19  4  Spectrum of cis-azobenzene i n 50 mole i> ethanol i n the presence of Cu(C10 Vand Co(C10 ) . 20 4 2 4 2 F i r s t order kinetic plots f o r uncatalyzed and HClO^catalyzed cis—trans isomerization of azobenzene i n 50 mole JfrTStOH at 60 C. 22  8.  Catalytic effect of perchloric acid at 60°C.  9.  Solvent dependence of the catalytic effects of perchloric and hydrochloric acids at 60°C. The broken line depicts the solvent dependence of the uncatalyzed rate (k ) • • • • • • • • • • • • 28 o F i r s t order kinetic plots for HCl catalysed cis-trans isomerization of azobenzene i n 50 mole % EtOH at 60 C« . . . . . . • 29  10. 11.  .25  Catalytic effect of hydrochloric acid at 60° C. . . . . . . . •  12. Concentration - dependence of HCl catalysis plotted according to equation (2l)  34 37  13. F i r s t order kinetic plots f o r Cu(C10 ) catalyzed cis-trans isomerization of azobenzene i n 50 mole 7^ ethanol at 60 C, , .  40  14. Catalytic effect of cupric perchlorate i n 50 mole i» EtOH at 60° C. . .  43  15. plot of log k  c a t  vs. log [Cu(G10 ) 4  2  . ,. 44  16. F i r s t order kinetic plots f o r the cis-trans isomerization of azobenzene catalyzed by Cu(Cl0 L i n the presence of 0.1 M HC10 i n 50 mole # EtOH at 60° C. . , . . . . . . . < , p 46 17. F i r s t order kinetic plots for the cis-trans isomerization of azobenzene catalyzed by Cu(C10.) i n the presence of 0.3 M HC10, i n 50 mole % EtOH a t 60*Ct 4  „  47  18. Combined catalytic effect of cupric perchlorate and perchloric acid i n 50 mole % EtOH at 60° C  48  19. Combined catalytic effect of cupric perchlorate and perchloric acid i n 50 mole fo EtOH at 60 °C  49  F i g u r e No.  (LIST OF FIGURES (Cont'd)  Page No.  20* F i r s t order k i n e t i c p l o t s f o r the c i s - t r a n s i s o m e r i z a t i o n of azobenzene c a t a l y z e d by CuCl- i n the presence o f 0.01 M HC1 i n 50 mole % EtOH a t 60°C.  51  21. F i r s t order k i n e t i c p l o t s f o r the c i s - t r a n s i s o m e r i z a t i o n of azobenzene c a t a l y z e d by CuGl i n the presence of 0.02 M HC1 i n 50 mole # EtOH a t 60°C. * . . .  52  22. Combined c a t a l y t i c e f f e c t o f c u p r i c c h l o r i d e and h y d r o c h l o r i c a c i d i n 50 mole % EtOH a t 60 C. . . . . . . . . . . . . . . .  54 A  INTRODUCTION Cis-Trans Isomerization of Azobenzene Cis-trans isomerism, doe to restricted rotation about a double bond, was f i r s t observed for ethylenic compounds and while i t had long been predicted also for molecules possessing -N=N- bonds, this has been confirmed only relatively recently. In 1937 Hartley (1,2) prepared cis-azobenzene by irradiation of a solution of the normal trans-form with ultraviolet light.  The cis-isomer,  although stable in crystalline form, undergoes spontaneous conversion into the trans-form in solution*  This occurs by both thermal and photochemical  mechanisms. Hartley (2) studied the kinetics of the thermal conversion spectrophotometrically and found i t to be of f i r s t order, with an activation energy of about 23 Kcal. per mole and a normal frequency factor (  10  11  l i t e r mole * sec* *").  The kinetics of the cis-trans isomerization of ethylenic compounds have interpreted (3) in terms of two types of mechanisms. The f i r s t , characteristric of reactions having high activation energies (40 - 50 Kcal) and normal frequency factors (10*^ - 10^ liter mole * sec  *) i s  adiabatic and involves rotation about the double bond without any change in the multiplicity of the electronic state of the molecule. The second, characterized by lower activation energies (15 - 25 Kcal) and abnormally 3 5 low frequency factors (10 - 10 l i t e r mole  — 1 - 1 sec ) i s apparently non-  adiabatic and proceeds through the triplet state of the molecule in which the barrier to rotation about the double bond i s removed. Although the activation energy of the isomerization of azobenzene i s low compared to  thai of most olefins, the normal frequency factor of the reaction suggests that its mechanism is of the former type.  SOLVENT EFFECTS The effect of solvent variation on the kinetics of the cis-trans isomerization of azobenzene has been examined by Hartley ( 2 ) and more extensively by Halpem, Brady and Winkler ( 4 ). Sates were determined in a large number of solvents and were found to depend inversely on the solvent polarity. Thus at 60° C, k, the f i r s t order rate-constant, increased from 0.000489 to 0.001415 min"  1  in going from a polar solvent such as  methanol to a non-polar one such as heptane. The activation energy correspondingly decreased from 24.8 to 22.8 Kcal/mole, but the effect of this was partially offset by a compensating decrease in the entropy of activation. These effects were correlated quantitatively with the internal pressure of the solvent in accordance with equations derived assuming regular solution theory behaviour. The behaviour in mixed solvents was more complex and showed evidence of preferential solvation of azobenzene by the better solvent.  SUBSTITUENT EFFECTS 4 The cis-isomers of a number of substituted azobenzenes and related compounds have been prepared and the kinetics of their isomerization examined. LeFevre and Northcott ( 5 ) found the rate of substituted azobenzenes to follow the order: CH^O > CH > C l ^ Br > H > N0 3  2  but were unable to correlate this with any property of the substituents.  - 3 T h i s order p a r a l l e l e d t h a t found f o r the i s o m e r i z a t i o n o f s u b s t i t u e d diazocyanides.  The e f f e c t of solvent v a r i a t i o n was  qualitatively similar  t o t h a t f o r azobenzene i t s e l f . A more extensive  study of the i s o m e r i z a t i o n of s u b s t i t u t e d azobenzenes  has r e c e n t l y been described by T a l a t i ( 6 ).  Fara-substituents  (both e l e c t r o n  donating and withdrawing) were found t o increase the r a t e of conversion again the r e s u l t s f a i l e d to c o r r e l a t e i n any simple way p r o p e r t i e s (e.g. 6"  values).  with substituent  Very high i s o m e r i z a t i o n r a t e s were observed  f o r the para-cyano and n i t r o compounds and, f o r these compounds the crate was a t i o n of t h i s i t was  but  i n a d d i t i o n , i t was  higher i n more p o l a r s o l v e n t s .  found t h a t In explan-  suggested t h a t while other s u b s t i t u e n t s operated only  through an i n d u c t i v e e f f e c t , f o r these compounds c o n t r i b u t i o n of resonance s t r u c t u r e s such as  was  (I)  p o s s i b l e , i n which the IT - e l e c t r o n d e n s i t y a t the azo-linkage  decreased, r e s u l t i n g i n reduced r e s i s t e n c e t o r o t a t i o n .  is  Furthermore,  because of the charge separation i n v o l v e d , such s t r u c t u r e s should be favoured by i n c r e a s i n g solvent p o l a r i t y .  CATALYTIC EFFECTS A c i d c a t a l y s i s of the c i s - t r a n s i z o m e r i z a t i o n of azobenzene was observed by H a r t l e y (2) who  first  d i d not however examine the e f f e c t i n d e t a i l .  - 4While this investigation was in progress, Schulte-Frohlinde (7) reported examining a large number of other substances for catalytic activity; Table (i), listing the inactive and active substances, is reproduced from his paper. The catalytically active substances are classified into two categories, electron donors and electron acceptors, acids being included among the latter.  _3 Bases and week acids (K < 10 ) were found to be inactive. The catalytic effect of thiourea was found to be greatly enhanced by acids, the combined activity being greater than the sum of the separate activities. Benzene appears to have been used as the solvent in these studies and many of the inorganic substances tested were presumably insoluble. It is reasonable to suggest that in some instances failure to observe catalytic activity may have been due to this factor. Schulte-Frohlinde has interpreted the pattern of catalytic effects in terms of the electron-transfer properties of the catalytic species and suggested that the catalytic path involves a transition state in which one or two electrons are transferred between the catalyst and azobenzene. The evidence for this concept does not appear to be convincing and there are notable exceptions to the suggested criterion for catalytic activity. It should be emphasized that the studies referred to here are for the most part qualitative and this do not lend themselves to detailed interpertations. GENERAL CONSIDERATION OF ACID CATALYSIS la considering the acid catalysis of the cis-trans isomerization of azobenzene, i t is of interest to review some concepts relating to acid catalysis in general.  TABLE (I)* LIST OF INACTIVE AND ACTIVE SUBSTANCES INACTIVE SUBSTANCES 1. LiOH  5. CaO  2. NaOH  6, BaO  3. NaCl  7. CaCl  4. Na S0 2  4  9. Cu-powder  13. Mu0  2  KMnO.  10, CufOOCCR^ 2  8. SiQ„ 2  n  -  4  PbO  HgO  12. MuSO, 4  16. (NH,) S 0 4 22 8  17. Ethanol 18. Phenol 19.  o—Dihydroxybenzene  20. p-Dihydroxybenzene 21. p-flydroxyazobenzene 22. Benzoic acid 23. o-Chlorobenzoic acid 24. p-Chlorobenzoic acid 25. Trans—cinnamic acid 26. o-Phthalic acid 27. Isophthalic acid 28.. p-Araino-benzoic acid 29. Pyridine 30. Piperidine 31. Diphenylamine 32. Carbazole 33. Urea 34. Phthalimide  •Quoted from Von Dietrich Schulte-frohlinde (7)  35 e  Diphenylbenzamidine  36.  syn-and anti-pyridine 2-aldehyd-phenyIhydrazone  37*  1,3 Dinitro-benzene  38.  2,4 Dinitro-phenol  39. m*Nitroaniline 40.. p~Nitroaniline 41.  3 Methoxy - 5 - ethylbenzoquinone  42. 4 ^ - j - j - Tetramethylbutylbenzoquinone (1,2) 43. p-Quinone 44. Metoxy-p-quinone 45. 46.  Phenanthrene-quinone 1,4- Dihydroxy-antbxaquinone  47. N» Phenyl-N  t  - benzoyl-hydrazine  ACTIVE SUBSTANCES a). Electron-donors  1.  Sodium  2.  Zinc dust  3.  SolidyU. -Sulphur in contrast to i^-Sulphur  4.  Colloidal sulphur  5.  Hydrogen sulphide  6.  Thiophenol  7*  Hydrazobenzene  8.  Phenylhydrazine  9.  Benzaldehyde-phenyl-hydr azone  A 7*  10*  Hydroxylamine-*ydrochloride  11.  Semicarbozide-bydrochloride  12. Ferrous-sulphate 13. Arsenious-acid  b).  ELECTRON ACCEPTORS  1.  Bromine  7. Lead oxide  2.  Iodine  8. Ferric sulphate  3.  Tetrachloro—p-quinone  9* Hydrochloric acid  4.  Tetrabromo-p-quinone  5. 6.  Tetrahydroxy-p-quinone Potassiunwiitroso - bisulphonate 13.  10.  Oxalic acid  11. Salicylic acid 12. Picric acid  2,4,6, Trinitro—benzoic acid  These have been found to have validity in relation to a great variety of reactions (8) and i t seems likely that they apply also to the present system* While most acid-catalyzed reactions are believed to proceed through the conjugate acid of the reactant a distinction is made between several different types of mechanisms. The f i r s t , called "unimolecular" postulates as rate determining step the decomposition of the protonated species found in a rapid pre-equili*brium, to give the products* A +H AH  +  +  ^  AH  —>  products  +  (fast)  (l)  (slow)  (2)  such a mechanism we will designate as (I). Mechanism (II) known as the "bimolecular" mechanism, also involves a rapid protonation pre—equilibrium, but in this case the rate-determining step i s the attack of a molecule of solvent (S) on the protonated species to give further intermediate AHS which decomposes to yield the products. +  A +H  ;=± AH*  +  —>  AH + S AHS  (fast)  AHS  +  —> products  +  (slow)  (3)  (fast)  (4)  The third type of mechanism involves the protonation itself as rate determining step A+H  —*  +  AH*  AH*  (slow)  (5)  products  (fast)  (6)  One of the most widely used criteria for distinguishing between these three types of mechanisms involves the Hammett acidity function, H . This Q  function was introduced by Hammett and Deyrup (8,9) and i s defined in terms of the ionization equilibria of a class of indicators, namely those functioning as uncharged bases, B+H  +  =±  BH  +  (7)  by the equation: H  =  - log BH+ C  + pK™+  (8)  % C  +  where  BH  i s the ratio of the concentration of the indicator in i t s acid  °B and basic form, directly measurable for a given solution by means of a spectrophotometer or a colorimeter, and  *- the thermodynamic ionization s  constant of the conjugate acid BH « This definition i s equivalent to H H  log  V"  *B  f^T" n  (9)  i s the hydrogen ion activity and fg and fgjj+ are molar concentra-  f  tion activity coefficients* It has also been found useful to define the function h  which i s related o  to H  Q  by the expression H = - log h o o  (10)  In the limiting case of ideal dilute solutions: H and  0  = pH  h = [H*] Q  A characteristic feature of the H  q  acidity scale for aqueous solutions  of the strong acids is that with increasing acid concentration h  Q  increases  more rapidly than [H ], Fig. 1 shows the deviation from linearity in the +  plot of h  versus [H l for HC10, and HC1 in aqueous solution (8)* Below 4 +  o  IM, the h values for HCIO^ and HC1 are nearly identical but the divergence Q  from linearity i s already appreciable. Zucker and Hammett (10,8) have suggested that i f the acid catalyzed reaction proceed through a mechanism of type (I) a linear dependence of the rate on h  Q  i s to be expected. On the other hand i f the mechanism is of type  (II) the rate should increase linearly with [H ]. +  - 10 -  11 -  In the third case (slow protonation step) a linear dependence of the rate on h o  i s again suggested; this case should also be distinguishable from the ^  other two by virtue of exhibiting general acid (rather than specifically hydrogen ion) catalysis* It i s necessary to point out that these criteria are by no means unambiguous* nor i s the theoretical foundation on which they rest of unquestioned validity (8)* Furthermore although empirical correlations of rates with H may be observed in acid catalyzed reactions of charged o molecules and in reactions of higher order, the significance of these (i*e* when applied to reactions other than those in which the i n i t i a l step involves a single uncharged molecule and an H able*  +  ion) is very question-  Furthermore since the variation of activity coefficients with change  in medium i s not well understood, caution must be observed in the interpretation of studies made in non-aqueous or mixed solvents* Although the H  q  concept i s applicable in principle to non-aqueous  solutions, including mixed solvents, relatively l i t t l e work has been done on such systems* Some measurements by Braude and Stern ( l l , 12, 13) on solution of HCl in aqueous ethanol, acetone and dioxane are reproduced in Fig* 2.  A common feature of these systems is that for a given concentration  of HCl, H  Q  passes through a minimum as the solvent composition i s varied  from pure water to the pure organic liquid: occurs in the vicinity of 50 mole i> H^O*  in each case this minimum Braude and Stern (12, 13) found  a similar dependence on changes of solvent composition for the rates of reactions known to follow an H - dependent mechanism* This is shown in .o Fig* 3 which depicts their results for the rearrangement of the ethynylpropenylcarbinol in solutions of 1M HCl in aqueous ethanol* some indication of the validity of the H  0  This provides  concept in these media*  0.0  1 0  50  100  M O L E io E T H A N O L FIGURE  3.  Dependence  f o l l o w i n g an H  q  on solvent composition of rate of r e a c t i o n s  - dependent mechanism  Ettaynylpropenyl-Carbinol  (12,13) (Rearrangement of  i n s o l u t i o n of 1 M HCl)  Structure of the Conjugate Aoid of Azobenzene The results of recent studies of the conjugate acids of azobenzene and its derivatives (14, 15) are consistent with the suggestion that acid catalysis of the cis-trans isomerization involves the formation of a protonated species* The protonation equilibrium of azobenzene has been studied spectro— photmetrically in aqueous sulphuric acid solutions containing approximately  20}t> ethanol  by volume* at the constant temperature of 25°C« Due to the  rapid isomerization only the trans-form (probably in equilibrium with small amount of cis-form) could be studied. A pK of - 1*6 for the conjugate ft  acid of azobenzene was calculated taking Hammetts*  scale as reference*  Klotz (16) had previously reported —2*48 for the pK of azobenzene, cala culated with reference to the acidity scale of Michaelis and Gxanick (17). The structure of AH has not been conclusively established* +  Klotz  and co-workers (16) assume structure (l) with the proton localized on one of the nitrogens through a bond involving lone pair electrons*  (I) Jaffa (14) on the other hand favors structure (II) in which the hydrogen ion i s equally shared by the two nitrogen atoms through delocalized bonds.  (II)  -  15  -  This suggestion i s based i n part on measurements on un symmetrically substituted azobenzenes which yielded a single Hammett plot i n spite of the presence i n the molecule of two non-equivalent N atoms*  The proposed  structure (II) i n which the H atom i s associated with both N atoms i s more readily reconciled with this result* than i s structure (I)* Jaffe  (14)  had also concluded on theoretical gounds, that the cis—conformation of structure (II) i s the more stable one* SCOPE OP PRESENT INVESTIGATION At the time this investigation was undertaken only the acid catalysis of the cis-trans isomerization of azobenzene had been reported and the i n i t i a l objective of the investigation was to examine the kinetics and mechanism of the acid-catalyzed reaction* Three acids were selected for study, HCIO^, HCl, CH^POOH* considered as examples of strong, intermediate and weak acids*  Furthermore, since i t  appeared l i k e l y that the acid catalysis involved protonation of nitrogen atoms or of the — N=N — bond* i t was also considered of interest to ex— amine the catalytic affects of metal ions such as Cu  * Ag  and Co  , whose  tendency to coordinate with nitrogen-containing ligands i s well known* Most of the measurements were made using aqueous alcohol as solvent, since both azobenzene and the catalysts were conveniently soluble i n t h i s medium*  A few measurements using organic acids and salts as catalysts were  made i n benzene solution*  - 16 •  EXPERIMENTAL  Materials Cis-azobenzene was prepared by the method of Hartley ( l ) .  100 grams of  commercial azobenzene* consisting almost entirely of the trans-isomer (M»pt* 67°C) was dissolved i n 1000 ml of Glacial Acetic Acid containing 20 ml of water, and the solution was irradiated with a GE (275 W) sunlamp for 24 hrs*  An appreciable  change i n color, due to increasing concentration of the cis-isomer, which i s darker, was observed*  900 ml of water was then added to precipitate the trans-isomer*  The latter was separated by f i l t r a t i o n and recycled ( i . e . subjected to further exposure i n acetic solution as described above*) The f i l t r a t e was further diluted with a l i t e r of water and the cis-isomer extracted with chloroform*  The  chloroform extract was evaporated i n a vacuum desiccator and the crude cis-isomer o recrystallized from petroleum ether u n t i l a constant melting point of 70 C was achieved*  The y i e l d of pure c i s after two exposures was about 8 grams* A l l  operations involved i n the separation and purification of cis-azobenzene were performed i n the dark since the two isomers are photochemically interconverted* Other Reagents D i s t i l l e d water and ethanol, r e d i s t i l l e d from KOH,  were used as solvents*  Cupric perchlorate (G.F.Smith reagent) was recrystallized from perchloric acid* Acids and other reagents were of reagent grade* Spectrophotometric  and Kinetic Measurements  Trans-azobenzene shows a strong absorption band i n the region 430-440 m (  max = 435 m  ) which has been ascribed (18,19) to the - N = N - linkage*  The  cis-isomer absorbs i n the same region (the band i s shifted s l i g h t l y toward the lower wave lengths) but with an extinction coefficient at 430 m of trans*  Thus the decrease i n optical density  about twice that  - 17 -  at 430  can be used to follow the isomerization reaction and this pro-  cedure, used in most of the earlier studies of the cis-trans isomerization of azobenzene, was also adopted in this investigation. Measurements were made with a Beckmann DU Spectrophotometer* In the kinetic experiments, solutions containing known concentrations —4 of azobenzene (generally 0*08 gm/l.= 4*39x10  moles/liter) were immersed  in a constant-temperature bath maintained at 60  C*  Samples were removed  at measured time intervals and transferred directly to the spectrophotometer cells* cooled rapidly by immersion in an ice—salt mixture (to quench the reaction) and the optical density determined* Since the absorption of both cis and trans isomers followed Beer's law, the concentration of the cis-isomer present at any time, could be computed by the formula [cis-A] =  (D— D«g)  [cis-A]  Q  (ll)  where: D, i s optical density at time, t Hj> , i s optical density at infinite time (i.e. of trans-azobenzene*) D , i s i n i t i a l optical density (i.e. of cis-azobenzene) and [cis-A] o is the i n i t i a l concentration of cis-azobenzene. The validity of this procedure was unaffected by the catalysts. Thus HC10 . HCl and AgClO^ in the concentrations used did not appreciably affect 4 4 the optical densities of solutions of either cis or trans-azobenzene (Fig. 4, 5.)  Cupric and cobaltic salts did absorb appreciably at the wave length  used but their optical densities and that of azobenzene apparently were additive (Fig. 6) so that effect was cancelled out by the subtraction of the optical density terms in equation ( l l ) .  - 18 -  O  TRANS - AZOBENZENE  A  TRANS - AZOBENZENE + O.IM HCIO or O.IM HC1 4 TRANS - AZOBENZENE + O.IM AgCIO 4  •  0.6 _  <  O M  s o  0.4 _  0.2 -  I  450  500  WAVELENGTH  400 m/A  FIGURE 4. Spectrum of trans-azobenzene in 50 mole % ethanol in the presence of HC10 , HC1 and AgClO^. 4  500  450  400  WAVE LENGTH mj* FIGURE 5. Spectrum of cis-azobenzene in 50 mole i° ethanol in the presence of HC10,, HC1 and AgCIO 4  - 20 -  1.0 0  CIS-AZOBENZENE  A  CIS-AZOBENZELNE +• 0,2M Co(C10 )  •  CIS-AZOBENZENE + 0.2M Cu(C10j 42  4  2  0.8  0.6  CO  i-3  < O >-l  E.  0.4  o  0.2  450  500  WAVE LENGTH  400 mjx  FIGURE 6. Spectrum of cis-azobenzene in 50 mole ethanol in the presence of Cu(C10 ) and Co(C10 ) 4  2  4  2  RESULTS AND DISCUSSIONS Of the three acids examined. HC10. and HC1, selected as typical 4  strong and intermediate acids, showed pronounced catalytic activity while the weak acid CfijCOOH was inactive. The effects of Cu(C10 ) , Co(C10 ) , AgC10 , CuCl , Cu(CH COO) were 4  2  4  2  4  2  3  2  also investigated; of these only the cupric salts showed pronounced catalytic activity* which was greatly enhanced in the presence of acids. These effects are considered separately below. In every case the kinetics were f i r s t order in cis-azobenzene, and conformed the rate law* - d [cis-A] /dt = :k [cis-^]  (12)  where the pseudo-first order rate constant k contained contributions from both the uncatalyzed (k ) and catalyzed 0 * . ^ ) Path, i.e. fl  k  "o k  +  k  cat.  (13)  k was measured separately and k ^ determined by difference Q  ca  (a) EFFECT OF PERCHLORIC ACID Three series of experiments, i n each of which the HC10 concentration 4  was varied, were performed using different solvent mixtures; 25, 50 and 75 mole percent ethanol. The temperature was kept at 60.1° C and the i n i t i a l concentration of cis-azobenzene was 0.08 gr. per l i t e r . The mnx-iimmi HC10 concentrations which could be studied were 1.1 M 4  in the 50$ solvent mixture, and 0.5 M in the other two solvents. Below these concentrations the kinetics were of f i r s t order in cis-azobenzene as shown by the kinetic plots in Fig. 7. At higher acid concentrations the f i r s t order rate plots deviated significantly from linearity, possibly due to side reactions of azobenzene with the acid.  - 22 -  5.8  30  60  90  120  150  180  TIME - MINUTES FIGURE 7. F i r s t order k i n e t i c p l o t s f o r uncatalyzed and HC10 catalyzed c i s - t r a n s i s o m e r i z a t i o n of azobenzene i n 50 mole % EtOH a t 6 0 C. 4  0  210  -  23  ~  ; The results for these series of experiments are summarized in Table (II) and Fig. 8*  The catalytic rate appears to be f i r s t order in per-  chloric acid corresponding to the rate law.  ethanol respectively. The solvent effect i s not large but the catalytic rate in 50 mole # ethanol, i s definitely smaller than at the other concentrations.  (Fig. 9)  The perchlorate ion and ionic strength effects do not appear to be important as indicated by the observation that 0.5M Na&O., either alone 4  or in the presence of 0.5M HC10., was without effect (Table VIII). 4  It seems likely that perchloric acid is completely dissociated i n these solutions and hence that the observed catalytic effect is due to H  +  (i.e. EjO or EtOH*) ions. The general considerations concerning the +  kinetics and mechanisms of acid catalyzed reactions, outlined in the introduction, would appear to be applicable to the present system which involves the reaction of a single uncharged molecule. Of the three general types of mechanisms considered, the third, involving the protonation of the reacting species as rate determining step, seems to be most unlikely in this case as the protonation of nitrogen atoms i s usually quite fast. The other two mechanisms however, appear plausible. The f i r s t assumes that the catalyzed isomerization involves the formation of the conjugate acid of azobenzene followed by a slow unimolecular step. i.e. cis-A + H cis-nAB*  trans-AH  cis-AH ===  +  traas~AH  (fast) +  ===== trans-A + H  +  (15)  (slow)  (16)  (fast)  (17)  - 24 TABLE II EFFECTS OF PERCHLORIC ACID AT 60°  Solvent  Acid moles/liter  75 raole# EtOH  0  k x 10  -1  sec  5  k  . x 10 cat  sec""''  7.04  0  »  oa  7.51  0.47  n  0.2  8.32  1.28  tt  0.3  8.96  1.92  II  0.4  9.60  2.56  0  5.76  0  tt  0.1  6.40  0.64  ft  0.3  7.68  1.92  II  0.5  8.62  2.86  ft  0.7  9.60  3.84  A  0.9  10.6  4.84  0.1  11.2  5.44  50 mole# EtOH  II  25 mole# EtOH  0  5.11  tt  0.1  5.72  0.61  tt  0.2  6.70  1.59  »t  0.3  7.39  2.28  II  0.4  7.96  2.86  tt  0.5  8.81  3.70  0  5  - 26 -  As for the structure of the protonated complex, referring to the discussion reported i n the introduction, model (I)  (I) would be more consistent with the enhanced ease of isomerization, since the localization of the proton on one of the nitrogen atoms should result i n polarization of TT electrons and a decrease i n degree of double bond character of the azo-linkage.  Model (II)  (II) on the contrary does not explain readily the enhanced isomerization, particularly i f , as suggested by Jaffe (14), the cis-conformation of this structure i s the more stable one* An alternative description of the slow isomerization step, which we are inclined to favour, involves the formation of the intermediate A H  +  ^ S - ^ ^ , N - N S  +  ?  (III)  ^  H  (III) i n which free rotation about the double bond i s possible*  Such a mechanism  i s consistent with either structure (I) or ( l l ) for the conjugate  acid*  The simple c r i t e r i o n suggested by Zucker and Hammett (10,8) for distinguishing between these two types of mechanisms, i*e* the expectation of a linear dependence of k , on h i n the f i r s t case and on [H 1 i n the cat o second, i s not readily applicable here i n view of the limited acidity range  accessible to kinetic measurement and of the absence of h data in mixed Q  solvent.  In aqueous solutions a plot of h versus [H*] diverges appreciQ  ably from linearity even at concentrations below 1 M HCIO^ (Fig l ) Such a divergence was not observed in our case. (Fig 8). On the other hand the dependence of k h  Q  . cn. solvent composition (Fig. 9) appears to follow cat  which, at least for HC1, has been shown to pass through a minimum in  the vicinity of equimolar concentration of B^O and EtOH. (Fig. 3). Because of the conflicting nature of these indications, and the questionable theoretical validity of the Zucker-Hammett hypothesis, i t i s not possible to distinguish conclusively between the alternative mechanistic possibilities. (b) EFFECT OF HYDROCHLORIC ACID Essentially the same procedure was followed as for HCIO^, but since the kinetics remained constantly f i r s t order in cis-azobenzene the measurements were extended up to about 1.5M HC1, above which the rates became too high for convenient measurement. Typical rate plots are shown in Fig.  10.  Four series of experiments were performed using 25, 50, 75 and 95 mole percent ethanol as solvent. The results axe recorded in Table III and attention i s directed to the following features: 1.  Except at very low concentrations, the catalytic activity of HC1  is higher than that for corresponding concentrations of HCIO^ in each solvent. 2.  i:). The kinetic dependence of k  . on the total HC1 concentration is cat apparently greater than f i r s t order, (Fig. 11).  - 28 -  k  . (0.5M HC10 )  cat  4  40 k  cat '  0 , 5 M  H C 1  )  30  in O  fe  20  4» S3  o  10  1  25  50  1 75  100  MOLE % EtOH FIGURE 9. Solvent dependence of the catalytic effects of perchloric and hydrochloric acids o  at 60 C. The broken line depicts the solvent dependence of the uncatalyzed rate (k ) o  - 30 -  TABLE III EFFECT OF HIDROCHLOEIC ACID AT 60° C.  Solvent  HCl. moles/liter  k x 10 -1 sec  5  k x 10 -1 see  !  +  95 mole % EtOH  0  6.70  0  w  0.1  10.0  w  0.3  22.6  15.9  0.5  42.9  36.2  0.7  76.1  69.4  *  *  0.85  104  *  1.0  139  3.30  97.3 132  75 mole # EtOH  0  6.75  0  »  0.1  8.05  1.30  "  0.3  13.1  »  0.5  23.1  16.3  "  0.7  35.7;  28.9  »  0.8  44.7  38.0  M  0.9  56.9  50.1  "  1.0  67.2  60.4  "  1.1  88.4  81.6  *  1.2  105  6.35  98.3  * 31  -  TABLE I I I (cont.)  50 mole % EtOH  0  5.76  0  o  0.2  6.40  0.64  n  0.3  7.78  2.02  tt  0.5  10.8  5.09  n  0.7  15.0  9.24  n  0.8  17.8  12.0  it  0.9  21.4  15.6  tt  1.0  24.2  18.4  tt  1.1  27.5  21.7  33.1  27.3  ti tt  1.3  40.8  35.0  it  1.5  54.7  48.9  it  1.7  72.8  67.0  - 32 TABLE III (cont'd) 25 mole % 0  4.78  0  tt  0.1  5.56  .78  n  0.3  6.39  1.61  »  0.5  7.78  3.00  it  0.7  10.0  5.22  tt  0.8  11.4  6.62  n  0.9  12.8  8.02  n  1.0  14.2  9.42  »  1  15.8  11,0  1.3  23.2  18.4  EtOH  II  n  1  - 33 TABLE IV EFFECT OF ACETIC ACID AT 60°C,  Solvent  HAlB  k x IO  moles/liter  sec  ^  3  k  .x 10  cat  sec""*  50 mole?S EtOH  0  5.76  0  0.1  5.50  0.26  •  0.5  5.50  0.26  •  0.7  5.50  0.26  •  0.9  5.50  0.26  n  -  3*  35  -  The rate shows an inverse dependence on the H^O content of the  solvent* that i s on i t s polarity* (Fig* 11) 4*  The rates approach (and in some cases actually f a l l below) those  for HCIO^ (i) at low HCl concentration and ( i i ) in solvents of high water content (Fig* 9, l l ) * These observations suggest that in addition to the H  +  ion catalysis  observed with HClO^, there i s an additional catalytic effect which may be due to undissociated HCl molecules*  As might be expected this becomes un-  important at low HCl concentrations and in solvents of high water content* since HCl tends to become completely dissociated under these conditions* The catalytic activity then approached that of the stronger acid HCIO^. The results of an experiment involving addition of NaCl are also of interest in this connection*  Due to the low solubility of the salt in  aqueous alcohol the maximum concentration of NaCl that could be added was about 0*2 M.  While this concentration of NaCl alone was without effect* i n  the presence of HCl i t produced a slight enhancement of the catalytic activity which may be attributed to an increase in concentration of undissociated HCl (Table VIII). As .mentioned previously corresponding addition of NaClO., 4  either alone or in the presence of HC10., was without effect* 4 The rate law predicted on the basis of this interpretation i s ,  where k i s the dissociation constant of HCl* HCl i s small, then [H ] £ [CI +  ~ [HC1]  If the fraction of undissociated  q  where [HCl] is the total HCl concentration and equation (20) becomes Q  k  c a t  / [HCl] - k ^ + ( k ^ / ) [HC1] b  Plots of k  K  (21)  0  ,/ [HCl] V S . [HC1 ] at the various solvent composition, cat o o •*  are shown in Fig. 12.  In accord with equation 21 these are fairly linear up  to HCl concentration of about 1 U,  Their intercepts yield values for k^f some-  what lower than those obtained with HC10 . One possible explanation for this 4  is that the H and that h  Q  +  - catalyzed reaction follows an h  Q  rather than H  +  - dependence  for HCl, even at low concentrations, i s appreciably lower than  for HCIO^. It i s interesting to note, however, that the values of k^+ obtained from the intercepts, like those for HClO^, pass through a minimum for the 50 mole percent EtOH solvent composition. A mechanism which would account for the catalytic effect of undissociated +  HCl molecules, involves the simultaneous coordination of H  m  and CI  at the  azo linkage to give an intermediate (IV), analogous to (III), in which the role of the solvent molecule (5) i s effectively assumed by CI**  jrj-  N  .H-<ry-M*.N-<fy CI  w  H (IV)  It i s interesting to note that a similar mechanism may account for the combined catalytic affect of H  +  and thiourea reported by Schulte-Frohlinde  (7). The kinetic dependence i s of the form. ^ ^ ° c a  [thiourea], and can  be accounted for by the formation of an intermediate analogous to (IV) through the reaction of azobenzene with the conjugate acid of thiourea. Also in accord with these suggestions are the results of Nozaki and Ogg (20) for the acid-catalyzed cis-trans isomerization of maleic acid. Different acids were found to exhibit widely differing catalytic activities (HCl, for  - 38 -  example* being more than 50 times as effective as HCIO^) suggesting that the anion also plays an important role i n the reaction* A mechanism* involving the simultaneous coordination of the anion and H (at the double bond and one of the carboxyl oxygen, respectively) was proposed, which i s in some respects analogous to that described above* the structure of the proposed intermediate i s  H0  H H  N  C o C- C - C HO' 1  0 - OH  An analogous mechanism, assuming the formation of an intermediate which involves coordination of both an electron donor and an electron acceptor has also been proposed (21,22) for the catalysis by secondary amines of the c i s trans isomerization of diethyl and dimethylmalmaleate*  Two molecules of  amine are found to react in the rate determining step; one conjugates through N at the double bond the other through H at the carbonyl oxygen* 0 H H » R' C » C - C-C - OR BO ' • BNR H B  \  NH0  V  (c) EFFECT OF ACETIC ACID As shown in Table (IV) no appreciable catalytic effect could be detected for concentrations of up to 1 M of acetic acid in 50 mole percent ethanol at 60°. Referring to our previous discussion on the effect of HCIO^ and HCl i t i s clear that for every proposed mechanism the acid catalysis appears to involve the formation of a protonated species between the uncharged azobenzene molecule and the acid*  It seems reasonable to suggest that the absence of any catalytic  activity for acetic acid i s due to the fact that this acid i s too weak to give rise to any protonated species analogous to (I), (II), or (IV).  -  39  -  (d) EFFECT OF CUPRIC PERCHLORATE The effect of cupric ions on the isomerization rate, vas examined systemo atically at 60 C in 50 mole percent ethanol*  The cupric perchlorate con—  centration vas Taxied up to IM, above which the rate became immeasurably fast*  Zero time uncertainties, due to fast rates, may also account for the  apparent "induction periods" (i.e. failure of the f i r s t order plots to extrapolate back through the i n i t i a l concentrations) reflected in the f i r s t order rate plots in Fig. 13,  Apart from this i n i t i a l effect, however, the  f i r s t order plots were convincingly linear.  The kinetic results are  summarized in Table (V) and Fig. 13 and show the following features: 1.  The specific catalytic activity of CutClO^)^ i s much higher than  that of the two acids previously studied. 2.  The kinetic dependence of k  .on the cupric perchlorate concat centration i s greater than f i r s t order (approximately 1.5) as shown by the plots of k . vs. [Cu(C10j J in Fig. 14 and of log k . vs. log [Cu(C10j 1 cat 4 2 cat 4 2 in Fig. 15. In the interpretation of these results i t seems reasonable to attribute the catalytic effect of Cu(C10.L to the coordination of Cu 4  with the aso-  2  nitrogen atoms to give an intermediate analogous to that suggested for acid catalysis. The significance of the observed deviation from f i r s t order dependence is not clear, but two contributing factors can be suggested: 1.  A medium (e.g. ionic strength) effect, superimposed upon the Cu  concentration effect, when the concentration of cupric perchlorate i s increased^ 2* Contributions from a "second-order" path (i.e. in Cu ) in addition to a f i r s t order one. This would imply a mechanism involving coordination of 11  two Cu  ions with azobenzene possibly through an intermediate.  -  _  40  -  . TIME - MIxNUTES  . .  FIGURE 13. First order kinetic plots for Cu(C10 ) catalyzed cis-trans isomerization 4  of azobenzene in 50 mole % ethanol at 60° C.  2  - 41 TABLE" V EFFECT OF CUPRIC PERCHLORATE AT 60° IN 50 mole % EtOH  HC10„  4  moles/liter  Cu(C10 > 4 2 moles/liter  kxlO  5  (sec *)  k x 10 cat A  (sec  0  •o  5.76  0  n  0.10  9.60  3.84  n  0.20  14.1  n  0.30  19.2  13.4  n  0.40  25.7  19.9  ti  0.50  35.0  29.2  tt  0.64  47.0  41.2  it  0.80  65.6  59.8  it  0.90  80.1  74.3  »t  1.00  90.4  84.6  .050  0  6.07  8.34  0.31  n  0.02  19.5  13.7  »  0.04  27.8  22.0  *  0.06  33.9  28.1  *  0.08  39.5  33.7  .10  0  6.40  0.64  «  0.01  24.1  18.3  »  0.02  30.7  24.9  »  0.04  46.0  40.2  "  0.06  59.8  54.0  - 42 * TABLE V  HC10. 4  Cu(C10 ) 4  2  moles/liter  0  (cont.)  k x 105  k  (sec" )  (see"  .x 10  cat ^  • 20 n  0.01  42.6  36.8  it  0.02  58.9  53.1  it  0.04  84.5  78.7  it  0.06  7.04  137  1.28  131  .32  0  it  .01  57.6  51.8  •i  .02  85.8  80.0  n  .04  137  131  tt  .06  184  178  tt  .08  215  209  .42  0  7.67  8.12  1.91  2.36  it  .01  it  •02  113  107  w  .04  177  171  it  .06  240  234  .52  0  69.4  8.60  63,6  2.83  tt  .01  tt  .02  121  115  II  .03  162  156  n  .04  184  178  II  .06  254  248  it  .08  316  310  81.2  75.4  )  43 -  100  -  44  -  fY (V) While there i s no a priori reason to expect such a complex to form, i t is of interest that the existence of an analogous stables cuprous chloride complex of azomethane i*e*  Cu>  ^CH-  has been reported* (23) The addition of small amounts of perchloric acid in the presence of cupric perchlorate proved to have an unexpectedly large effect, the combined catalytic  effect being many times greater than the sum of the separate  effects of the two catalysts*  This effect was investigated systematically  in a series of experiments in which the HC10, and Cu(C10 ) concentrations 4  4  2  were simultaneously varied (from 0*1 to 0*5 M and from 0*005 to 0*06 M respectively)*  Again, because of high rates and the i n i t i a l complications  discussed earlier, measurements could not be made for Cu  concentrations  above 0*06 M* Typical f i r s t order rate plots for these experiments are shown in Fig, 16 and 17.  At constant acid concentration, plots of k  . versus cat  [CufClO^J^] show a good linear dependence with the exception of the region of very low [ C u ^ J , where the curves bend sharply towards the origin (Fig, 1 8 ) . ^ versus [HCIO^] at constant [-Cu**] (Fig* 1.9) are linear except  Plots of k  for high [H ] (about 0,5 M HCl) ) where they tend to level off. From these +  4  observations i t appears that, over a considerable concentration range of Cu and H , the predominant contribution to the catalytic rate i s of the kinetic +  form:  -  20  4 0  6 0  46  -  8 0  1 0 0  1 2 0  1 4 0  TIME - MIN. FIGURE 16, First order kinotic plots for the cis-trans isomerization of azabenzone catalyzed by Cu(C10 ) in tho presence of 0,1 M HC10 in 50raolo% EtOH at 60° C. 4  2  4  - 50 Rate = k [A][CU "][H*'] H  «  k  cat »  *MC***]  This implies a catalytic path involving the simultaneous interaction of H and Cu  with azobenzene* One possible interpretation is in terms of an  intermediate complex analogous to (V) in which H  and Cu  are coordinated  to the two N atoms* (f)  EFFECT OP CUPRIC CHLORIDE BeoauB* of complications due to solubility limitations the effect of  cupric chloride could be examined only in the presence of hydrochloric acid* Three series of experiments were performed with 0*1, 0.2, 0*3 M HCl and various concentrations of up to 0.04 M CuCl^*  The acid and the cupric salt  concentration we re restricted to this low range because of the extremely high catalytic rate* Typical f i r s t order rate plots are shown in Figs* 20 and 21 and the kinetic results are summarized in Table VI and Fig* 22, It i s seen that at constant acid concentration* plots of k- ^ versus c  [CuClg] are linear over the entire concentration range examined* The tendency for the catalytic activity to be enhanced by the acid i s less marked than for cupric perchlorate* but this may be due to the much lower concentration range to which these measurements were confined* At low HCl concentrations k ^. i s first order in CuCl^ and is much greater than for ca  Cu (CIO ) at comparable HC10  concentrations*  (g) EFFECT OF CUPRIC ACETATE Again solubility limitations precluded addition of cupric acetate alone* One experiment was made with approximately 0*2 M cupric acetate in the presence of 0*5 M acetic acid* A measureable catalytic effect was noted (Table VIII) which was, however, much small than for corresponding perchlorate or chloride solutions* No further quantitative studies were attempted*  - 51 -  20  40 TIME - (MINUTES)  60  FIGURE 20. First order kinetic plots for the cis-trans isomerization of azobenzene catalyzod by CuCl  in the presence of 0,01 M HCl in 50 mole # EtOH  - 52 -  20  40  60  TIME - (MINUTES) FIGURE 21. First order kinetic plots for the cis-trans isomerization of azobenzene catalyzed by CuCl^ in the presence of 0.02 M HCl in 50 mole % EtOH at 60° C.  80  TABLE VI EFFECT OF CUPRIC CHLORIDE VI Solvents  HCl moles/liter  50 mole % EtOH - Temp, 60° C,  CuCl. 2 moles/liter  0.01 M  0  kxlO  5  (sec" ) 1  k .x 10 cat (sec ) -1  5.76  0  n  0.01  35.8  30.0  *  0.02  69.0  63.2  *  0.03  101  "  0.04  169  163  "  0.06  316  310  95.2  *  ,2U NaC10  105  99  »  .2M NaCl  155  149  .02 M  4  0 0.001  5.76  0  8.96  3.20  "  °.005  21.0  15.2  "  0.01  39.8  34.0  *  0.02  78.2  72.4  "  0.03  "  0.04  .03  0  110  104  5.76  0  M  0.005  25.6  1>.8  n  0.01  49.8  44.0  "  0.02  94.0  88.2  "  0.03  146  140  0.04  191  185  w  - 54 -  TABLE V I I EFFECT OF CUPRIC ACETATE Solvent 50 mole # EtOH - Temp. 60° C.  CH COOH  Cu(CH COOH)  mole/liter  moles/liter  3  3  •5M  0  ,5M  0.2  2  k x 10 (sec  5.51 15.3  5  k  x 10  (sec"" ) 1  0 9.50  5  - 54 A -  - 55-  (h) RESULTS IN BENZENE A few catalytic measurements vrere also made in benzene solutions, using heptanoic acid and cupric heptanoate*  Measurements were confined to low  concentration (-2M) of the latter, both because of i t s low solubility and i t s intense colour which interfered with the spectrophotometric method used to follow the reaction* Under these conditions no catalytic effect was noted and this may be attributed to the low concentrations of H and Cu  and their corresponding  low coordinating tendencies in these solutions* (i) MISCELLANEOUS EFFECTS Among the other salts examined were CoCClO^)^ and AgClO^ and some results for these are listed in Table VIII. 0*5 M Co(Cl0 ) (in the presence of 0*1 4  2  M HCIO^) did not appear to have any effect, while 0*5 M AgClO^ gave a catalytic rate only slightly higher than that due to the acid alone. Since the coordination o f the catalyst to the azo-nitrogen atoms appears to be a common feature of a l l the catalytic effects noted, this behaviour may be related to the very much smaller tendencies of Co Cu  and Ag (relative to  ) to form complexes with nitrogen ligands in general*  - 56 * TABLE VIII MISCELLANEOUS RESULTS a. Salt effect Solvent 50 mole $> EtOH - Temp* 60 C. Acid 10.mole/liter 4  k x IO  Salt NaClO.mole/liter 4  k  x IO —1 (sec" )  5  (sec" ) 1  0  0  5.76  0  0  .5  5.76  0  .5  0  8.62  2.86  .5  .5  8.62  2.86  HCl mole/liter  NaCl mole/liter  k x IO  k x IO *_l (sec )  5  c a  (sec "") 0  5.76  0  •2  5.50  0  6.51  •75  6.71  .95  b*  •  2  5  A  0  •2  5  c a  0  Results in benzene— Temp* 60° C*  Heptanoic acid moles/liter  Cupric Heptanoate moles/liter  k x 1O .v (sec"* '  k x IO cat (sec" )  5  ;  5  A  1  0  0  12.7  0  .2  0  11.2  0  •2  .2  14.3  1.6  c* Results with Co(C10 ) and AgC10 in the presence\ of HC10 4  2  4  4  Solvent 50 mole % EtOH - Temp. 60°  co(cio )  k x IO  moles/liter  HCIO^ 4 moles/liter  0.5  0.1  6.4  4  2  5  (sec "")  k  x IO cat_^ (sec" ) 0.6  5  -  57  TABLE V I I I ( c o n H )  AgC10  4  moles/liter  0.5  HCl moles/liter  0.1  k x IO (sec *")  8.3  5  k  c a t  x IO  (sec ) 1  2.5  5  - 58 SUMMARY AND CONCLUSIONS Although a number of common features have been recognized for the different catalysts examined, i t has not been possible to suggest a single mechanism to account for a l l of them* In the case of the H  +  - catalyzed reaction (HCIO^) i t has been postula-  ted as most probable the formation of a protonated species by action of H on azobenzene*  The  +  suggested mechanism involves the subsequent reaction of  the protonated complex with a molecule of solvent to give the intermediate (III) i n which H  +  and the molecule of solvent S are coordinated with the  nitrogen atoms at the azo linkage*  S  The proposed structure f o r (III) iss  H  (III) i n which free rotation about the N-& bond i s possible* The greater catalytic a c t i v i t y shown by HCl i s interpreted i n terms of an additional effect of undissociated HCl molecules* —  involves simultaneous conjugation of CI  The proposed mechanism  +  and H  at the azo-linkage to give  an intermediate (IV) analogous to (III)  Q-H-Q Cl  H  (17) This suggestion i s supported by the fact that analogous mechanisms, assuming the formation of an intermediate which involves both an electron donor and an electron acceptor, have been proposed i n other cases of cis-trans isomerization*  - 59  The predicted rate law i s then  and the results were found to be in accord with this. Acetic acid was found to be inactive. The kinetic dependence of k .on cupric perchlorate was found to be cat of order greater than f i r s t .  Coordination of Cu  with the azo-nitrogen atoms.  to give an intermediate analogous to (III), appears to be probable; to explain higher order path, contribution of two cupric ions to give a complex, Azobenzene.2 Cu  has been postulated.  The greatly enhanced catalytic activity of Cu(C10 ) in the presence of HC10. i s attributed to the formation of a similar complex by simultaneous 4  4  +-  coordination of H  2  11  and Cu  at the azo-nitrogen atoms.  Similar results have been obtained for CuCl^ although the enhancement by HCl i s less marked in this case. +•  4+  The much lower catalytic activities of Ag and Co  ions are attributed  to their lesser tendency to coordinate with nitrogen ligands.  - 60 -  SUGGESTION FOR FURTHER WORK  I t would be of interest t o extend t h i s work and to examine further several problems which developed during the course of this investigation, i n connection with acid catalysis and with the mechanism of cis-trans isomerization. 1.  Among these are:  Determination of h^ f o r the solutions and temperatures used i n this  study with a view to establishing more conclusively whether the acid catalyzed reaction follows an h  Q  o r H* dependence, and hence applying the  Zucker-Hammett criterion to differentiate between the various alternative mechanisms. Alternatively i t might be possible to achieve this by extending the measurements to higher acid concentrations (this may be possible at lower temperatures) or t o aqueous solutions f o r which h^ i s already known. The latter involves some experimental d i f f i c u l t i e s because of the s o l u b i l i t y of azobenzene i n water, but i t might be possible to overcome these. 2.  To test the suggestion of an additional catalytic path due to •fr  undissociated HCl by varying the H using a suitable chloride s a l t .  —  and CI concentrations independently  Some measurements of this type were attempted  with NaCl which, however, proved too insoluble i n aqueous alcohol. 3.  To examine other acids (particularly those of intermediate strength)  with a view to ascertaining whether the behaviour observed f o r HCl i s general and to correlating catalytic a c t i v i t y quantitatively with acid strength. 4.  I n the same connection other substances, known to have both nucleo-  p h i l i c and electrophilic properties, such as amines, should be investigated. These are known (21,22) to catalyze cis-trans converstion of olefenie compounds and i t would be of interest t o ascertain whether they are effective also f o r azobenzene.  - 61 5*  The effect of temperature, should also be investigated in order to  determine the activation energy for the catalysed conversion* 6*  Finally, i t would be very interesting to extend the quantitative  catalytic studies to substituted azobenzenes*  - 62 REFERENCES 1.  Hartley, G.S., Nature, 140, 281 (1937)  2.  Hartley, G.S., J . Chem. Soc. 633 (1938)  3.  Magee, H. Shand, W. and Eyring, II., J . Am. Chem. Soc. 63, 677 (4l)  4. Halpern, J . Brady, G.W. and Winkler, C.A. Can. J . of Research, B. 28. 140 (1950) |  5. Le Fevre, J.R. W. and Northcott, J . J . Chem. Soc. 867 (1953) 6. T a l a t i , E.R. Ph.D. Thesis, Ohio State University. 1957 (Supervised by Dr. Earl W. Malmberg) 7.  Schulte-Frohlinde, D. Liebigs, A. Chem. 612, 131, 138 (1958)  8. Long, F.A. and Paul, M.A.  Chem. Revs. 57. 1, 935 (1957)  9. Hammett, L.P. and Deyrup, A.J. J . Am. Chem. Soc. 54, 2721 (1932) \0. Zucker, L. and Hammett, L.P. J . Am. Chem. Soc. 61, 2779, 2785, 2791 (1939) 11.  Braude, E.A. and Sterne, E.S. Nature 161. 169, (1948)  12.  Braude, E.A. J . Chem. Soc. 1948, 794, 1971.  13.  Braude, E.A. and Sterne, E.S. J . Chem. Soc. 1948, 1976, 1982  -4  14. Jaffe, H.H. and Gardner, R.W.  -*  15. Jaffe, H.H. Yeh Si-Yung, and Gardner, R.W. Molecular Spectroscopy, Z, 120 (1958)  J . Am. Chem. Soc. 80, 319 (1958)  16. Klotz, I.M., Fiess, H.A., Chen Ho, J.Y. and Mellody, M. J . Am. Chem. Soc. 76, 5136 (1954) 17. ^  Hichaelis, L. and Granick, S. J . Am. Chem. Soc. 64, 1861 (1942)  18. Le Fevre, R.J.W. and Vine, H. J . Chem. Soc. 1938 P. 431 19.  Cook, A.H., Jones, D.G. and Polya, J.B. J . Chem. Soc. 1315 (1939)  20.  Nozaki, K., and Ogg, R. J r .  J . Am. Chem. Soc. 63, 2681 (1941)  - 63 21. Nozaki, K. 22.  J . Am. Chem. Soc. 63, 2681 (1941)  Davies, M. and Evans, P.P. Trans. Faraday Soc. 51, 1506 (1955)  23. Dunitz, J.D.  (1957) unpublished.  Quoted by Orgel, L.E.  i n "Hetals and Enzyme Activity", Biochemical Society Symposium No. 15, Cambridge University Press (1958) page. 8.  


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