摘要:

t Vice-Pres&nts Prof D. €I. Everett MBE MA DSC Prof P. Gray MA SCD Prof I. N.Murrell PRD Prof W.C.Price SCD n"~ PRS Honorary Semtay &day lXt6-hM&rs on the Primary Jounurls Committee Pmf R. N. Dixon BSC PHD Prof D. H.Everett MBB .MA DSC bf J. H.h'dl WL PHD SCD FKIC Prof F. C. Tompkins DSC FRIC FRtl AJsiFtmt Editors Prof T. M.Su@n AMSCD FRS Pmf J. S. RowinSon MADPH~me PBS Dr H. k S~BAm D Prof F.C,Ton~pMnsDSC mc m Prof F.C. TompMns DSC mc m Prof p. MA SCD I)r D. HUS& BSC PHDFRIC Prof H.M.FRYBSC DPHIL Dr D.A. Young PRD DSC Prof D.A. King BSC PHD Prof F.C. Tompkins DSC mc mts P.G. Hall BSC M.J. Grant BSC D.A.Young PHD DSC Burlington House London WIV DBN tekphom 01-734 9864

ISSN:0301-5696    

DOI:10.1039/FS97510FX001

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Spiers Memorial Lecture The Development of Ideas about Proton Transfer Reactions BY R. P. BELL Department of Chemistry University of Stirling Scotland Received 29th September 1975 The amount of quantitative information available about proton transfer (acid-base) reactions in solution probably exceeds that for any other class of reaction and this is true both for equilibrium and for kinetic measurements. In view of the importance of such reactions in biological systems and their mechanistic simplicity (at least on paper) it is easy to see why their interpretation and further study has attracted much attention including a Faraday Society Discussion in 1965 and two recent This lecture will attempt to give some account of the early history of the subject and to call attention to some of the main points of current interest.In keeping with most of the contributions to this Symposium the emphasis will be mainly on kinetic problems. The recognition that certain reactions involved the transfer of a proton between two species was implicit in the early days of the ionic theory. Thus if it is accepted that ammonium acetate is ionised its formation from ammonia and acetic acid becomes NH3+ CH3C02H-+ NH; + CH,CO; though even here many chemists would have supposed that in solution this reaction necessarily proceeded through the dissociation of the reactants to give OH-and H+ respectively which then combined. However the generality and importance of proton transfer reactions was first realised in Bronsted's acid-base deBnition first published in 1923,4 and subsequently elaborated in his monograph " Acids and Bases ".5 This definition gave a unified presentation of the dissociation of acids and bases neutralisation hydrolysis of salts indicator equilibria buffer action etc.all of which were traditionally treated as separate topics much to the confusion of the student. In order to include the dissociation of acids in the same scheme it was necessary to acknowledge that the hydrogen ion in aqueous solution is a hydrated proton and an extension of the same ideas showed that the behaviour of acids and bases in non-aqueous media is largely determined by the acid-base properties of the solvent itself. Bronsted's definition owed much to the analogy between proton transfer and electron transfer in redox systems and table 1 compares the two types of function in aqueous solution.It may be noted that many of the differences in usage arise from essentially practical points. In particular redox systems are characterised by potentials rather than by equilibrium constants because they rarely enter into mobile equilibrium with solvent species. For the same reason the accessible range of redox power in aqueous solution is enormously greater than the range of acid-base strength since it is possible to obtain stable aqueous solutions of many species which are thermodynamically capable of reducing water to hydrogen or of oxidising it to hydrogen peroxide or oxygen. Acid-base strengths are normally characterised by the equilibrium constants of their mobile reactions with solvent species but it would 7 SPIERS MEMORIAL LECTURE be equally logical to use the potential of a hydrogen electrode in a solution containing equal concentrations of conjugate acid and base this has been proposed,6 but never generally adopted.TABLE OF ACID-BASE AND REDOX SYSTEMS IN AQUEOUS SOLUTION 1.-COMPARISON ac i d-base redox definition A + B+H+ definition R + 0+e-reaction Al +B2 + B1+A2 reaction R1+ O2+ O1+R2 standard system standard system H30+ + H20+ H+ +HZ+H20+ H,O++e-Reactions nearly always Reactions frequently slow fast hence equilibria hence equilibria difficult to directly measurable. measure. Equilibricm with solvent Equilibrium with solvent always present. usually absent.Hence measurable range of Hence range of redox potentials dissociation constants about -3 V to +3 V corresponding limited to about 10l2 to 1O1O0 in equilibrium constants. corresponding to 0.7 V. Equilibrium constants more Redox potentials more convenient convenient to use than to use then equilibrium constants. "acidity potentisls ". The pursuit of verbal definitions is not usually a profitable pastime in the physical sciences but Bronsted's acid-base definition did have immediate consequences in the field of acid-base catalysis. This had long been attributed solely to hydrogen and hydroxide ions respectively (or the analogous species in non-aqueous solvents) and for many organic molecules plausible mechanisms had been suggested by which the addition or removal of a proton could lead to the observed reaction.Once it was realised that hydrogen and hydroxide ions were representatives of whole classes of acids and bases the possibility arose that the catalytic power of a solution could depend not only on its pH but also on the concentrations of other acids or bases present especially in buffer solutions. This behaviour general acid-base catalysis had already been observed by Dawson and his collaborators in their work on the acetone-iodine reaction,"' but its full significance was not realised until the publica- tion by Bronsted and Pedersen * of their work on the base-catalysed decomposition of nitramide. This 50-page paper which appeared one year later than the acid-base definition is truly remarkable in its scope and repays careful reading to this day it will be referred to several times in the course of this lecture.Bronsted and Pedersen show that the rate of decomposition of nitramide is independent of [OH-] over a considerable range but varies linearly with the concentration of other basic species such as carboxylate anions or amines and they correctly interpret this as evidence for slow proton transfer from nitramide (or an isomer) to these bases. It is interesting to note that the original object of the nitramide work was to obtain a kinetic method for measuring [H+] or [OH-] and hence to study salt effects on the dissociation of weak electrolytes. It is not clear whether the study of general base catalysis was stimulated by the acid-base definition or vice versa but in any event the coincidence was a fortunate one.* The quantitative estimates made by Dawson for the catalytic effect of undissociated carboxylic acids need revising in the light of later views on salt effects and the interpretation of conductivity data but his qualitative conclusions remain largely unaffected. On the other hand several other claims to have detected general catalysis by acids do not survive such re-examination. R. P. BELL The experimental distinction between specific and general acid-base catalysis depends upon the relation between the velocity constants of the proton transfer step and of other processes which follow or precede it. This was first clearly formulated by Bronsted for reactions involving a single proton transfer and extended by Pedersen to processes such as prototropic isomerisations which involve two successive proton transfers.Many of the expressions thus derived have been re- discovered by later workers in connection with particular reactions. Bronsted also showed how experiments could best be planned to avoid complications due to salt effects on equilibria or velocity constants and his procedures are now generally accepted though not always adhered to. log K FIG.1.-Carboxylate ion catalysis in the decomposition of nitramide reproduced from ref. (8). Many studies of general catalysis have followed the pioneer work of Bronsted and his collaborators and no attempt will be made to describe them. In principle simpler behaviour might be anticipated in aprotic solvents where there is no possibility of catalysis by solvent species but in practice complications are often caused by association and lack of information about relevant equilibria.1° Most of the useful results for catalysed reactions therefore relate to aqueous solutions and a considerable amount of information about the rates of proton transfer reactions can be obtained by a correct analysis.However there are severe limitations to the types of system which can be studied in this way since the existence of a catalysed reaction implies that one of the partners in the proton transfer reaction must be unstable or reactive. In most reactions exhibiting general acid-base catalysis proton transfer either takes place to or from carbon (formation of carbonium ions or carbanions) or else is accompanied by drastic changes in other parts of the molecule as in the mutarotation of glucose and related addition reactions of the carbonyl group.Moreover the conventional techniques for studying catalysed reactions were limited to rather slow processes while the much faster thermodynamically favourable reactions of simple oxygen and nitrogen acids were inaccessible to direct measurement. This position has of course been transformed during the last twenty five years by the introduction of new techniques for studying fast reactions notably by Eigen and his collaborators SPIERS MEMORIAL LECTURE and published velocity constants for proton transfers now range over some twenty powers of ten and a wide variety of chemical types.In spite of this wealth of information its quantitative interpretation in molecular terms was slow to develop and is still the subject of much enquiry as shown by the papers presented to this Symposium.* Such interpretations often depend upon the discovery of some empirical quantitative relation and subsequently of deviations from it and for proton transfers discussion has centred around the Brunsted relation. This was put forward by Bronsted and Pedersen * to correlate the catalytic effect of anion bases in the decomposition of nitramide with their basic strength ; the excellence of this correlation is shown by fig. 1 reproduced from their original paper.? Similar results have been obtained for many other reactions in both water and non-aqueous solvents and also for directly measured rates of proton transfer though there are very few instances in which the correlation is as good as for the decomposition of nitramide possibly because the small size of the nitramide molecule minimizes deviations due to steric interactions.The Bronsted relation represents the earliest example of a linear free energy relation and can be expressed as dR(AG*) = P6R(AGo) (1) where AG* is the free energy of activation AGO the standard free energy change for the overall proton transfer reaction (both corrected for statistical factors) dR denotes a change caused by substitution in one or both reactants and /I is a constant originally assumed to be positive and less than unity.The Bronsted relation differs from most rate-equilibrium correlations in that AG* and AGO refer to the same reaction and this together with the low steric requirements of the proton probably accounts for the accuracy with which it is frequently obeyed. Many problems of interest can be expressed in terms of the range of validity of the Bronsted relation and in particular the following questions may be asked :$ (a) What types of structural variation are consistent with the validity of eqn (I)? (b)Has /I the same value for substitution in both reactants and can 6 include variations in the nature of the solvent ? (c) Need the value of p lie between zero and unity? (d)Does p remain constant over a large range of AGO? These queries will now be considered separately.(a) It was always envisaged that a single Bronsted relation would apply accurately only to a series of similar acids or bases for example carboxylic acids or ring- substituted anilines and individual deviations were soon recognised and to some extent explained. Thus in the decomposition of nitramide basic catalysts of differing charge were found to generate different Bronsted relations and plausible explanations were given similar charge effects have been recently reported for the acid catalysed hydrolysis of ethyl vinyl ether.12 Similarly primary secondary and tertiary amines are found to give separate Bronsted plots in the decomposition of nitramide,13 which * Many early workers especially Bronsted himself showed little curiosity about the molecular mechanism of the reactions which they studied.Their attitude might be likened to that attributed to Dr. Johnson who explained his lack of interest in horse-racing by saying " Sir it is already known to me that one horse can run faster than another! " f This plot should strictly speaking be modified slightly to allow for statistical factors which Bronsted and Pedersen applied in an incomplete form. However the correlation is equally good when the modified factors are used. $ Recent developments in the Bronsted relation are discussed in two reviews by Kresge," to which the writer is much indebted. R. P. BELL can be accounted for in terms of the solvation of the amine cations.14 Because of the small size of the proton steric hindrance is not often prominent in proton transfer reactions but there are a number of instances in which 2,6-substituted pyridines react more slowly than anticipated l5 conversely cases are known in which the presence of large polarisable groups in both reactants leads to an abnormally fast reaction presumably because of hydrophobic attraction between these groups.The above effects are all relatively small but if comparisons are made between species of widely differing structures very large deviations from eqn (1) appear. For example although phenol acetylacetone and nitromethane have approximately the same pK-values the velocity constants for their reactions with hydroxide ions are about lolo lo5and 10 dm3 mol-1 s-l respectively. The original explanation of the slow reaction of nitromethane l7 regarded it as a pseudo-acid which could only react with bases after its slow isomerisation to the "true " acid CH2 :NOOH.According to current views the aci-form could only be produced after the reaction of the normal form with base to produce the anion. The reason for the slowness of this and similar reactions is still not fully understood but it is almost certainly connected with the extensive structural and electronic reorganisation involved which may result in a two-stage process. Returning to acid-base systems of similar structures Bronsted and Pedersen pointed out a number of ways in which apparent deviations from the Bronsted relation (regarded as an empirical correlation) could be used to obtain information about equilibria or structures in solution.For example since solutions of carbon dioxide involve the equilibria C02+H,0 $ H2C03*H++HCO; with the ratio [C02]/ [HZC03]= 270 the catalytic power of the bicarbonate ion will be related to the "true " pK of H2C03 3.89 rather than to pK = 6.35 measured by conventional means. Bronsted and Pedersen were not able for technical reasons to test this idea for the decomposition of nitramide but it has been found subsequently that the bicarbonate ion does have an abnormally low catalytic effect in other reactions,20 and this offers in principle a method for determining the degree of hydration of carbon dioxide. Various types of isomerisation equilibrium can also be investigated and we have used this principle recently to investigate the lactol-ketoacid equilibrium HO 0 RCO v\ RC CO I I I I 1 lactol keto-acid carboxylate ion for 27 aliphatic and aromatic keto-acids.2 The overall dissociation constant was determined by conventional means and the "true " dissociation constant of the keto-acid form by measuring the catalytic effect of the carboxylate ion in the decom- position of nitramide.In the few cases where comparison is possible our values for the [lactol] :[keto-acid] ratio agree well with other sources. Mention may be made of two other applications proposed by Bronsted and Pedersen for nitramide kinetics which of course apply in principle to any systematic study of rates of proton transfer for a series of related acid-base pairs. The first refers to the dissociation of an unsymmetrical dibasic acid according to the scheme Hf +-XYH +HXYH +HXY-+H+ where the measurement of the rate of proton transfer to the equilibrium mixture of the ions -XYH and HXY-offers a method of determining the ratio of their concentrations.The second relates to the structures in solution of oxyacids and their anions for example phosphorous acid might exist 1- SPIERS MEMORIAL LECTURE as P(OHj3 OPH(OH), OzPH20H or OPOH. Bronsted and Pedersen show that the different statistical factors which these formulations imply lead to different kinetic consequences and hence to the possibility of distinguishing between them. It would seem worthwhile to investigate further the potentialities of these and similar applica- tions.(6) Although a simple molecular picture of proton transfers predicts that p should have the same value for substitution in both reactants there is no thermodynamic necessity for this to be the case. There are not many investigations in which syste- matic variations of both acid and base have been carried out but it is already clear that unequal values of p will be frequently encountered particularly when the two acid-base pairs are of very different structural types. The most striking example relates to the nitro-alkanes discussed in the next section but the reaction of diketones and keto-esters with carboxylate anions can also give unequal p-values as shown by the results in table 2.22 TABLE 2.-BRONSTED EXPONENTS FOR SUBSTITUTION IN BOTH REACTANTS (SIXCOMPOUNDS OF EACH CLASS WERE INVESTIGATED).acid base P C6H5CH2( MeCO)CHC02E t XC6H4CH2( MeCO)CHC02Et xco; CH,CO 0.44 0.77 xco 1.oo CH2ClCO 0.98 A related question is how far changes of AGO produced by changing the solvent can be represented by eqn (1) with the value of p derived from substitution in the reactants. Although this appeared to be the case in some instances,23 it is now clear that in general the effect of solvent upon the rate of proton transfer bears little relation to its effect on AG". This has been shown particularly clearly by Cox and Gibson,24 and certainly reflects the importance of changes in solvation as discussed in a later section. (c) The assumption that 0 < < 1 implies that the properties of the transition state are intermediate between those of the reactants and the products and in particular that substituent effects which affect only the transition state are unimportant.These conditions are satisfied for most systems but recent work by Bordwell 25 on the reactions of nitroalkanes with bases has revealed a number of exceptions. Whereas the reaction of a single nitroalkane with a series of bases (e.g. carboxylate ions) obeys a " normal " Bronsted relation with p z 0.5 the reaction of a single base with a series of nitroalkanes can give values of p which are either negative or greater than unity thus representing an extreme case of the behaviour described in the last paragraph. These " deviant " Bronsted exponents have been rationalised by different authors in different ways including the superposition of two effects which vary differently with the extent of proton transfer and the dissection of the process into two stages.Since this topic is the subject of another paper in this Symposium 25 it will not be pursued here. (n)The original formulation of the Bronsted relation represented an integrated form of eqn (1) with p assumed constant and it was realised by Bronsted and R. P. BELL Pedersen that plots of In k against In K (i.e. of AG* against AGO) for a series of similar acids or bases would remain linear only over a limited range. In practice it is difficult to detect curvature experimentally without covering several powers of ten in k and K, and such ranges are often attainable only by introducing chemical variations which may conflict with the requirement of a “ similar” series.For example it has recently been claimed 26 that one of the frequently quoted examples of a curved Bronsted plot (for the reaction of bases with compounds containing the carbonyl group) is better represented by two straight lines one for monocarbonyl and the other for P-dicarbonyl compounds. Nevertheless there is now good semi- quantitative evidence for such curvature for a number of reactions 27 and much interest attaches to its theoretical interpretation or prediction. There are three main causes of such curvature. The first arises when the reaction involves two or more consecutive stages and a change in reactivity causes a shift in the rate-limiting step.It has been considered particularly by Jencks,28 and will not be discussed further here. The second is usually referred to as Eigen curuature though it was in fact predicted in the 1924 paper of Bronsted and Pedersen. The argument is that for a highly exoergic proton transfer reaction will take place at every encounter independent of AGO SO that P = 0 similarly for a highly endoergic transfer the rate of the reverse reaction is independent of AGO giving p = 1 for the forward reaction. The complete curve will therefore resemble fig. 2 reproduced from Bronsted and Pedersen with a FIG.2.-Relation between catalytic power and acid-base strength over a wide range reproduced from ref. (8). curved transition region in which p is varying with AGO. The same argument was produced subsequently by Eigeq2’ who gave a more quantitative treatment and showed that the whole transition region would extend effectively over only a few powers of ten in k and thus could not explain the more extended linearity (with p < 1) commonly observed.The third type of effect is termed Marcus curvature and focuses attention on the chemical activation barrier. There is no general reason why the height of this barrier should be a linear function of AGO over an extended range and the simplest inter- pretation of the Bronsted relation in terms of intersecting potential energy curves 30 shows that linearity will be maintained only over the range in which the two inter- secting curves themselves remain linear. Predictions of Bronsted curvature therefore demand some assumption about how the energy profile varies with AGO and this is SPIERS MEMORIAL LECTURE what is done in Marcus' theory originally developed for electron transfers 31 and subsequently applied to proton transfers.32 By assuming that AG* is determined by the point of intersection of two identical parabolae and that the only effect of varying AGO is to displace these parabolae vertically with respect to one another he derives the following expressions AG* = -RTln(k/Z) = (1 +AGo/4AG:)2AG,f p = dAG:/dAGo = +(1 +AGo/4AGz) dp/dAGo = 1/8AG where Z is a collision number and AG: (assumed constant for a given type of reaction) is the so-called "intrinsic barrier " corresponding to AGO = 0.Eqn (2) has also been derived by procedures which superficially differ fundamentally from Marcus' procedure but they all contain the same assumption about intersecting parabolic energy curves or else the equivalent assumption that j? is a linear function of AGO.AG"/AG$ FIG. 3.-Relations between Bronsted exponent and AGO based on intersection Morse curves. (R. P. Bell to be published). There is little doubt that Marcus' theory gives a qualitatively sound basis for understanding the curvature of Bronsted plots and in particular its prediction that the curvature (d/?/dAG") will increase with the intrinsic rate is borne out by experiment. However it suffers from the limitations of the model on which it is based and care must be taken in applying it quantitatively. Thus the fact that p is often found to have different values for variation of the two reactants is evidence against the model of identical intersecting parabolae a fixed distance apart and Koeppl and Kresge have shown 34 that if the curvatures and separation of the parabolae are allowed to vary the resulting plots of p against AGO are sigmoid rather than linear as in eqn (2).Although such plots may be effectively linear over a limited range their slopes differ considerably from 1 /8AG$ and application of Marcus' equations would lead to incorrect values for AG;. If the picture of intersecting potential energy curves is adhered to Morse curves may be more realistic than parabolae and I have recently made calculations on the basis of this Fig. 3 shows the results for three R.P. BELL 15 typical Morse curves it is clear that the curvature may be either greater of less than the Marcus value and may even have the opposite sign. The picture of intersecting curves represents a particular approximation (small overlap) and the opposite extreme is represented by the BEBO treatment which has been applied extensively to hydrogen atom transfers. Marcus has shown 32 that this assumption leads to relations similar to eqn (2) but the quantitative results are again different. The real difficulty here lies in devising a model to represent the forces acting on the proton during its transfer and it is not clear at present whether these differ in any major way from those operating in hydrogen atom transfers. The above considerations apply to the process of proton transfer between two reactants which are correctly positioned and solvated.It is now generally believed that the observed free energies of activation contain a substantial contribution from the energy needed to bring the separated reactants together and to re-organise the solvent and this process is envisaged as taking place before the proton transfer.* This can be allowed for by adding to eqn (2) two further energy terms (for the forward and reverse reactions) usually denoted by wr and wp. In principle it is possible to determine both AG and wr (and sometimes also wp) from experimental curved Bronsted plots and this kind of analysis has been carried out by a number of authors. 27 However the values of AG and w thus derived must be regarded with caution partly because the curvatures are not accurately defined experimentally and partly because of the reservations about energy profiles expressed in the last paragraph.? These problems are of course appreciated by the authors concerned and some of the points are considered in the papers by Marcus and by Hassid Kreevoy and Liang in the present Symposium.The terms w' and wp in the last paragraph are one example of a general problem the role of the solvent in proton transfer reactions. The solvation of the proton and of other ions is of course of prime importance in determining the position of proto-lytic equilibria but we are concerned here more with the kinetic effect of changes of solvation which must accompany the redistribution of charge during a proton transfer.This may be thought of in terms of the solvation of the transition state though this implies an equilibrium situation and it is still an open question whether the re- orientation of the solvent may lag behind the movement of the proton so that equilibrium solvation exists only in the initial and the final states. Much interesting information has been derived from the effect of varying the solvent and more should soon be available from the study of proton transfers in the gas phase though so far the latter have yielded more information about equilibria than about kinetics. In amphiprotic solvents such as water there may be a more intimate involvement of solvent in which proton transfer takes place through one or more intervening solvent molecules.The oldest example of this is the Grotthus chain mechanism for explaining the high mobilities of hydrogen and hydroxide ions in water and a similar explanation has been given for the high rate of the reaction H,O++OH-+ 2H20 though it should be noted that more recent determinations 46 gave a considerably lower value for this rate. Direct evidence for proton transfer through a water molecule was first obtained in n.m.r. studies by Grunwald Lowenstein and Meibo~m,~~ who showed from the broadening of the water singlet that in proton exchange between amine molecules and their cations the indirect process was more important * The separation of these two processes is probably justifiable for electron transfer but less clearly so for proton transfers.iIf the intrinsic barrier heights AG$ are really as low as concluded in the analyses quoted a treatment in terms of parabolae becomes more reasonable however there is an obvious danger of a circular argument here! SPIERS MEMORIAL LECTURE than direct exchange. Many other systems were found subsequently to show similar behaviour which seems to be fairly general for oxygen and nitrogen acids. A special situation arises in the reversible addition reactions of water and similar substances to carbonyl compounds which are catalysed by acids and bases. Reason-able mechanisms involving two successive proton transfers are as follows Acid cataZysis R2C(OH) +HB +R2C(OH)OHi+B-+R2C0+H20+HB Basic catalysis R,C(OH) +B-+R,C(OH)O-+HB +R2C0+H20+B-.However it was pointed out by Eigen 37 that the observed rates in some systems would imply individual velocity constants in excess of the diffusion-controlled limit. He therefore proposed a "one-encounter " or "intimate " mechanism in which both proton transfers take place within a single encounter. Such mechanisms are sterically more plausible if the transition state contains one or more extra water molecules leading to mechanisms such as the following H H H H \ / \ / 0-H-0 0--H-0 I \ I \/ C H C H 1 L /\ /\ I 0 H-0 0-H--0 \ \ H H Support for this view comes from the high orders with respect to water and the large negative entropies of activation which are found when such reactions are studied in solutions of water in non-aqueous though there is no direct evidence for a cyclic transition state and the position may be different in aqueous solution.In these and similar cases it is also debatable whether the two or more proton transfers involved take place synchronously or step-wise. In general it is true to say that there is a growing realisation that solvent participation is important in a large proportion of proton transfer reactions but that there is still much uncertainty about the details of such participation. The final section of this survey deals with the use of liydrogen isotopes in proton transfer reactions. Soon after the discovery of deuterium in 1932 it was realised that isotopic exchange provided a means of studying rates of proton transfer and a number of mechanistic investigations by Ingold and others depended on a comparison of rates of deuterium exchange with other processes such as racemisation or bromina- tion.Subsequently tritium has proved a more useful tool in this respect especially for studying very slow proton transfers. However more interest attaches to kinetic isotope eflects and I shall consider only primary effects though it should be mentioned that in principle measurements in H20+D20mixtures can give information about the number of water molecules involved in the transition state. The current theory of kinetic isotope effects conveniently described as semi-classical derives through transition state theory from the corresponding theory for equilibria first formulated by Urey and by Mayer and Bigeleisen.It has been fully described many times and will not be discussed here.39 In the harmonic approxima- tion the expression for the isotope effect involves only the normal vibration frequencies of the initial and transition states and since computer programs are available for calculating normal frequencies in terms of force constants even for relatively compli- cated systems isotope effects can be calculated for a wide range of models. Since R. P. BELL the transition state frequencies are inaccessible to experiment comparisons of theory and experiment can serve as a method of probing transition state structures and have been widely used for this purpose. A particular problem which has attracted much attention is the variation of kH/kD(or kH/kT)with AGO in a series of similar reactions or with the symmetry of the transition state.It was first suggested by Westheimer 40 that kH/kDwould have a maximum value for a symmetrical transition state for which AG' should be close to zero. This has been confirmed experimentally for a number of proton transfer reactions and fig. 4 gives a further example recently studied in this laboratory.*22 AGO FIG.4.-Hydrogen isotope effects in proton transfer from 3-nitrocamphor to anion bases. (Unpub-lished results by Dr. S. Grainger). Even when no maximum is observed the dirzction in which kH/kDis changed by changes of reactivity can be used to judge whether the transition state is reactant-like or product-like and the result compared with evidence from other sources for example the deviation of the Bronsted exponent from the value one half.Such comparisons make good qualitative sense but it is doubtful whether any quantitative agreement can be expected between different approaches. The concepts of " transi-tion state symmetry " or " extent of proton transfer " are not well defined since they could refer to various properties of the transition state-geometrical position of the proton bond orders ratio of force constants charge distribution etc.-and there is no reason to expect that these properties will all vary in parallel. The semi-classical theory of isotope effects is consistent with a large proportion of the experimental data for proton transfer reactions but there is growing evidence that it does not tell the whole story.A recent survey 41 listed 27 examples of proton transfer reactions for which the isotopic dependence of either the reaction velocity or the Arrhenius parameters (or both) could not be explained by semi-classical theory as well as 25 further reactions involving transfer of hydrogen atoms or hydride ions. I have argued elsewhere 41 that these discrepancies are due to a neglect of the twinel L>flect i.e. the quantum correction which applies to the passage of light particles across energy barriers several authors 42 suggested around 1930 that this *It has also been suggested that the same type of relation would hold between variations of kH/kD and AGO caused by changes in solvent composition and such behaviour has been observed in a few instances.However the work reported by Cox and Gibson at this Symposium *' shows clearly that in general neither the rates nor the isotope effects of proton transfer reactions can be correlated with solvent effects on AGO. SPIERS MEMORIAL LECTURE effect might be important for the movement of protons in chemical reactions but it is only recently that experimental confirmation has been forthcoming. It should be stressed that the tunnel correction has just the same logical status in quantum theory as zero point energy since both are a direct consequence of the uncertainty principle and it is interesting to note that in the harmonic approximation the tunnel correction can be incorporated in the semi-classical expression for the isotope effect merely by including one imaginary frequency to represent motion along the reaction co-ordinate.Moreover model calculations 43 suggest that the tunnel correlation is at least as important as the real vibrations of the transition state and in particular that the variation of kH/kDwith AGO is attributable almost entirely to a variation in tunnel effect the contribution of the real vibrations being negligible. If this result applies generally it means that we can still accept the correlation between isotope effect and transition state symmetry though the reason is no longer that given originally by We~theimer.~~ A number of authors have invoked tunnelling to account for experimental results on proton transfer reactions and in particular Caldin 44 has used it to explain some very large isotope effects and deviations from the Arrhenius equation at low tem- teratures.His latest work reveals a remarkably large solvent dependence of some isotope effects probably attributable to the influence of solvent interaction on the effective mass of the proton and hence on the tunnel c~rrection.~~ This promises a new approach to the investigation of coupling between proton transfer and solvent motion. It should however be mentioned that there are difficulties in the theory of the tunnel correction which do not arise for real vibrations of the transition state. In a complete quantum-theoretical treatment tunnelling would not arise as a separate issue and the usual procedure of multiplying the transition state expression by a tunnel correction is strictly valid only when this correction is small and the same limitation applies to the assumption that tunnelling takes place along a single separable coordinate.Moreover the shape of the barrier for specific reactions is even more difficult to guess than the transition state frequencies though it can be derived from a complete treatment of any model. Thus our conclusion must be that we can have some confidence in the fundamental ideas behind the tunnel correction but that there is still much to be done in relating our models to the reactions which we investigate. The same conclusion applies to several other aspects of proton transfer studies and it is interesting to see how many of the questions raised at the 1965 Faraday Society Discussion on proton transfers are still matters of current concern.Disc. Faraday SOC. 1965 39. R. P. Bell The Proton in Chemistry (Chapman and Hall London 2nd. edn. 1973). E. F. Caldin and V. Gold (ed.) Proton-Transfer Reactions (Chapman and Hall London 1975). J. N. Bronsted Rec. Trav. Chim. 1923 42 718. J. N. Bronsted Om Syre- og Basekatlyse (University of Copenhagen 1926) English translation Chem. Rev. 1928 5,231. E. Wiberg 2.phys. Chem. A 1934 171 1. 'H. M. Dawson and F. Powis J. Chem. SOC. 1913 2135 ; H. M. Dawson and C. K. Reiman J. Chem. SOC.,1915,1426. J. N. Bronsted and K. J. Pedersen,2.phys. Chem. 1924 108 185. K. J. Pedersen Den almindelige Syre- og Busekatalyse (Copenhagen 1932) ; J. Phys. Chem.1934 38 581 ; Trans. Faraday SOC. 1938 34 237. lo For references see ref (2) p. 148. l1 A. J. Kresge Chem. SOC. Reu. 1973 2 475 ; also Chapter 7 of ref. (3). l2 A. J. Kresge and Y. Chiang,J. Amer. Chem. SOC. 1973 95 803. l3 R. P. Bell and A. F. Trotman-Dickenson J. Chem. SOC.,1949 1288; R. P. Bell and G. L. Wilson Trans. Farnday SOC. 1950 46,407. R. P. BELL l4 A. F. Trotman-Dickenson J. Chem. SOC. 1949 1293 ; A. G. Evans and S. D. Hamann Trans. Furuduy SOC. 1951 47 34. l5 R. P. Bell M. H. Rand and K. M. A. Wynne-Jones Trans. Faruday SOC., 1963 85 1773 ; J. A. Feather and V. Gold J. Chem. SOC.,1965 1752. l6 R. P. Bell E. Gelles and E. Moller Proc. Roy. SOC. A 1949 198 308. l7 A. Hantzsch Ber. 1899 32 575. R. P. Bell and W. C. E.Higginson Proc. Roy. SOC. A 1949,197,141 ; R. P. Bell J. Phys. Chem. 1951 55 885. l9 F. G. Bordwell and W. J. Boyle J. Amer. Chem. SOC.,1975.97 3447. 2o F. J. W. Roughton and V. H. Booth Biochem. J. 1938 32 2049; A. R. Olson and P. V Youle J. Amer. Chem. SOC., 1940 62 1027. 21 R. P. Bell B. G. Cox and B. A. Timimi J. Chem. SOC.B 1971 2241 ; R. P. Bell B. G. Cox and J. B. Henshall J.C.S. Perkin 11 1972 1232 ; R. P. Bell and J. B. Henshall J.C.S. Perkin II 1975 39; R. P. Bell and A. D. Covington J.C.S. Perkin II 1975 1343. 22 Unpublished measurements in this laboratory by Dr. S. Grainger. 23 R. P. Bell and B. G. Cox J. Chem. SOC B 1970 194; 1971 783. 24 B. G. Cox and A. Gibson Furaduy Symp. Chem. SOC., 1975 10 107. 25 For references see F. G. Bordwell Faraday Symp.Chem. SOC.,1975 10 100. 26 D. S. Kemp and M. L. Casey J. Amer. Chem. SOC. 1973,95 6670. 27 For references see ref. (11). 28 W. P. Jencks and J. M. Sayer Faraday Symp. Chem. SOC.,1975 10,41. 29 M. Eigen Z. phys. Chem. (Frankfurt) 1954 1 176 ; Angew. Chem. (Int. Edn.) 1964 3 1. 30 R. P. Bell and 0. M. Lidwell Proc. Roy. SOC. A 1940 176 114. 31 R. A. Marcus J. Chem. Phys. 1956 24,966; Disc. Faruday SOC. 1960 29 21 ; J. Phys. Chem. 1963 67 853 2889; Ann. Rev. Phys. Chem. 1964 15 155; J. Chem. Phys. 1965 43 679. 32 R. A. Marcus J. Phys. Chem. 1968 72 891; J. Amer. Chem. SOC.,1969 91 7224; A. 0. Cohen and R. A. Marcus J. Phys. Chem. 1968 72 4249. 33 V. G. Levich R. R. Dogonadze and R. M. Kuznetsov Electrochim. Acta. 1968 13 1025; Elektrokhimiyu 1967 3 739 and later papers ; J.R. Murdoch J. Amer. Chem. SOC.,1972 94 4410. 34 G. W. Koeppl and A. J. Kresge J.C.S. Chem. Comm. 1973 371. 35 R. P. Bell to be published. 36 E. Grunwald A. Loewenstein and S. Meiboom J. Chem. Phys. 1957 27 630. For later summaries see A. Loewenstein and T. M. Conner Ber. Bunsenges. 1963 67,280 ; also chapter 4 of ref. (3). 37 M. Eigen Disc. Faruday SOC. 1965 33 7. 38 R. P. Bell J. F. Millington and J. M. Pink Proc. Roy. Soc. A 1968,303 1 ; R. P. Bell and P. E. Sorensen J.C.S. Perkin 11 1972 1740. 39 See e.g. C. J. Collins and N. S. Bowman (ed.) Isotope Eflects in Chemical Reactions (Van Nostrand New York 1970). 40 F. H. Westheimer Chem. Rev. 1961 61 265. 41 R. P. Bell Chem. SOC. Rev. 1974 3 513. 42 F.Hund Z. Phys. 1927 43 805 ; D. G. Bourgin Proc. Nat. Acud. Sci. 1929 15 357 ; R. M. Langer Phys. Rev. 1929 34 92 ; S. Roginsky and L. Rosenkewitsch Z. phys. Chem. B 1930 10 47; E. Wigner Z. ghys. Chem. B 1932 19 203; R. P. Bell Proc. Roy. SOC.A 1933 139,466; C. E. H. Bawn and G. Ogden Truns. Furudzy SOC. 1934 30,434. 43 R. P. Bell W. H. Sachs and R. L. Tranter Truns. Faruduy SOC. 1971 67 1995. 44 E. F. Caldin Chem. Rev. 1969 69 135 and later papers. 45E. F. Caldinand C. J. Wilson FarudqSyrnp. Chern. SOC., 1975 10 121. 46 G. Brikre and F. Gaspard J. Chem. Phys. 1967 64 463 ; G. C. Barker and D. C. Sammon Nature 1967 213 65.

ISSN:0301-5696    

DOI:10.1039/FS9751000007

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Enthalpiesof Formation and Solvation of Some Organic Anions BY EDWARDM. ARNETT,DALEE. JOHNSTON,LEONARD AND E. SMALL AND DUMITRU OANCEA Department of Chemistry University of Pittsburgh Pittsburgh Pennsylvania 15260 U.S.A. Receiued 30th April 1975 Procedures are described for measuring AH, the enthalpy of deprotonation for organic Bronsted acids in DMSO. Evidence from a number of sources is presented to demonstrate that the deproton- ation process is clean and complete. Correlation with Bordwell's pKa data shows that AH; = AGb in DMSO for the compounds considered here. Good correlations are also found between A&(DMSO) and AHD(gas) for different classes of compounds. Solvation enthalpies are calculated by combining heats of deprotonation in the gas phase and in DMSO with heats of solution of the precursor acids.Comparison is made between resonance- delocalized carbanions enolate anions halide ions and alkoxide ions. The results are rationalized in terms of ionic size charge density and steric hindrance to solvation. The question of anion solvation is an essential one for elucidating acid-base interactions in solution. In addition such matters as salt effects on rates and equilibria the lyotropic series of protein denaturation salt rejection in reverse osmosis and solvent effects on nucleophilicity are of great topical significance in many areas of physical biological and applied chemi~try.~ Of particular concern to the organic chemist however are the many base-catalysed reactions-substitutions eliminations and condensations-which compose about 50 % of the " bread-and-butter " reactions of synthetic chemistry.These reactions are initiated through the conversion of an unreactive neutral organic molecule to a highly reactive anion through abstraction of a proton by a strong base. It is now recognized that solvation and ion-pairing of such organic ions can play a deciding role in their reactivity. This article presents a quantitative analysis of the effects of structural change on the solvation energies of a broad variety of organic anions from the gas phase to dimethyl sulphoxide (DMSO). We will use a simple Born cycle in conjunction with newly available experimental data. Although some of the data may require modification as new results are forthcoming the overall consistency of the results strongly enforces their validity and the (fairly obvious) conclusions which we will draw from them.It is widely understood that solvent effects on isodesmic proton-transfer equilibria A-+A'H+ A'-+AH (1) can be resolved into components attributable to the neutral acids AH and A'H and to their conjugate base anions. Some tentative conclusions can be made on the relative importance of specific solvation of neutrals and anions through studying their thermodynamic properties for transfer from one solvent to an~ther.~ However 20 E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA the ultimate assignment for the role of solvation to each species can only be made by measuring reaction (I) in the gas phase in the absence of solvent and then separating the solvation energies of the neutrals and their ions through the following cycle.AHD A-+A’H -+ AH + A’-(gas) I1 I I AH,(A-) AH,(A’H) AHs(AH) I AHs(A‘-) 5.5. 1 1 AHD A-+A’H + AH + A’-(DMSO) We will be concerned here with enthalpy comparisons since it is much easier to measure or estimate heats of solution of neutral molecules [i.e. AH,(AH)] in DMSO than to measure their Henry’s Law constants as needed to obtain the corresponding free energy terms. Furthermore we have developed a simple procedure for corn-paring heats of deprotonation in DMSO which is applicable to a great number of organic acids including many compounds for which comparable free energy data (i.e.? pKa’s) cannot be easily measured.We shall see however that the enthalpy of deprotonation in DMSO and corresponding free energy are generally equivalent. To apply the above cycle to the calculation of anion solvation enthalpies the necessary data are obtained and treated in the following way (a)All acids are compared to cyclopentadiene as the standard acid. Its anion is the standard anion. Thus 6AHf’DMSo(AH)means “the partial molar heat of solution (i.e.? solvation enthalpy) of AH from the gas phase to DMSO minus the corresponding value for cyclopentadiene” . (b)Heats of deprotonation in the gas phase have been determined by McIver ’-’ using ion cyclotron resonance; by Bohme using flowing afterglow and by Kebarle using high pressure mass spectroscopy. We shall use Kebarle’s recent results l3 for most of our carbon acids.These will be compared with published data for alcohols and alkoxide ions. (c)Heats of deprotonation in DMSO using the potassium lyate salt (K+DMSYL-) are readily determined by solution calorimetry using a method described fully in previous publications.14* (d) The most dubious property in the cycle for many of the compounds reported here is the heat of solution from the gas phase to DMSO i.e. 6AH:’DMSo (AH)* In principle this is readily calculated by subtracting the heat of vaporization (or sublimation) from the heat of solution of pure liquid or solid AH in DMSO. In practice we find that the necessary heats of vaporization or sublimation have either not been measured or (even worse) have been reported but are obviously wrong.’ 6* l7 In order to compensate partially for the lack of vaporization enthalpies we have introduced a substitute approach (see Experimental and Results).We note that the heats of vaporization per se are not needed for our analysis or even the actual heats of solution of the various compounds from the gas phase to DMSO. What we do need are the rehire heats of solution from the gas phase to DMSO. There is a considerable body of evidence l8 to show that in general such differences are rather small and additive unless hydrogen-bonding to solvent occurs. Furthermore the relative strength of hydrogen-bonding interactions in 6AHj’DMSo(A) should be equal or proportional to the “ relative ” heat of transfer from an “ inert” solvent (such as CCl,) to DMSO.We may therefore propose that relative heats FORMATION AND SOLVATION ENTHALPIES of transfer from CC14 to DMSO should be roughly proportional to relative heats of solution from the gas phase to DMSO. As seen below this correlation holds to about k 0.5 kcal/mol for five model compounds whose heats of vaporization are reliably known. We doubt if it produces an error greater than 2 kcaI/mol in any compound in this study. Finally (e) In order to calculate relative heats of solution for anions from the gas phase to DMSO we combine terms as follows l4 6AH!'DMSo (A-) = SAHD(DMS0) -6AHD(gaS) -k 6AHf'DMSo (AH). EXPERIMENTAL MATERIALS All compounds were commercially available but were purified until homogeneous to gas liquid chromatography and the physical properties such as refractive index or melting point agreed with reliable literature values.DMSO was carefully purified as before15 by vacuum distillation from n-butyl lithium. Water concentration was maintained below 50 ppm (Karl Fischer titration) by storage over 4-A Molecular Sieves under argon in a black-taped bottle. K+DMSYL- was prepared as before15 using potassium hydride the working concentration of 0.1 M of lyate salt being established by quenching an aliquot in water and titrating with standard HCl. CALORIMETRY The 250 ml adiabatic solution calorimeter and its use for determining AHs and A& have been described fully.15 All measurements were made at 25°C and involved seven to twelve replica injections of 50-1OOmg.of AH into the calorimeter liquid -DMSO K+DMSYL- solution or CC14. Most enthalpy values are precise within a standard deviation of & 0.25 kcal/mol. Reproducibility of the system was checked frequently by determining A& for fluorene the value -18.2k0.4 kcal/mol being a well established standard in this laboratory. RECOVERY EXPERIMENTS As part of the evidence (infra vide) that the acids in this study were undergoing clean reversible deprotonation a standard quenching and recovery routine was developed. A 2 pl sample of the acid under study at 0.005 M in DMSO was injected into the water carrier stream of a DuPont 830liquid chromatograph at 1500 p.s.i. pressure. After passage through an octadecylsilane column the elution time and peak area of the acid were determined using a u.-v.detector at 254 nm. Following this a sample of the K+DMSYL- solution containing deprotonated acid in DMSO was injected. The anion of the acid and DMSYL- anion were neutralized instantly and in all cases except cyclobutanone cycloheptatriene and acetonitrile the elution volume and peak area of the quenched acid corresponded to that for the original DMSO solution of the acid. RESULTS Evidence that deprotonation of all compounds considered here was clean complete instantaneous and reversible in 0.1 M K+DMSYL/DMSO rests on the following facts. (1) Exothermic heats of reaction were displayed on the calorimeter strip chart recorder immediately after sample injection following which the recorder trace continued parallel to the base line.Heats of reaction were not concentration-dependent. Furthermore addition of dicyclohexyl-18-crown-6 polyether which we have shown l5 to dissociate potassium ion pairs through cation complexing had no effect on AH, for acetone fluorene phenylacetonitrile cyclopentadiene or acetyl-acetone B. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA (2) Proton magnetic resonance spectra of DMSO-DMSYL-d6 solutions of deprotonated acids corresponded in every way with DMSO-d6 solutions of the acid precursors except for the presence of the acidic proton. (3) Recovery experiments using liquid chromatography described above showed no decomposition for the compounds considered here. In view of the sensitivity of many of these compounds to base-catalysed reactions other than deprotonation this evidence is crucial.(4) A fourth piece of evidence (with many more ramifications) arises through comparison of our AHD data with the pK,’s for the same compounds determined in Bordwell’s laboratory. Although Bordwell’s measurements involve an indicator titration at roughly one-hundredth the base concentrations used by us the correlation of our results with his for 43 compounds has a correlation coefficient of 0.988 and a slope of 1.32f0.03 as shown in fig. 1. This slope corresponds within experimental error to 2.303 RTat 25” with the extraordinary implication that the relative standard free energies for deprotonation are exactly equal to their heats of deprotonation. This close agreement strongly supports the validity of both sets of measurements.pK vs AHEMSO -20--15--10-8 10 12 14 16 I8 20 22 24 26 28 30 32 PK FIG.1.-Correlation of pKa’s of Bronsted acids in DMSO (kindness of Prof. Bordwell) with heats of deprotonation in DMSO AH,. DISCUSSION The heat of deprotonation AH,(DMSO) measured in K+DMSYL-in DMSO is a physical property of singular value for comparing the Bronsted acidities of organic acids. Its particular virtues are (a)it is the only method currently available which can be applied directly without extrapolation or interpolation to the entire range of proton donors from strong oxygen acids such as carboxylic or phenolic ones whose pKa’s lie midway in the pH range to relatively weak carbon acids such as acetonitrile or diphenylmethane with pKa’s close to 34-the value for DMSO itself; (b)the FORMATION AND SOLVATION ENTHALPIES DMSO/DMSYL- system has been investigated extensively so that our results are directly relatable to pKa studies by Bordwell and Ritchie,lg to a host of kinetic and ion-pairing studies 2o and to an enormous variety of recent synthetic work; (c) it is relatively convenient to determine heats of solution of neutral organic acids in DMSO in order to calculate the solvation enthalpies of their conjugate base anions as we shall do here.In this section we shall comment briefly on the relationship between AHD(DMS0) values and two other criteria of acidity for the same compounds-their pKa’s in DMSO and their acidities in the gas phase.We shall then consider the factors which influence solvation energies of the anions by comparing them with the heats of anion formation in the gas phase. CORRELATION OF AHD(DMS0) WITH pKa (DMSO) Fig. 1 portrays the surprisingly good correlation between AHD and pKa for a rich variety of organic acids in DMSO. As noted the correlation is significant not only because of its generally clean linearity over a wide range of energy for many structural types but also because its slope implies equality between AH; and AG;. One consequence here is that AS& the standard entropy for deprotonation for these manifold species is nearly constant and probably equal to that for deprotonating DMSO since the deprotonation reaction is AH+K+DMSYL- K+ +A- +DMSO. In an earlier publication l5 we noted a rather good correlation between the first AHD values we determined and a few pKa’s in DMSO which were available at that time.It is both gratifying and surprising to us to find further experimentation not only supporting the original conclusions but strengthening them. At this point the only series of compounds which shows a systematic displacement from the line is the nitroalkanes. Cyclobutanone is also far from the line but as we have shown by the recovery experiment it undergoes a fast exothermic secondary reaction. CORRELATION OF AHD(DMS0) WITH AHD(gaS) Bordwell and coworkers 21 have recently reported a close correlation between the pKa’s of a number of carbon acids in DMSO and corresponding gas phase acidities reported by McMahon and Kebarle.22 In view of the correlation of our data with Bordwell’s a linear relation of AH,(DMSO) to gas phase acidities is required.Such a plot is shown in fig. 2 in which we include many new data graciously provided by Prof. Kebarle. Also included are some previously reported l4 gas phase data from McIver’s laboratory. It is important to recognize here that heats of deprotonation in the gas phase are usually equal within experimental error to standard free energies of deprotonation for isodesmic processes like those considered here. Kebarle’s values were obtained by high pressure mass spectrometry at 600 K while McIver’s acidities were determined by ion cyclotron resonance at 298 K. Taft 23 has shown the generally good agree- ment between data collected by the two groups.However some discrepancies do exist the most glaring being 3.7 kcal/mol for p-chlorophenol. In order to minimize artifactual differences between methods and observers all of the AHD(gas) values used here are taken arbitrarily from Kebarle’s work except for the aliphatic alcohols. The overall trend towards correlation of our results with the gas phase acidities is obvious from fig. 2. Furthermore we find as did Bordwell a partial separation of hydrocarbon acids from oxygen acids such as alcohols ketones phenols and carboxylic acids. The separation suggests (but does not require) an added exothermic E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA term for the ionization of the latter compounds in solution.Special factors also seem to be affecting the ionization of the alcohols thus rendering them much more sensitive to structural change than the other compounds. 310 t -AHD(DMSO)/kcal mol-' FIG.2.-Plot of enthalpies of deprotonation in DMSO [AHD(DMSO)] against enthalpy of deproton-ation in the gas phase [AH~(gas)]. All values are in kcal/mol and gas phase data were provided by Prof. Kebarle and McIver. Correlation coefficient of points on top line = 0.990 and on bottom line = 0.998. Obviously the differences between acidity orders in solution compared to those in the gas phase can be caused by solvent interactions with the neutral precursors or the anions or both. It is important to resolve these factors since such knowledge can suggest the proper point of attack for manipulating proton transfer energies through variation of molecular structure or solvent.This we shall do in the next section. SOLVATION ENTHALPIES OF ORGANIC ANIONS Heats of solution of a number of typical organic anions from the gas phase to DMSO were calculated using the approach described in the introductory section of this paper. The necessary data for the calculation have been published recently 14* 24 for many of the anions treated here. However we have added several more for the present discussion. The solvation energy of an ion is that part of its energy in solution which derives from its net interactions with the solvent. It is informative to compare this external interaction of the ion with the energy required to form it in the absence of solvent- AH,(gas).In fig. 3 the two properties are plotted. Although the overall pattern is FORMATION AND SOLVATION ENTHALPIES a random scatter we have been so bold as to separate it into several lines which include ions with common structural features. The interpretation of the factors which lead to these groupings seems so reasonable that we do not hesitate to repeat them here. 8 -s-23 --E 0;- -50 -40 -30 -20 -10 0 10 20 30 4 I deprotonation energy SAH:/kcal mol-l FIG.3.-Correlations of solvation enthalpies of anions from the gas phase to DMSO [8AH:*DSMo (A31 with deprotonation enthalpies of corresponding acid precursors in the gas phase SA&(gas). All values referred to cyclopentadiene and cyclopentadienyl anion as standards.All ions except alkoxide ions fall on three lines which are mutually parallel as shown. On each of these lines the ions which enjoy the highest degree of charge delocalization because of resonance or large size have the least exothermic values of 8AH:-DMSo(A-). The series of ions whose solvation energies are the least are the hydrocarbanions at the top of the graph. Included with these are p-nitrophenolate and malononitrile anion whose charges must be very widely dispersed. The next line includes delocalized oxyanions of all sorts enolates carboxylates and phenolates. We believe it is significant that the carbanions and enolate anions of most localized charge fall at the intersections with the lines generated by the alkoxide ions whose charge is mostly localized on oxygen.We have rationalized the steep slope of the alkoxide ion line in terms of steric hindrance to solvation of the localized charge a factor which differentiates alkoxides from the other ions in the plot. In the interests of objective scepticism we also warn the reader that this series of ions was studied in the gas phase under different conditions from the others. Also it is possible but unlikely that in solution they are not completely ionized due to their notorious proclivity for homoconjugate ion formation RO--HOR.25 We doubt if these factors are a problem here. Finally we note that the halides which are relatively small and sterically unpro- tected have relatively high solvation energies and generate a line parallel to the two above it.Solvation energies of ions and molecules are generally related to their volumes-in E. M. ARNETT D. E. JOHNSTON L. E. SMALL AND D. OANCEA the former case charge density and in the latter the energy of cavity formation are the chief variables. One might suppose in the present case that in keeping with electro- static principles the solvation energies of the organic ions considered here would be related to their reciprocal radii. Since the ions are generally non-spherical and their partial molar volumes unknown we have crudely estimated r-l for the ion through using the molar volume of the neutral acid precursor and then taking the reciprocal of its cube root. Fig. 4 shows that even with the many rough approximations in this approach a clear trend towards linearity is found.This gives added support to the general consistency of the methods and data used in deriving the solvation energies. 15 t ._ c -13- 0 EtOH rcl x m WE -17-2l- 0 MeOH cl .- crj u >- -25- % -29 - H2O -331 I I 1 0.I 0.2 0.3 0.4 FIG.4.-Correlation of anion solvation enthalpies with estimated reciprocal ionic radius. We are glad to acknowledge the helpful contributions of Prof. Bordwell Kebarle and McIver to this research. Supported by NSF grant G.P. 6550-X. On leave from the Department of Physical Chemistry University of Bucharest Romania. See E. M. Arnett N. J. Hornung and R. J. Minasz Proceedings of C.N.R.S. Conference on Water in Biological Systems Roscoff France June 1975 for recent references relating to these topics.B. G. Cox in Ann. Rep. Progr. Chem. A 1973 70. ’R. T. McIver JI. J. A. Scott and J. M. Riveros J. Amer. Chem. SOC.,1973 95 2706. R. T. McIver Jr. and J. H. Silvers J. Amer. Chem. SOC., 1973 95 8462. ’R. T. McIver Jr. and J. S. Miller J. Amer. Chem. SOC.,1974 96 4323. * D. K. Bohme E. Lee-Ruff and L. B. Young J. Amer. Chem. SOC.,1972 94 5153. D. K. Bohme E. Lee-Ruff and L. B. Young J. Amer. Chem. Soc. 1971,93,4608. FORMATION AND SOLVATION ENTHALPIES lo T.B. McMahon and P. Kebarle J. Amer. Chem. SOC.,1974,96 5939. P. Kebarle presented at Structure-Energy Symposium San Juan R.R. ; 1974. l2 P. Kebarle in Modern Aspects of Electrochemistry ed. B. E. Conway and J.O’M. Bockris (Plenum Press N.Y. 1974) 9 1. l3 P. Kebarle personal communication. l4 E. M. Arnett L. E. Small R. T. McIver Jr. and J. S. Miller J. Amev. Chem. SOC.,1974 96 5638. Is E. M. Arnett and T. C. Moriarity J. Arner. Clzem. SOC.,1971 93,4908; 1973 95 1492. l6 J. S. Chickos J. Chem. Ed. 1975 52 134. l7 G. W. Thomson in Technique of Organic Chemistry Vol. 1 Part 1 ed. A. Weissberger (Wiley Interscience New York 3rd Edn. 1965 ) p. 401. R. Fuchs and R. F. Rodewald J. Amer. Chem. SOC.,1973 95 5897. Also personal com- munication with Prof. Fuchs. l9 F. G. Bordwell Faraday Symp. Chem. SOC. 1975 10 100. Provides references to the relevant literature. 2o D. J. Cram Fundamentals of Carbanion Chemistry (Academic Press New York N.Y. 1965) ; M. Szwarc Ions and Ion Pairs in Organic Reactions Vol.2 (John Wiley and Sons New York N.Y. 1974). 21 F. G. Bordwell J. E. Bartmess W. S.Matthews,G. E. Drucker and Z. Margolin,J.Amer. Chem. SOC.,1975 97 3226. T. B. McMahon and P. Kebarle J. Amer. Chem. Soc. 1974,96 5940. 23 R W. Taft in Proton Transfer Reactions ed. E. Caldin and V. Gold (Chapman and Hall London 1975). 24 E. M. Arnett Dale E. Johnston and Leonard E. Small J. Amer. Chem. SOC., 1975 97. 25 J. H. Exner and E. C. Steiner J. Amer. Chem. SOC.,1974 96 1782.

ISSN:0301-5696    

DOI:10.1039/FS9751000020

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Hydrogen-Bonded and Ion-Pair Complexes in Aprotic Solvents BY J. E. CROOKS Chemistry Department University of London King’s College Strand London WC2R 2LS AND B. H. ROBINSON Chemistry Building The University Canterbury Kent Received 2 1st May 1975 The proton-transfer reactions between acids and bases in aprotic solvents of low relative permit- tivity lead to the production of ion-pair complexes via hydrogen-bonded reaction intermediates. For most systems the rate-limiting step is the formation of the hydrogen-bonded intermediate. The rates of these reactions as measured by the temperature-jump technique agree well with those calculated by a refined theory of diffusion-controlled rates which takes into account rotation within the encounter complex.The reactions of Bromophenol Blue with pyridine basis are anomalously slow the rate-limiting step being proton-transfer within the hydrogen-bonded complex to form the ion-pair. Values of the enthalpy entropy and volume of activation and the primary kinetic isotope effect enable the structure of the transition state to be deduced. The reaction between a proton-donor and a proton-acceptor in an aprotic solvent of low permittivity may lead either to the formation of a hydrogen-bonded complex or an ion-pair. For example pyridine and phenol form a hydrogen-bonded complex whereas triethylamine and 2,4-dinitrophenol form an ion-pair. A considerable amount of charge generation and separation occurs in the forma- tion of an ion-pair as shown by the high dipole moments of these complexes,1 so that a considerable amount of energy must be supplied.Ionization is often observed in aqueous solution where the strong ion-solvent interactions provide the necessary energy but these interactions are much weaker for aprotic solvents. The con-sequence is that only hydrogen-bonded complexes will be formed unless energy can be supplied by some other process. Electron delocalisation may provide the necessary 29 HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS energy. If the loss or gain of a proton by one of the molecules permits more extensive electron delocalisation or in other words establishes a conjugated system then ion-pair formation is favoured. For example the negative charge on the 2,4-dinitrophenol anion is delocalised over the nitro-groups.The establishment of a conjugated system is typically accompanied by a large bathochromic and hyper- chromic shift of the visible/u.-v. absorption spectrum so that ion-pair formation may be observed spectrophotometrically. For example the 2,4-dinitrophenol-triethylamine ion-pair is yellow having an extinction coefficient of 8130 dm3 mol-l cm-l at 400nm in toluene solution whereas 2,4dinitrophenol itself only has an extinction coefficient of 102 dm3 mol-1 cm-l at this wavelength. It seems reasonable that the formation of an ion-pair from an acid and a base must involve a hydrogen-bonded complex as reaction intermediate. An ion-pair differs from the corresponding hydrogen-bonded complex only in the location of the proton; the conversion of a hydrogen-bonded complex to an ion-pair usually requires only the motion of the proton along the hydrogen bond.A general scheme for the formation of an ion-pair may thus be written as k12 k23 AH+B +C +D k21 k32 where AH is the acid B the base C the hydrogen-bonded complex and D the ion-pair. If the structures of AH and B are such that charge delocalisation does not occur as a consequence of proton transfer the equilibrium constant for the second step K23 is extremely small If this is so the only observable product of the reaction is C. If on the other hand charge delocalisation does occur the predominant product is D. The existence of a small quantity of C in equilibrium with D may be overlooked as the absorption spectrum of the mixed products is dominated by that of D.However if K23 is near unity the value of the overall equilibrium constant K where (calculated assuming that [C] is negligible) is found to be concentration-dependent. 3-5 If the existence of C is taken into account the absorbance A of the ion-pair peak of a solution containing stoichiometric concentrations [AH] and [B] of acid and base is given by ti where E is the extinction coefficient of D. Thus a plot of [AH],e/A against [B] has gradient K-' and intercept (1 +K231-l. Hence values of Kl 2 the equilibrium constant for the formation of C may be found since K12 = K/K23. (4) Values of K12 may also be estimated for systems for which C is low by comparison with data from model systems. For instance it is found that values of KI2 for phenol with aromatic amines are on average 2.3 times greater than for Magenta E (I) with aromatic amines probably due to formation of an intramolecular OH .. .Br hydrogen bond in Magenta E. Values of K12 for Magenta E with the stronger aromatic amines can be measured directly by use of eqn (3) but values of K23 for Magenta E with J. E. CROOKS AND B. H. ROBINSON aliphatic amines are too large for K12to be evaluated in this way. However values of K1 for Magenta E with aliphatic amines may be estimated by dividing the values found for phenol with aliphatic amines by a factor of 2.3. OH I Corresponding values of K, for Bromophenol Blue (11) have been estimated from infra-red absorption measurements in CC14 solution by the use of 2,6-dibromophenol as proton-donor since it resembles Bromophenol Blue except in that it lacks the ability to form a conjugated system on proton loss.OH I1 A third way of measuring K12is to observe the hydrogen-bonded complex in the visible/u.-v. spectrum since a small bathochromic shift (-20 nm) occurs on hydrogen bond formation. Values of K, for Magenta E with the weaker aromatic amines have been measured by this method. Thermodynamic parameters obtained by these methods are listed in table 1. FAST REACTIONS INVOLVING THE FORMATION OF ION-PAIR COMPLEXES The kinetics of formation of ion-pair complexes in aprotic solvents for a range of acids and bases have been studied by the temperature-jump technique. The reaction is monitored by kinetic spectrophotometry of the ion-pair absorption peak.The experimental data takes the form of values of the relaxation time robs for solutions of various values of [AH]and [B]. For all the solutions studied it is found that each solution has only one observable relaxation time ;plots of 7%; against [AH]+[B] are straight lines with positive gradients and finite intercepts. This is in accordance with the simple kinetic scheme kr AH+B+D kb 32 HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS for which 7-l = kf([AH]+[B]) +kb. (5) The ratio of the gradient to the intercept kf/kb,is predicted to be equal to the observed overall equilibrium constant K as defined by eqn (2) and this is in fact observed.For all systems studied except those involving Bromophenol Blue reacting with aromatic amines which are discussed later values of kf are found to be in the region TABLE OF HYDROGEN-BOND FORMATION 1.-THERMODYNAMICS K12 at 25°C -AH102 -AS92 /dm3mol-1 /kJ mol-1 /JK-1 mol-1 ref. phenol t rie thy lamine 71 33 74 4 (solvent CC14) tri-n-propylamine 18 25 59 4 tri-n-butylamine 23 29 71 4 2-met hylpyridine 61 29 63 4 2,6-dimethylpyridine 83 29 60 4 2,4,6-trimethylpyridine 117 31 66 4 2,6-dibromophenol pyridine 7.7 20 50 5 (solvent CC14) 2-met hylpyridine 7.7 21 53 5 2,6-dimethylpyridine 7.4 26 71 5 2,4,6-trimethyIpyridine 10.9 24 59 5 Magenta E (solvent C6H5C1) 2-methyl p yridine 28 25 55 4 2,6-dime thylpyridine 31 31 75 4 2,4,6-trime thylpyridine 55 32 74 4 Bromophenol Blue 3-chloropyridine 3 -9 19 51 5 (solvent C6H5Cl) Bromophenol Blue 2,4,6-trimethylpyridine 53 29 64 3 (Solvent C6H5Cl) (second complex)* pyridine 20 26 60 14 pyridine 25 27 64 pyridine 29 23 48 pyridine 35 32 77 *Bromophenol Blue is a dibasic acid.Strong amines react quantitatively with the first acidic group at the concentrations used and the data refer to reaction of the amines with the second acidic group. This reaction is very similar to that between Magenta E and bases. 10s-109 dm-3 mol-' s-l as shown in table 2. Values of kf are independent of K but decrease with increasing steric hindrance round the base. The results can thus be interpreted in terms of a simple mechanism in which the rate-determining step is the diffusion together of AH and B the subsequent proton-transfer being fast.However the possibility of a hydrogen-bonded intermediate cannot be ignored. A system involving two equilibria has in principle two relaxation times although only one may be experimentally observable. The general solution is complex but may be simplified by making either of two assumption^.^ J. E. CROOKS AND B. H. ROBINSON If the first equilibrium is established much more quickly than the second the relaxation times are given by A plot of Ti1 against [AH]+[B] thus gives a straight line for which the ratio of gradient to intercept is k12/k21,i.e. K12. As the observed ratio is K r1 cannot be robs. Eqri (7) is however compatible with the data if as is so in practice Kl2([AH]+ [B]) < I which implies K239 1.On this scheme the gradient of the r$ against [AH]+[B] plot i.e. k, is identified with k23K12. However if proton transfer were the rate-determining step k23 and hence k, should increase with increasing K which is not observed for the reactions in table 2. TABLE 2.-RATES OF FORMATION OF ION-PAIR COMPLEXES (CHLOROBENZENE SOLVENT) K 10-8 x kt temp. acid base /dm3 mol-1 /dm3mol-ls-l /"C ref. 2,4-dinitrophenol quinuclidine 119000 20 25 7 triet h yl ami ne 23 500 17 20 8 t ri-n-propylamine 4 900 5.1 25 8 tri-n- butylamine 3 350 3.O 25 3 tri-n-pentylamine 5 700 3.3 20 8 t ri-n-octylamine 4 800 2.4 24 7 t ri-n-nonylamine 9 700 3.3 20 8 Magenta E trimethylamine 18800 18 13 4 triethylamine 119OOO 16 25 4 tri-n-propylamine 21 600 8 25 4 tri-n-butylamine 33 loo 5.2 24 4 2,4,6-trimethylpyridine 12 800 7 -26 4 Bromophenol Blue trimet hylamine 3 400 13 25 6 (second complex) triet hylamine 33000 12 25 6 t ri-n-propylamine 4900 5.6 25 6 tri-n-butylamine 7 200 5.5 25 6 tri-n-octylamine 6 300 2.3 25 6 If the second equilibrium is established much more quickly than the first the relaxation times are given by Since -cobs is a function of reagent concentration eqn (8) cannot apply.Eqn (9) is compatible with the experimental data if as is usually found in practice K23 B 1. Thus k may be identified with kI2,the rate constant for formation of the hydrogen- bonded intermediate which accounts for its independence of K and its variation by steric effects.The experimentally observed variation of k with solvent and with temperature shows that k12is not a simple diffusion-controlled rate. For a simple diffusion- controlled reaction the observed activation energy should be due solely to changes in solvent viscosity with temperature and so be equal to the activation energy of s 10-2 34 HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS viscosity of the solvent. Not only is this not so but for one reaction a negative value of AH has been observed as shown in table 3. On the simplest model the rate of a diffusion controlled reaction is given by the Smoluchowski equation l2 kD = 44Dm+DB)(rAH+rB) (10) where rAHis the radius of AH,considered as a spherical molecule and DAH is the diffusion coefficient of AH.Since D,H is related to the solvent viscosity by the Stokes-Einstein equation DAII = kT/4nqrAH k = q/(2+rBr,&+rAHr;l) kT fi17/4kT. Thus a plot of k;' (in units molecules m-3 s against (in units kg m-ls-') has gradient (4kT)-l (in units J K). If k and q are in the more conventional units of mol d~n-~ s and Poise respectively the gradient is (6.02 x x 4kT)-' mol dm-3 s Poise-' which has the numerical value of I .09 x lo-* at 25°C. The observed values of this gradient are much greater as shown in table 4 for solvents ranging in viscosity from chlorobutane (q298 = 4.27 x Poise) to iodobenzene = 15.8 x Poise). Furthermore the plots of k;' against q have a positive intercept on the k;l axis. TABLE 3.-ACTIVATION ENERGIES FOR ION-PAIR COMPLEXES FORMATION acid base AH^* AH& solvent /kJ rnol-1 /kJ mol-1 ref.2,4-dinitrophenol tri-n-butylamine chlorobenzene -3.5 +8.8 10 Magenta E trimet hylamine chlorobenzene +3 +8.8 4 Magenta E triethylamine chlorobenzene +9 +8.8 4 Magenta E tri-n-propylamine chlorobenzene +12 +8.8 4 Magenta E tri-n-but ylamine chlorobenzene $6 +8.8 4 Magenta E 2,4,6-trimethylpyridine chlorobenzene +2 +8.8 4 Magenta E tri-n-butylamine chlorobutane +5 +7.5 11 Magenta E tri-n- bu tylami ne chloropentane +13 +8.4 11 Magenta E tri-n-butylamine bromobenzene +8 +10.5 11 These anomalies can be resolved if the formation of the hydrogen-bonded complex C is broken down into two kinetically distinct steps. The first step is the translational diffusion together of reagent molecules to form an encounter complex AH B in which AH and B are oriented at random.The second step is a rotational diffusion so that AH and B rotate until they are correctly oriented for the formation of a hydrogen bond. The hydrogen-bonded complex AH.. . B,then reacts at rate k to give the ion-pair product. kt kr AH+B+AH,B+AH..B-+A-..HB+. k; trans-rotation reaction lational diffusion A detailed analysis of this process has been given by Solc and St~ckmayer,'~ J. E. CROOKS AND B. H. ROBINSON k is the rate of reaction which would be observed if the translational and rotational diffusive processes were infinitely fast and so is a bimolecular rate constant in eqn (1 3). It may be expressed as a unimolecular rate constant the rate for the unimolecular conversion of hydrogen-bonded complex to ion-pair i.e.k23,by dividing by the equilibrium constant for hydrogen-bonded complex formation K,2. The other parameters are defined as below a = (rAH +rB) 4AH = fraction of surface area of AH considered as a sphere available for reaction AAH = ($AH +k;z~H)l(l +kl~*H) zAH = correlation time for rotational diffusion ~-' = (1 -AAH)-'(l -AB)-' +(1 -AAH)-'(AB-&)-' +(1 -AB)-'(AAH-+AH)-' Values of rAH and rB may be estimated from space-filling molecular models (Catalin). The value of k may be taken as k K,& where k is identified with kDand evaluated from eqn (10) and KAHB the equilibrium constant for encounter complex formation may be evaluated from l2 KAHB = ha3.(14) (This is valid for dissimilar YAH and rB if rAH = rBr a value of 4na3is preferred for KAHB). TABLE 4.-PARAMETERS FOR THE SOLC-STOCKMAYER CALCULATIONS rm mI reaction 2 Magenta E+ tributylamine 560 400 1.6 1.57 0.15 0.15 6.2 1.6 36 10 2,4-dinitrophenol+tributylamine 560 400 2.8 3.00 0.08 0.15 3.3 3 72 19 *Using values of #AH and +B tabulated. The value of zAH may be calculated l2 from the Debye equation ZAH = hqriH/kT. Hence Ic~zAH = &(rid + ri ')/a2. (16) The gradient of the plot of k< against q according to eqn (13) is inversely propor- tional to $AH& and insensitive to the other parameters so that it is more convenient to adjust 4AH and 4Bto obtain the experimental gradient and see if sensible values are obtained.Table 4 shows values of the gradient calculated from the values of 4AH and $B given which inspection of molecular models shows are not unreasonable. The high value of the gradient for 2,4dinitrophenol as acid may be seen to be attribut- able to a low value of the fractional surface area of the molecule available for reaction. HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS This may be due to the restriction of the rotation of the OH group by hydrogen bonding to the ortho NO1 group. The model implicitly ignores the possible efTect of the intramolecular hydrogen bond causing an activation energy for the formation of the hydrogen-bonded complex. Solvation effects are similarly ignored. It is not possible to use eqn (13) to derive a relationship between the observed activation energy and the activation energies for the final step and for viscosity.However the small or negative values for the observed activation energy may be explained in qualitative terms. The observed rate is a function not only of k but also of the concentrations of hydrogen-bonded complex and encounter complex. Increasing the temperature increases k, but decreases the concentrations of reaction intermediates since AH" for their formation is negative. Thus an increase in temper- ature may lead to an anomalously small or negative increase in the observed rate. ANOMALOUSLY SLOW REACTIONS INVOLVING THE FORMATION OF ION-PAIR COMPLEXES The indicator acid Broniophenol Blue 11 reacts with pyridine bases to form ion- pair complexes of type 111.0 Br HO Br 111 The kinetics of this reaction have been investigated using a laser tempera-ture-jump apparatus,l monitoring the progress of the reaction by spectro-photometric detection of 111which has an absorption peak at 405 nm. Relaxation times were found to be in the range 500 ps-25 ms whereas relaxation times for the systems listed in table 2 were in the region 2-5Op. The relaxation HO IV J. E. CROOKS AND B. H. ROBINSON times varied with concentration in accordance with eqn (5). The kinetic data obtained are listed in table 5. It can be seen that by contrast with the values listed in table 2 k is very dependent on base strength and well below the values expected for a diffusion-controlled reaction.Values of k are almost independent of base. Values of AH are negative for three of the four bases studied. TABLE 5.-THERMODYNAMIC AND KINETIC PARAMETERS FOR THE REACTION BETWEEN BROMOPHENOL BLUEAND PYRIDINE BASES IN CHLOROBENZENE SOI~UTION AT 298 K 10- 5 kr -AS" -Asf* amine 10-3 K /dm3rnol-1 /dm3 mol-' s-l kb Is-' -AH" /kJrnol-' -AH* /kJ mol-' /Jmol-' K-' /Jrnol-l K-' pyridine 2-methylpyridine 2,6-dimethylpyridine 2,4,6-trimethylpyridine 2.09 31.6 123 1 000 1.16 14.6 96.2 970 50 44 70 98 43 58 55 60 9.3 -5.3 -6.6 -15.4 81 108 87 88 117 145 133 144 The negative values of AH strongly suggest the kinetic significance of an inter- mediate complex which may be identified with the hydrogen-bonded complex IV.The kinetic scheme is thus kiz k23 AH+B$C+D k21 k3z in which C is the hydrogen-bonded complex IV D is the ion-pair complex 111 and the rate-determining step is the conversion of IV to 111. The formation of IV from Bromoplienol Blue and base is an exothermic process so that raising the temperature reduces the concentration of IV and hence reduces the rate of the overall reaction. The first equilibrium is rapidly established ; k, is of the order of the diffusion- controlled rate as shown for the analogous systems listed in table 2 for which kfhas been identified with k12. For the systems listed in table 5 the observed relaxation TABLE 6.-THERMODYNAMIC AND KINETIC PARAMETERS FOR THE INTERCONVERSION OF HYDROGEN-BONDED AND ION-PAIR COMPLEXES OF BROMOPHENOL BLUEIN CHLOROBENZE~E SOLUTION AT 298 K -AH53 AHA -AS53 -ASX 10-2 523 k3z !kJ /kJ /J K-1 /J K-l base K23 15-is-' mol-1 rnol-1 rnol-1 mol-1 3-chloropyridinea 1.2 2.6 220 14 34 44 83 pyridine 300 150 50 23 29 41 2-met hy I pyridi ne 4 200 1 900 44 37 16 55 92 2,6-dimethylpyridine 18 000 13 000 70 29 20 28 96 2,4,6-trimethylpyridine 92 000 89 000 98 37 8 29 85 Notes a from eqn (3) using 2,6-dibromophenol with base in CCI as model system for evaluation of KI2.time is then related to the individual rate constants by eqn (7). For all the bases studied except 3-chloropyridine K23 9 KI2,so that K,,([AH]+[B]) 4 1. There-fore k is identified with k23K12 k with k32,and AH with AH:2+AHG. Values of K, and AH:2 have been taken from the model systems listed in table 1 to give values for the kinetic parameters for the rate-determining step which are listed in table 6.A pronounced solvent effect on the kinetics has been observed." Values of kr and kb have been obtained by both stopped-flow and temperature-jump techniques HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS which were found to be concordant and are listed in table 7. The primary isotope effect has been measured by the differential stopped-flow technique,'* and found to be close to unity. Values of AV? have been measured l9 by use of a high-pressure observation cuvette incorporated in the laser temperature-jump apparatus.'' These data are listed in table 8. TABLE7.-sOLVENT EFFECTS ON THE KINETICS OF THE REACTION BETWEEN BROMOPHENOL BLUEAND PYRIDlNE AT 298 K solvent relative permittivity ET /kJrnol-' 10-3~-10-39 /drn3 mol 1 /dm3 mol- s-l 1O-'?3 Is- a kb(= F2) Is- C6H5CH3 2.38 141.7 1.15 6.9 3.4 6.0 C6H6 2.27 144.2 1.82 13.8 5.6 7.5 C~HSC~ 5.62 156.8 2.12 117 40 55 CHZCI 8.89 171.8 6.2gb 641 184 71 b Notes a kz3 = kf/K, ; values of K, taken as those for the hydrogen-bonded complex between Magenta E and pyridine (see table 1).b K # kf/kbin this instance because of the presence of a further equilibrium involving dimerisation. TABLEPRIMARY ISOTOPE EFFECTS AND VOLUME DATA FOR THE REACTION BETWEEN BROMOPHENOL BLUEAND PYRIDINE AT 298 K -AV,* -AVO -AV$(= -At'&) solvent Ka/Kn (kf)E/(kf)D /cm3 mol -1 /cm3 mol-1 /cm3 mol -1 CrjHsCHs 1.04 1.04(+0.01) --C6HsCl -16 14 2 The formation of 111 differs from most other reactions in which proton-transfer results in ion-pair complex formation in that the negative charge on the deprotonated acid is located at some distance from the site of proton loss.The sultone ring in I1 has opened to form the new site of the negative charge. Five processes can be recognised as occurring along the reaction co-ordinate namely (i) formation of the hydrogen-bonded complex IV (ii) proton transfer along the hydrogen bond (iii) solvent re-organisation associated with (ii) (iv) opening of the sultone ring during which the negative charge disappears from the phenolate group and appears on the sulphonate group (v) migration of the protonated amine from the phenolate to the sulphonate site.It is clear from the data in table 7 that the solvent plays a dominant role in the kinetics of the reaction. The values of AS:3 and AVT suggest that there is extensive solvent reorganisation on forming the transition state for the reaction of IV to give 111. One reason why this reaction is slow in such non-polar solvents as chlorobenzene is the need for solvent involvement. This may be because proton transfer can only occur when the solvent has adopted a particular configuration that appropriate for the solvation of the ion-pair. The rate kZ3,will thus depend on the probability of this rearrangement occurring. The rate increases rapidly with increasing polarity of the solvent and there is a good correlation between log kz3and the ET polarity value of the solvent.21 Extrapolation suggests a value of 10l2 s-I for the rate in aqueous solution.The large value of -A V suggests considerable electrostriction of the solvent in the transition state. Furthermore the value of AV& suggests that the solvation of the transition state resembles that of the ion-pair complex. The negative J. E. CROOKS AND B. H. ROBINSON value of AS2 suggests that the transition state is more ordered than the ion-pair complex and as this is not due to solvation changes this may be associated with the existence of the sultone ring in the transition state. The closeness of the primary isotope effect to unity contrasts with the large effects recently observed for proton transfer from carbon acids to amine bases in similar solvents.22 This indicates that for Bromophenol Blue as acid either proton migration in the rate-limiting step is strongly coupled with other heavy atom motion (e.g.solvent rearrangement or sultone ring opening) or proton-transfer is not rate-limiting. If the latter is valid the actual proton migration step can only affect the observed rate via a pre-equilibrium constant. There are two plausible transition states HO V VI In V proton-transfer is synchronous with ring-opening whereas in VI these two processes are uncoupled. We believe that the weight of the evidence favours VI. VI is more ordered than 111 but solvated to a similar extent as required by the evidence from AS; and AVG. The kinetic isotope effect is close to unity because proton- transfer is almost complete in the transition state VI.The rate of the back reaction k32,is little affected by base strength because proton-transfer has hardly started in the transition state for the reverse process. The proton-transfer step is slow i.e. k23is low because proton transfer is opposed by both an entropy and an enthalpy of reaction. The negative values of AS2 are due to the need for solvent reorganisation around the highly dipolar transition state. The comparatively high values of AH2 are due to the need to supply energy for charge generation and separation which in the transition state is not adequately supplied by delocalisation energy. The dependence of AH2 on the base strength shows the effect of electron delocalisation in the amine.Aromatic amines are weak bases because delocalisation energy is lost on protonation but this loss is reduced by electron donation from substituent methyl groups. Proton transfer in the systems listed in table 2 is fast because delocalisation is synchronous with charge generation so that the energy of the system decreases continuously as the proton moves across the hydrogen bond. A. A. Maryott J. Nat. Bur. Stand. 1948 41 1. R. P. Bell and J. E. Crooks J. Chem. SOC.,1962 3513. J. E.Crooks and B.H. Robinson Chem. Commun. 1970,979. E. F. Caldin J. E. Crooks and D. O'Donnell J.C.S. Faraday I 1973 69 1000. J. E. Crooks and B. H. Robinson Trans. Faraday SOC.,1971 67 1707. J. E. Crooks P. J. Sheridan and D. O'Donnell J. Chem.SOC.B 1970 1285. 'E. F. Caldin J. E. Crooks and D. O'Donnell J.C.S. Faraday I 1973 69 993. HYDROGEN-BONDED AND ION-PAIR COMPLEXES IN APROTIC SOLVENTS * K. J. Ivin J. J. McGarvey E. L. Simmons and R. Small J.C.S. Faraday I 1973 69 1016. E. F. Caldin and J. E. Crooks J. Chem. Soc. B 1967 959. lo K.J. Ivin J. J. McGarvey E. L. Simmons and R. Small Trans Faraday Soc. 1971 67 104. l1 G. D.Burfoot E. F. Caldin and H. Goodman J.C.S. Faraday I 1974,70 105. l2 A. M. North The Collision Theory of Chemical Reactions in Liquids (Methuen 1964). l3 K. Solc and W. H. Stockmayer Int. J. Chem. Kinetics 1973 5 733. l4 B. H. Robinson unpublished data. l5 E. F.Caldin J. E. Crooks and B. H. Robinson J. Phys. E 1971 4 165. l6 J. E. Crooks and B.H. Robinson Trans. Faraday SOC.,1970 66 1436. l7 G. Gammons B.H. Robinson and M. J. Stern J.C.S. Chem. Commun. 1972 1157. l8 K. J. A.Hargreaves and B. H. Robinson unpublished data. l9 T. Altinata B.H. Robinson and C. J. Wilson unpublished data. 2o E. F. Caldin M. W. Grant B. B. Hasinoff and P. A. Tregloan J. Phys. E 1973 6 349. 21 K. Dimroth C.Reichardt T. Stepmann and F. Bohlmann Ann. 1963 661. 22 E. F. Caldin and S. Mateo J.C.S. Chem. Commun. 1973 854.

ISSN:0301-5696    

DOI:10.1039/FS9751000029

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Structure and Mechanism in Complex General Acid-base Catalyzed Reactions BY WILLIAM AND JANEM. SAYER P. JENCKS Department of Biochemistry Brandeis University Waltham Massachusetts 02154 U.S.A. Received 28th April 1975 The mechanisms of acid and base catalyzed carbonyl and acyl group reactions are determined largely by the lifetime of the initial addition intermediate. Catalysis through diffusion-controlled trapping by relatively strong acids or bases niust occur if reversion of the intermediate to reactants is faster than proton transfer involving solvent and product formation. If breakdown to reactants is faster than separation of the intermediate and catalyst the intermediate is formed within a solvent cage containing the catalyst through a pre-association mechanism.If the " intermediate " is still more unstable or if there is no barrier for proton transfer a stepwise reaction is impossible and the reaction must be concerted. Thus changes in structure of the reactants can be correlated with changes in the lifetime of the intermediate and the mechanism of catalysis. We would like to know the nature and the magnitude of the rate enhancements that are brought about by general acid and base catalysis of carbonyl and acyl group reactions a field in which R. P. Bell has been a pioneer.l* A few examples are described here in which changes in the structure of reactants and catalysts and the lifetimes of intermediates provide some insight into the mechanism and driving force of these reactions. The mechanisms are divided into two broad classes and several subclasses although there is not always a sharp dividing line between them.I. TRAPPING OF ADDITION INTERMEDIATES A. STABLE INTERMEDIATES When a strong nucleophile attacks a reactive carbonyl compound the initial product is likely to have enough stability to be trapped by proton transfer involving the solvent before it reverts to reactants so that trapping by added catalysts is unnecessary. For example the pK of the oxygen atom in the addition compound that is formed from the attack of trimethylamine on formaldehyde is 9.3 so that the rate constant k for protonation of T* by water is 4 x lo5 s-l (eqn (1); T' based on this pK and diffusion-controlled deprotonation by hydroxide ion in the reverse direction).Since the rate constant kl for amine expulsion to form the unstable formaldehyde molecule is only 3.4 x lo3 s-l (eqn (1)) the addition compound T"will always go on to products and added buffers do not catalyze the rea~tion.~' The more basic addition compounds formed from anionic nucleophiles generally have pK values of the order of 11-13 and will abstract a proton from the solvent 41 GENERAL ACID-BASE CATALYZED REACTIONS 102-103faster; they are correspondingly less likely to exhibit buffer catalysis. Thus cyanide hydroperoxide hydroxide bisulphite and basic thiol anions add to carbonyl compounds without buffer catalysis. 0 0-HO kl t I ks I HX + ,C I1 =HX-C-X-C-(2) k-I I k- I If there is a proton on the attacking nucleophile the intermediate can be trapped by a proton switch mechanism with a rate constant k of ca.106-108 s-l (eqn (2) see later) so that if expulsion of the attacking nucleophile is slower than this no buffer catalysis by trapping will be seen. B. DIFFUSION-CONTROLLED PROTON TRANSFER When a less stable intermediate breaks down to starting materials faster than it is trapped by proton transfer involving solvent there must be catalysis by buffers. When methoxyamine a less basic amine attacks p-chlorobenzaldehyde a less reactive aldehyde expulsion of the attacking amine (k.-l = 3 x lo8s-I) is faster than either proton abstraction from water (k2 = 3 x lo4s-l) or a proton switch (k,= 6 x lo6s-l) so that trapping of the intermediate by encounter with a moderately strong buffer acid increases the observed rate (eqn (3)).5 When the acid strength is too ti MeONH2 + >C=O kl +I cA-lMeONH2-C-OH -2 MeONH2-C-O-kACHA’ ~ k-1~3 x 108s-1 I k-A 1 (3) T’ weak to give a thermodynamically favourable proton transfer trapping will not occur on every encounter and the Bronsted plot becomes an “ Eigen curve ” with a change in slope from limiting values of 0 to -1 with increasing pK of the acid.When the catalyst concentration is increased sufficiently trapping becomes faster than reversion to reactants methoxyamine attack becomes rate determining and the plot of kobsagainst buffer concentration levels off as the rate becomes independent of buffer concentration. Trapping of T by diffusion-controlled reaction with hydroxide ion or proton transfer to a second molecule of attacking amine occurs in the reactions of piperazine and sarcosine with ~yridine-4-aldehyde.~.Similarly the synthesis and cleavage of ureas RNHCONH, shows no buffer catalysis when RNHz is basic but the reaction of 4-anisidine with cyanic acid shows 0 k + RNH + C II RNH2-I1 k-1 N H W. P. JENCKS AND J. M. SAYER general acid and base catalysis with nonlinear Bronsted plots and similar maximum rate constants for strong acids and bases.' With basic amines the intermediate T' is relatively stable (k-l < k, eqn (4)) but as the amine becomes less basic the expulsion of amine from T* becomes faster than proton transfer to form urea ( k- = 3 x 10' s-l k = 10' s-l) and the intermediate can be trapped by added catalysts through k and k,.The reaction with anisidine and the cleavage of 1-phenylcarbamoylimidazole exhibit a change in rate-determining step from proton transfer to C-N bond formation or cleavage as the buffer concentration is increased.8* Addition intermediates formed from the attack of nucleophiles on acyl compounds are usually less stable than those formed from aldehydes and ketones because of electron donation by resonance from the 0 N or S atom that helps to expel the nucleophile and stabilize the acyl compound. The value of k- for expulsion of the aliphatic amino group from T* in the intramolecular aminolysis of S-acetylmer- captoethylamine is 6.6 x 10' s-' and the intermediate is trapped by encounter with acids giving a nonlinear " Eigen-type " Bronsted plot (eqn (5)).lo T' A separate proton transfer step is also required by the dependence on acid concen- tration of the aminolysis rate and the product distribution from the hydrolysis of 2-methyl-A2-thiazoline which proceeds through the same intermediates.lo*' The estimated value of k- for the addition compound T* formed from phenyl acetate and methylamine is 3 x lo9 s-' and the available data are consistent with a mechanism for ester aminolysis in which general acid and base catalysis involves trapping of this intermediate by proton transfer (eqn (6)).12 The addition inter- mediate T' can also be generated upon hydrolysis of the corresponding imido ester. T+ 0 0- ZNH t II COR 'H I-N-C-OR I I T' kf I+ TO Amide - T- Amide The partitioning of this intermediate between ester and amide is independent of the pK of the leaving phenolate ion and shows nonlinear Bronsted plots for general acid catalysis as expected if the partitioning is controlled by proton transfer.I2 The rate constants of aminolysis reactions characteristically exhibit a large dependence on the GENERAL ACID-BASE CATALYZED REACTIONS basicity of the amine (Pnuc= 0.8-1.0) when the rate-determining step is a reaction of T* involving (a) encounter-controlled catalysis by general acids or bases (b) a proton switch to give TO;or (c) direct breakdown by expulsion of a relatively good leaving group via k*.c. PRE-ASSOCIATION OR " SPECTATOR " MECHANISMS When the rate constant for the breakdown of an intermediate becomes larger than that for diffusion apart of the intermediate and catalyst (e.g.kLi and k- H kt 'H I HN -t ;C-0 HN-C-0' R k-I R I -Products H k'l +H I HI BH+.N-C-O'-B.HN.>C-O B.HN-C-0-w-R kLt RI RI respectively in eqn (7)) the pathway of lowest free energy for the breakdown and formation of the intermediate must involve a preliminary association of the reactants and catalyst in an encounter complex; the catalyst may be present as a " spectator " during heavy atom rearrangement l4 When the catalyst does not provide appreciable stabilization of the transition state for formation of the intermediate the Bronsted a or p value for strong acids or bases will be zero but for weak acids or bases the subsequent proton transfer and separation of product and catalyst will become rate determining so that the Bronsted plot exhibits a break similar to that in simple diffusion-controlled proton transfer reactions.Since the pre-association mechanism provides a lower energy faster pathway than the diffusion-controlled mechanism the larger rate constants for strong acid or base catalysts will cause the break in the Bronsted plot to be shifted as shown in the upper line of fig. 1.15*l6 PKBH + FIG.I.-Bronsted plot illustrating that the break in the curve for a base-catalyzed pre-association mechanism is at a higher pK than that for a diffusioncoiitrolled proton transfer mechanism. W. P. JENCKS AND J. M. SAYER This mechanism has been suggested for general base catalysis of the attack of 2-methylthiosemicarbazide on p-chlorobenzaldehyde which shows a break at least 1.4 pK units above the expected pK of 3.1 for T* and gives a calculated Bronsted plot in good agreement with experiment based on a value of kL = 5 x loll s-l.Such a shift can provide presumptive evidence for a preassociation mechanism; the absence of a dependence of the rate on solvent viscosity l7 would provide further evidence. IT. CATALYSIS WITH TRANSITION STATE STABILIZATION A. HYDROGEN BONDING Structure-reactivity considerations predict that general acid catalysis of the attack of a nucleophile on the carbonyl group will involve more proton transfer (larger Bronsted a) for a more weakly basic nucleophile.l** l9 In accord with this the Bronsted a value of 1.0 for the formation and breakdown of hydrogen peroxide- p-chlorobenzaldehyde addition compounds is much larger than the a values for aldehyde and ketone hydration.2o Although the interpretation of Bronsted co- efficients has been questioned,21 it is an experimental fact that this a value of 1.0 means that polar substituents on the catalyst have the same effect on the stability of the transition state as they do on the equilibrium ionization reaction so that we can say that so far as substituent effects are concerned the catalyst in the transition state resembles the conjugate base of the acid. The fact that acid-catalyzed breakdown is 6 to 18 times faster for the p-methoxybenzaldehyde than for the p-chlorobenzaldehyde adduct means that significant C-0 bond cleavage has occurred in the transition state since even complete protonation should be favoured by only about two-fold in the p-methoxy compound.These data suggest that the rate-determining step is the formation and cleavage of the C-0 bond and that the essentially complete proton transfer in the transition state is stabilized by hydrogen bonding to the catalyst (eqn (8)). The rate constant for the solvated proton deviates downward by approx- imately 50-fold from the Bronsted line for carboxylic phosphoric and arsenic acids. This makes the detection of general acid catalysis possible in spite of the a value of 1.0 and suggests that bifunctional hydrogen bonding to the two acidic protons in the transition state may account for the relatively high activity of the latter catalysts.B. CONCERTED CATALYSIS The strongest evidence for concerted acid-base catalysis is for reversible additions of ROH to unsaturated centres with Bronsted p values (for ROH addition) and a values (for ROH expulsion) near 0.5. In the general acid catalyzed breakdown of an alcohol-phthalimidium addition compound (eqn (9)) the value of a increases steadily R R k GENERAL ACID-BASE CATALYZED REACTIONS from 0.49 to 0.74 as the leaving alcohol becomes more basic in the series from tri- fluoroethanol to As the catalyzing acid becomes stronger there is a decrease and then a reversal in the dependence of the rate on the pK of the leaving alcohol (Qlg changes from -0.23 for acetic acid to +0.24 for the proton).These results are inconsistent with a stepwise mechanism of catalysis and also cannot be explained by a hydrogen-bonding mechanism in which the proton rests in one well of a double potential well hydrogen bond during C-0 bond formation and break- down.22*23 If the alcohol leaves as the anion with hydrogen bonding to the catalyst there must be negative charge on the oxygen atom in the transition state (1). This is inconsistent with the faster acid catalyzed expulsion of ethanol than of trifluoroethanol. I 2 If the alcohol is first protonated and leaves with hydrogen bonding to the buffer base there must be positive charge on oxygen in the transition state (2). This is inconsistent with the faster expulsion of trifluoroethanol than of ethanol catalyzed by the weaker acids.The linear Bronsted plots and the gradual change in a and in sensitivity to leaving group pK support a concerted mechanism rather than a sudden shift from one to the other of the hydrogen-bonding mechanisms 1 and 2. The data may be described by reaction coordinate contour diagrams such as that shown in fig. 2gfor the 'acetic acid-trifluoroethanol reaction. Changes in the position +I 1 A-HO-C-AH O-C-RI RI FIG.2.-Possible reactioncoordinate contour diagram for the breakdown of a trifluoroethanol- phthalimidium ion addition compound catalyzed by acetic acid. The horizontal axis represents proton transfer and the vertical axis represents cleavage and formation of the C-0 bond. of the transition state with changing structure of the reactants and catalysts provide a rationalization for the observed changes in a and leaving-group effects.These changes correspond to an interaction coefficient l/c = 0.07 in the Cordes equation that interrelates these parameters. 22 The analogous general acid catalyzed hydrolysis W. P. JENCKS AND J. M. SAYER of benzaldehyde phenyl acetals exhibits an even larger sensitivity of changes in transition state structure to changing structure of the catalyst and leaving with l/c z 0.2. Evidently the energy gradients in the contour diagrams that permit these large shifts are unusually shallow. These shifts and the normal or low catalytic constants for the proton in these reactions appear inconsistent with an S-shaped reaction coordinate with only vertical motion for C-0 cleavage and the proton in a stable potential well in the transition state.The general acid catalyzed dehydration of carbinolamines to imines (eqn (10)) is a similar reaction in which the value of a increases from 0.62 to 0.73 as the basicity of the amine decreases in the series hydrazine (pK 8.3) to thiosemicarbazide (pK 1.9).25 This change is as expected from the increase in the energy of the lower part of the reaction coordinate diagram and corresponds to an interaction coefficient l/c2 = 0.02. The addition of methoxyamine to p-chlorobenzaldehyde is catalyzed by the proton through a class I1 (concerted or hydrogen bonded) mechanism (eqn (1 1)) concurrently with the previously mentioned class I stepwise mechani~m.~ As the pK of the attacking amine is decreased the dipolar addition intermediate becomes less stable the stepwise mechanism is correspondingly less favourable and only the class I1 mechanism is observed for the reaction with 2-methylthiosemicarbazide.Electron-donating substituents on the aldehyde also favour amine expulsion and destabilize the dipolar intermediate relative to the aldehyde and to the transition state of eqn (1 1). Accordingly the addition of semicarbazide to p-nitrobenzaldehyde proceeds through concurrent class I and class I1 mechanisms whereas only the class I1 mechanism is observed for p-methoxybenzaldehyde.26 Thus "harder " reactions that require more unstable intermediates are more likely to proceed through class I1 mechanisms. c.DIFFUSION-CONTROL LED CATA L Y Z ED RE ACT I0 N S When structural changes make an intermediate progressively less stable a point may be reached at which the reaction occurs with a strong catalyst at a diffusion- controlled rate but with a weak catalyst at a slower rate through a class I1 mechanism. This will result in a break in the Bronsted plot and other structure-reactivity correla- tions. For example strongly basic thiol anions add to acetaldehyde with rate- determining nucleophilic attack (kl eqn (12)) no buffer catalysis and essentially no dependence of the rate on nucleophile basicity but weakly basic anions exhibit a I k2 I RS-+ ;c=o RSCOH.A-RSCOH 4-A-I k-2 1 sharp downward deviation in the structure-reactivity correlation and a value of Pnuc = l.0.27928 This is because the rate-determining step becomes the separation of hydroxide ion from the product (k2,eqn (12) A-= OH-); in the reverse direction the diffusion-controlled encounter of hydroxide ion with the addition compound is rate determining (k20.9 x 1O1O M-1 s-l) .When the thiol is weakly basic the = intermediate T-is unstable so that thiol anion expulsion is faster than protonation by water; the estimated rate constants for methyl mercaptoacetate (pK = 7.8) are 48 GENERAL ACID-BASE CATALYZED REACTIONS k- = 5 x lo8 s-' and K,k2 = 2.5 x lo8s-'." For still less basic thiols k- is estimated to be in the range 1010-1012.4 s-' and a general acid catalyst must be present in the reacting complex for the reasons given in I-C.Accordingly these reactions exhibit general acid catalysis with a Bronsted a value of 0.2 and a small dependence of the rate on thiol anion basicity (Pnuc = 0.15 for two thiophenols). This is consistent with either hydrogen bonding or a concerted mechanism of catalysis. The position of the break in the Bronsted plot for this and other reactions of this kind does not correspond to the pK of the intermediate. The breaks in the nonlinear Bronsted plots for general base catalysis of the hydrazinolysis of acetylimidazole and for general acid and general base catalysis of the methoxyaminolysis of acetyl- triazole 29 are more than 2 pK units away from the estimated pK values of the addition intermediates. This is consistent with the mechanism of eqn (13) for these reactions 0 0-0 11 kl + I kd[Catj kc n RNH 4-,C-X RNH,-C-X L[T'.Cot] ,CNHR t Cat (I3) k-I I k-d in which strongly basic or acidic catalysts react with the intermediate T' at every encounter (k,) and weak catalysts cause a slower concerted breakdown of Tk (kc); the rate constants for strong acid and base catalysts in the acetyltriazole reaction are equal corresponding to a common value of kd.Mechanisms for acid catalyzed addition of a nucleophile N to a carbonyl group are summarized schematically in fig. 3. The preferred mechanism depends largely on the lifetimes of the intermediates. It is obvious (although not always recognized) that if an " intermediate " has too short a lifetime to exist (<10-'3-10-'4 s) a reaction cannot be stepwise and must be concerted.What is not yet known is whether concerted catalysis is possible when the intermediate does exist. The concurrent Type I and Type I1 reactions of methoxyamine and semicarbazide with benzaldehydes demonstrate that the intermediate T' has a significant lifetime (ca. s) with a barrier for C-N cleavage and that concurrent reaction pathways I I FIG.3.-Mechanisms for general acid catalysis of the addition of a nucleophile N to a carbonyl group. The mechanisms are drawn in the form of a reaction coordinate diagram but contour lines are omitted. The dashed lines represent free energy barriers for association and diffusion processes. * For HA = H20 based on pKa = 12.4 for To,from a measured pKa of 12.4for HOEtSCH20H (R. Kallen personal communication) and structure-reactivity correlation^,^^ and an assumed value of k-2 = 10'' M-' s-'.Note that if the breakdown is concerted the K; step is not required. W. P. JENCKS AND J. M. SAYER 49 inside and outside the central box of fig. 3 are possible. If there is no barrier for proton transfer in the T'*HA complex this intermediate does not exist and the reaction must be concerted. However we do not know whether there is such a barrier nor whether coiicerted catalysis is possible when the existence of this barrier provides a potential well for the complex. The mechanism is likely to be stepwise when the proton transfer occurs through one or more water molecules as it does in many simple proton transfer reactions 30 whereas a concerted mechanism is favoured when there is direct proton transfer between the reactant and catalyst and an initially unfavourable proton transfer suddenly becomes strongly favourable in the course of the reaction.R. P. Bell Acid-Base Curalysis (Oxford Univ. Press London 1941). R. P. Bell The Proton in Chemistry (Cornell Univ. Press Ithaca N.Y. 2nd ed. 1973). T. D. Stewart and H. P. Kung J. Amer. Chem. SOC., 1933 55 4813. J. Hine and F. C. Kokesh J. Anrer. Chenr. SOC. 1970 92 4383. S. Rosenberg S. M. Silver J. M. Sayer and W. P. Jencks J. Anier. Chern. SOC.,1974,96 7956. H. Diebler and R. N. F. Thorneley J. Amer. Chenz. SOC.,1973 95 996. R. N. F. Thorneley and H. Diebler J. Anrcr. Clrern. Soc. 1974. 96 1072. A. Williams and W. P. Jencks J. C. S. Perkin II 1974 1753 1760.A. F. Hegarty C. N. Hegarty and F. L. Scott J. C. S.,Perkin I/ 1974 1258. lo R. E. Barnett and W. P. Jencks J. Amer. Cliem. SOC., 1969 91 2358. R. B. Martin and R. I. Hedrick J. Anzer. Chetn. SOC.,1962 84 106; R. B. Martin R. I. Hcdrick and A. Parcell J. Org. Cliern. 1964 29 3197. A. C. Sattcrthwait and W. P. Jencks J. Atner. Cficm. SOC. 1974 96 7018 7031. l3 W. P. Jencks and K. Salvesen J. Amer. Cfietn. SOC. 1971 93 1419. L. D. Kershner and R. L. Schowen J. Amer. Clienr. SOC.,1971 93 2014. l5 M. I. Page and W. P. Jencks J. Amer. Chem. SOC. 1972 94 8828. J. kl. Sayer and W. P. Jencks J. Amer. Chenr. SOC.,1973 95 5637. l7 C. Cerjan and R. E. Barnett J. Phys. CJietn. 1972 76 1192. E. H. Cordes and W. P. Jencks J. Amer. Cfienr. Soc. 1962 84 4319.l9 W. P. Jencks Chem. Rev. 1972 72 705. 2o E. Sander and W. P. Jencks J. Anrer. Chern. SOC., 1968 90 4377. 21 F. G. Bordwell and W. J. Boyle Jr. J. Anier. Chetn. SOC.,1972 94 3907; A. J. Kresge Curiud. J. Clienr. 1974 52 1897. 22 N. Gravitz and W. P. Jencks J. Amer. CJiern. SOC. 1974 96 507. 23 J. Hine J. Amer. Cliern. Soc. 1972 94 5766. '' B. Capon personal communication. 2.i J. M. Sayer M. Peskin and W. P. Jencks J. Anrer. Cfreni. SOC. 1973 95 4277. 26 J. M. Sayer B. Pinsky A. Schonbrunn and W. Washstien J. Amer. Clrem. SOC.,1974,96 7998. 27 G. E. Licnhard and W. P. Jencks J. Anier. Chenr. SOC. 1966 88 3982. 2R R. E. Barnett and W. P. Jcncks J. Ainer. Chenr. SOC., 1969 91 6758. 2y J. P. Fox and W. P. Jencks J. Anrer. Chem. Soc. 1974 96 1436. 30 D. Roscnthal and E. Grunwald. J. Anrer. Chenr. SOC.,1972 94. 5956 and references therein.

ISSN:0301-5696    

DOI:10.1039/FS9751000041

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

GENERAL DISCUSSION Dr. J. R. Jones (University of Surrey) (communicated) The reference to the decomposition of nitramide a reaction that has contributed greatly to the development of modern theories of acid-base catalysis prompts me to report some isotope effect data which I believe to be the first example of an isotope effect maximum for a nitrogen acid. The results in the table (kHrefers to ordinary nitramide in H20and in the presence of phenolate bases kD to deuteriated nitramide in D,O and also phenolate bases) represent the product of a solvent isotope effect (which probably lies between 2 and 3) and a primary isotope effect. The presence of the former does not however alter the fact that the primary isotope effect is at a maximum for catalysis by 2.4 6-trichlorophenolate and pentachlorophenolate anions and as the pK of nitramide is 6.48 this behaviour is similar to that observed for carbon acids.TABLE 1 .-ISOTOPE EFFECTS IN THE BASE-CATALYSED DECOMPOSITION OF NITRAMIDE AT 198.2 K base pK kH/1 mol-1 s-1 kHlkD water -1.74 8.5x 5.1 2,4-dini trophenol nitramide ion 4.09 6.48 0.0166 0.0368 3.5 5.4 pent achlorophenol 2,4,6-trichlorophenol2-nitrophenol 5.25 6.00 7.17 0.0320 0.449 1.84 10.2 9.0 6.3 This reaction also affords the unusual opportunity of studying catalysis by the anion of the substrate-that the isotope effect is lower than expected is probably related to the fact that the anion in aqueous solution is thought to possess the aci-form structure of nitramide. It is significant that the catalytic coefficient for the nitramide anion is much lower than that predicted from the Bronsted relationship established by phenolate bases.Prof. W. H. Saunders (University of Rochester) said How do the heats of solution of the neutral acids compare to those of the anions ? If the former are small compared to the latter it is easy to see why variations would be essentially negligible for the present purposes. If not changes in the energy of cavity formation with molecular size should be a significant factor. Prof. E. M. Arnett (University of Pittsburgh) said Because of electrostatic solvation the heats of solution of all the anions are much larger than those of the neutral acids. However in terms of the results presented in in this paper we are concerned with the relative effect of structural change on the anions compared to cyclopentadienyl anion and of the neutral acids compared to cyclopentadiene.In these terms again the calculated heats of solution of the anions are more sensitive to structure than are their neutral precursors-the range for the organic anions being about thirty kcal/mol and for their acids about six kcal/mol. The role of cavity size is indeed demonstrated in fig. 4 of our paper where a rough relationship between solvation energy and a crudely estimated reciprocal ionic radius is shown. I don’t believe that it is appropriate to interpret this correlation in terms of any particular solvation model since molar 50 GENERAL DISCUSSION volume terms play a significant role in the regular theory of solutions for non- electrolytes as well as in most electrostatic treatments of ions.The Born theory of ionic solvation does not work very precisely even for small spherical ions such as sodium or chloride. Its application to large complicated organic ions through correlations such as fig. 4,therefore represents the broad applicability of cavity terms which contribute both to solvation of the acid and its anion. Prof. El. P. Bell (University of Stirling) said The use of relative heats of solution in carbon tetrachloride in place of heats of vaporization is a reasonable procedure but I would like to ask whether this has been checked by using results for other non-polar solvents or by applying Trouton’s rule to the boiling points of the solutes.Prof. E. M. Arnett (University of Pittsburgh) said Yes indeed we have. A number of other non-polar solvents such as benzene and dichlorobenzene have been tried in our laboratory showing the same type of cancellation. A number of comparable measurements reported by Fuchs Drago Friedman and their students demonstrate the same type of cancellation. We admit to its vulnerability and believe that it will affect some of the results reported here but not very seriously. Prof. F. G. Bordwell (Northwestern University) said I would like to amplify and clarify the statement made in Arnett’s paper regarding “ acids. ..for which com- parable . . . pKs cannot be easily measured. ” The method for measuring pK’s outlined in our paper is applicable to most acids for which enthalpies can be deter- mined in DMSO.The one exception is alcohols. Here the formation of strongly hydrogen-bonded species e.g. RO- . . . H-OR causes interference. It is for such compounds that measurement of heats of deprotonation can provide supplemeiitary pK data. Prof. E. M. Arnett (University of Pittsburgh) said Bordwell is basically correct about his titration method with regard to which our claim sounds somewhat pejorative. When I made that statement I had in mind not only the problem of the alcohols but also the fact that all free energy measurements covering a broad spectrum of acidity require extrapolations with cumulative errors from compound to compound or from solvent to solvent. The particular value of our calorimetric procedure is that all acids both strong and weak are compared in exactly the same medium by the same method under the same conditions.Bordwell’s titration method is simple precise and elegant. Fig. 1 of our paper demonstrates the close parallel between his results and ours. Prof. E. F. Caldin (University of Kent) said For very fast proton-transfers in aprotic solvents such as those of substituted phenols considered by Crooks and Robinson the question is whether the rates and activation parameters can be explained by viscosity control of encounter or of rotation within the initial complex or whether we must consider also “chemical ” interactions such as desolvation breaking of internal hydrogen bonds and stabilisation of the initial complex by dispersion and dipole-dipole forces.When the experimental results are compared with those calculated by the simple Smoluchowski theory in which the reactant molecules are treated as spheres with uniformly reactive surfaces and translational diffusion alone is considered the rate constants k are smaller than predicted the slope of the plot of k-l against viscosity for a given reaction in a series of solvents is larger by a factor of the order of 10 than the predicted value 1/4RT and the activation enthalpy in a GENERAL DISCUSSION given solvent is often smaller than the value calculated from the viscosity one value being even negative. Crooks and Robinson find however that when they apply more sophisticated diffusion theory (due to solc and Stockmayer) in which only a fraction of the surface of each molecule is available for reaction the slope of the plot of k-' against viscosity can be explained if these fractions have values that appear not unreasonable (see their table 4).They note that the model ignores any effects of solvation and of internal hydrogen bonding. Recent work by Burfoot 'has added two more reactions to the list of those whose rates have been studied in a series of solvents and show anomalous viscosity- dependences (table I). An application of solc and Stockmayer's treatment (for which we are indebted to a personal communication from Dr. solc). in which the molecular surface area available for reaction was taken as the van der Waals area of the N or 0 atom gave predicted values for the slope of the plots of k-' against viscosity that were 2-10 times larger than the experimental values.The values for the reactions of 2,4,-dinitrophenol and picric acid differ by a factor of 5 although if internal hydrogen bonding is important it will affect the geometry of both acids ; and the value for picric aicd is about the same as for trichloroacetic with the same base. The variations of the enthalpy of activation of a given reaction from solvent to solvent and from reaction to reaction is a given solvent do not seem to be explicable by viscosity and geometrical factors alone. It appears that " chemical " factors such as solvation internal hydrogen bonding and complex-formation by dispersion forces must be considered as well. This can be done in terms of the rate constants for the individual steps of a three-step scheme such as given by Crooks and Robinson.More experimental work is needed to make clear the roles of translational and rotational diffusion. Dr. J. E. Crooks (Uttii*ersityof London) and Dr. B. H. Robinson (Utiirersitj of Kent) said We would not claim that the solc and Stockmayer equation gives exact numerical answers because even their model is highly simplified. However their treatment does explain from first principles why the plot of k-' against q has the observed form and why the slope is much greater than that predicted by the Smoluchowski equation. The values of AH for the reactions of Magenta E with a range of amine bases show a poor correlation with AH&c.but the average value of AH is equal to the average value of AH&.within the limits of experimental error. The negative value of AH for 2,4-dinitrophenol as acid may well be associated with the strong intra- molecular hydrogen bond. Prof. L. Melander (Giireborgs Utzirersitei) said Could the writers give some further comments on the tentative transition-state structures V and VI? As they stand it seems somewhat unnatural that the charge should be less delocalised in VI than in V in spite of the fact that the proton has been fully transferred to the nitrogen atom in VI while it is about half-way in V. If VI is the best picture according to the experimental facts is it the a-bond energy in the sultone ring that should be expected to offer resistance to the development of a delocalisation of the same kind as in V? Dr.J. E. Crooks (University of London) and Dr. B. H. Robinson (Utii~~ersiry of Kent) said Transition state V is intended to show a state of the system where the proton-transfer is occurring synchronously with ring opening. VI shows a transition ' K. Solc and W. H. Stockmayer Ifif. J. Cherri. Kirretics 1975 5 733. G. D. Burfoot and E.F. Caldin J.C.S. Furuduy I 1976 72 1O00. GENERAL DISCUSSION 53 state where the proton has transferred completely before ring opening has started i.e. the two processes are decoupled. This is the essential difference we are trying to make between V and VI. If the sultone ring is intact it is expected that most of the charge will be localised on the phenolate oxygen rather than in the benzenoid ring.In fact VI probably represents a highly energetic unstable intermediate and the transition state for the overall reaction will likely correspond to a species where rupture of the sultone ring has just begun. However it will be closest in structure to VI. The delocalization of charge in the particular case of BPB is slow for two reasons (a)As stated by Melander delocalization is slow because a C-0 0bond must be broken in the sultone ring for which some activation energy will be required. (When no 0bond rupture is needed for delocalization as in the case of proton transfer from Magenta E then this process will be very rapid.) (b) In a low-polarity medium the close proximity of the positive charge on the protonated amine will inhibit delocalization of negative charge.We conclude therefore that (i) Delocalization of charge (and by implication sultone-ring opening) is not synchronous with proton-transfer from BPB. (ii) The rate constant for ring-opening is still faster than that for proton transfer to weak bases. A value of lo6 s-' seems reasonable. (iii) Very slow ring-opening/delocalizationcan be observed when a sterically hindered acid is used. (See separate comment by Robinson and Parbhoo). Prof. E. F. Caldin (University of Kent) said The interpretation of a value for the activation volume AV* in a single solvent for a reaction producing an ion-pair is not simple. There are two contributions to AVF one from changes of bonding as the reactant molecules approach each other (AV,') the other from changes in the arrange- ment of solvent molecules (AV:).To an approximation these are additive AV* = AVT+AVZ. (1) If we identify AV; with the volume change AV& due to electrostriction produced by the charge-separation and calculate this for the formation of a point dipole (p*) in a cavity of radius r* in a medium of dielectric constant D,we obtain (putting q = (D-1)/(2D+ I)) This expression is evidently sensitive to the values of p* and r* which are not known. (With p* = 10 D and r * = 5& it gives AV& = -6 cm3 mol-'). In a series of solvents however AV* should be linear with (dq/dp), the intercept at (dqldp) = 0 will be AV and for any particular solvent AVZ can then be found as AV* -LIP':. For the Menschutkin reactions of pyridine with methyl iodide and triethylamine with ethyl iodide in chlorobenzene at 50°C this treatment gives AVZ = -7 cm3 mol-' for each reaction.The charge-development in the transition state of these reactions is probably comparable with that for the reaction of bromophenol blue with amines (the slope of the plot gives p* = 8 D) so we may seek to compare this value of AVF with one H. Hartmann H. D. Brauer H. Kelm and G. Rinck Zeit. phys. Chem. (Frankfurt) 1968 61 53. GENERAL DISCUSSION derived from Crooks and Robinson's value of A V* in chlorobenzene (-16 ~m~mo1-l). In the absence of experimental values in other solvents we estimate AV from a molecular model and use eqn (1). If we suppose that AVlf is the volume difference between OH.. N at the van der Waals distance (0. . N = 3.75 A) and the hydrogen- bonded distance (0.. N = 2.78 A),2 we find AVT = -17 cm3 mol-'. This value is comparable with the experimental value of AV* and would suggest that reorganisa- tion of solvent molecules is not responsible for any large contribution to AV*. The calculated value is however sensitive to the distances assumed. Dr. J. E. Crooks (University of London) and Dr. B. H. Robinson (University of Kent) said As stated by Caldin it is important to be able to dissect AVT into its separate contributions. In the terminology used in our paper the two contributions are A V;, the volume change on forming the intermediate hydrogen-bonded complex (analogous to AV:) and AV,. (AVg) the further volume change on forming the transition state which presumably is dominated by solvation changes in response to charge development (electrostriction).Then (cf. Caldin's eqn (l)) we have AV = AVY,+AV,',. The question is raised as to which of these terms dominates for our system and Caldin presents evidence in support of AV,",. In principle an estimate of AV& can be obtained from (2). However the calculation is sensitive to the size of the cavity. If p* = 3 D and Y* = 1.5 A (corresponding to proton-transfer along a hydrogen-bond) then A V2+-becomes -20 cm3 mol-' which is close to A Vf'. It would also be useful to make the proposed plot when more data become available but there must always be doubt as to whether conclusive evidence can be obtained when the analysis demands that macroscopic parameters are used for interactions in a microscopic environment of the system.(However viscosity works well). The calculation of AV; similarly must ignore any contribution due to changes in orientation and packing imposed by hydrogen-bond formation but these are very difficult to estimate. Direct experimental evidence on the magnitude of AV," is limited but several authors seem to favour a value of -5 cm3 m~l-'.~-' If this is the case then the large value of -AV would be associated mainly with solvent reorgan isat ion. However it is a relatively simple matter to measure AVr2 experimentally for the BPB-pyridine system using the MagE-pyridine system as a model (as used previously for the dissection of AH and AS:).All that is required is a visible spectrophotometer with an optical cell which can be pressurised to a few kbar. Then AYY2 is obtained from the pressure dependence of the equilibrium constant for the model system. It will then be possible to estimate with some confidence the absolute magnitudes of A V,02and A VA. AVf" has recently been measured for BPB and the more sterically hindered base 2t-butylpyridine in chlorobenzene. Although the initial H-bond formation is much weaker (KI2(2tBupy) = 0.66 M-l KI2(py) = 29 M-I) AV is large and negative (-21 cm3 mol-') and close to that observed for the pyridine system. C. D. Hubbard C. J. Wilson and E. F. Caldin J. Arner. Chem. Suc. 1976 98. S. N. Vinogradov and R. H. Linnell Hydrogen Bondin.9 (van Nostrand New York 1971).p. 178. W.J. Le Noble and T. Asano J. Ainer. Chem. Soc. 1975 97 1778. E. Whalley Adv. Phys. Org. Chem. 1964 2 93. E. Fishman and H. G. Drickamer J. Chem. Phys. 1956 24 548. T. Altinata B. H. Robinson and C. J. Wilson unpublished work. GENERAL DISCUSSION Dr. D. M. Parbhoo and Dr. B. H. Robinson (University of Kent) said Some preliminary experiments have been carried out on the system Bromothymol Blue (I) + Pyridine which is related to the Bromophenol Blue (I1 in paper) + Pyridine system. The two indicator acids differ only in the ring substituent groups. Bromothymol Blue (BTB) is more sterically hindered to rotation about the central carbon atom than Bromophenol Blue (BPB) due to the methyl group in the 3 position of the ring.Thermodynamic and kinetic parameters (by the stopped-flow method) have been measured for ion-pair formation from I. From consideration of equilibrium constant values the acids appear to behave similarly. However kinetic measurements indicate two discrete rate processes zf and z (7 > 5zl) for BTB in the stopped-flow time range. In contrast a single relaxation is observed for BPB under the same experimental conditions. For BTB zfis base concentration dependent but z is almost independent of base concentration. (C The results are consistent with the detailed mechanism proposed in the paper for BPB. It would appear that the faster rate process for BTB refers to the formation of an ion-pair species D’ resembling VI in the paper (i.e.proton transfer has been effected but the sultone ring is still intact). The slow process could then be associated with the opening of the sultone ring to form a more stable ion-pair. This involves rupture of a a-bond and delocalisation of charge from the phenolate ion to the sulphonate ion. D’ would have a lower extinction coefficient than D and so such a process can be observed spectrophotometrically. (The migration of the protonated amine cannot be observed in this way). It is possible that the ring-opening reaction is slower with BTB due to the effect of steric hindrance. It is perhaps also significant that BTB fails to form an ion-pair complex (A,, = 560nm)with the much stronger aliphatic amine bases due to the loss of the second proton. The results thus lend support to the idea of proton-transfer (step IV +.VI) decoupled from ring opening (VI -, 111). Prof. M. M. Kreevoy (University of Minnesota) said Many hydrogen bonded complexes in which the basicity of the two basic sites is about equal are mixtures of tautomers as required by the kinetic analysis given by Crooks and Robinson.’ However there are systems in which only a single intermediate structure can be observed. The u.-v. spectra of pyridine-1-oxide and its complexes with a series of acids of varying strength are shown in fig. 1.2 In fig. 2 is shown the spectrum of a solution containing a two-fold molar excess of pyridine- 1-oxide over trifluoro-methanesulphonic acid. The vibrational spectra and freezing point depressions of S. N.Vinogradov and R. H. Linneli Hydrogen Bonding (Van Nostrand Reinhold Co. New York 1971) pp. 163-169. ’K.-C. Chang Ph.D. Thesis (University of Minnesota 1975). GENERAL DISCUSSION such solutions show that simple one-to-one complexes are involved.'* This is confirmed by the X-ray diffraction pattern of a related solid which also shows that the oxygen-oxygen distance is only a little over 2.4A.3 These ultra-violet spectra are inconsistent with the existence of a tautomeric equilibrium in these complexes. The spectra in a series of tautomeric mixtures should show a steady decrease in the A A Inm FIG.1.-The u.-v. spectra of pyridine-1-oxidc and a series of its complexes with carboxylic and sulphonic acids in sulpholane solution. Curve A is for 0.081 M pyridine-l-oxide path length 5.01 pm ; curve B for 0.065 M complex with CH,COOH path length 6.15 pm; curve C for 0.076 M complex with CHCl,COOH path length 6.73 pm ; curve D for 0.101 M complex with CF,COOH path length 14.0 pm ; curve E 0.151 M complex with CH3S020H path length 8.87 pm ; curve F 0.093 M complex with CF3S020H path length 14.3 pm.The progressive blue shift of the charge transfer band is indicated by --; the location of the benzenoid band by---. Because both the concentration and the path length were variable the molar absorptivities A/cl are indicated at one or more points on each spectrum to make a comparison of intensities easier. intensity of the bands due to the free base with a concomitant increase in the bands due to the protonated base without much change in the wavelengths of maxiinurn absorption.If the spectra are superimposed one or more isosbestic points should be observed. Vinogradov and Grunwald and their coworkers have observed such patterns. In contrast each spectrum in the present series contains only one set of bands the location of which shifts continuously with the increasing strength of the acid from that characteristic of the free base to that characteristic of the protonated base. Thus no distinguishable tautomers can exist for more than about s (the K.-C. Chang Ph. D. Thesis (University of Minnesota 1975). J. Husar and M. M. Kreevoy J. Arner. Chenr. Soc. 1972 94 2902. L. GoliE and F. Lazarini Vestnik Sloc. Kern. Dr. 1974 21 17. 'S. N. Vinogradov and R. H.Linnell Hydrogen Bonding (Van Nostrand Reinhold Co. New York 1971) pp. 163-169. D. Eustace and E. Grunwald J. Anier. Chem. SOC.,1974 96 7171. GENERAL DISCUSSION characteristic time of an ultra-violet experiment). The homoconjugate complex of pyridine-1-oxide and its conjugate acid (fig. 2) and also that of 3,5-dinitrophenol and A 200 250 300 h/nm FIG.2.-The u.-v. spectrum generated by a solution of 0.148 M pyridine-1-oxide and 0.074 M CF3S020H,in sulpholane. The path length is 5.91 pm. The molar absorptivity is indicated. its conjugate base,” behave similarly but that of p-nitrophenol is a tautomeric mixture by this criterion.2 In the cases where tautomers do not exist the potential function for hydrogenic motion along the line connecting the two basic sites must have only a single minimum or if there are two the central maximum must not rise significantly above the lowest allowed vibrational It would be of interest to carry out the Crooks and Robinson experiments with substances of this sort since eqn (9) then simplifies to eqn (5).Prof. J. H. Fendler (Texas A & M University) said As part of our interest in the behaviour of surfactants in nonpolar solvents we have investigated the formation of hydrogen bonded ion-pair complexes between indicators and polyoxyethylene(6) nonylphenol in dry benzene as well as apparent dissociation constants in restricted volumes of water solubilized in benzene by surfactants. Bromophenol blue in dry benzene transfers protons to the fraction of polyoxyethylene(6) nonylphenol which is ionized to from two complexes with absorption maxima at 41 1-412 nm (complex I) and 580-590 nm (complex IX) respectively.The equilibrium constant for the for- mation of c~mplex I K, has been calculated at absorbances at 41 1-412 nm from Ao-A 1 log -= log -+log [polyoxyethylene (6) nonylphenol] (1) A-A KI where Ao A and A are absorbances of bromophenol blue in benzene in the absence of polyoxyethylene(6) nonylphenol that in its presence and that when all of bromo-phenol blue is in the form of complex I. The value obtained for K, 66.6 M-I is considerably smaller than those determined for the interaction of bromophenol blue with aliphatic amines in chlorobenzene by Crooks and coworkers. The slope of the line in the plot of the left hand side of eqn (1) against log [polyoxyethylene(6) nonylphenol] at relatively low surfactant concentration is 0.94 implying a 1 1 I.M. Kolthoff M. K. Chantooni Jr. and S. Bhowmik J. Anrer. Chem. SOC.,1966 88 5430. T.-M. Liang unpublished results. J. C. Speakman Structure and Bonding 1972 12 141. GENERAL DISCUSSION stoichiometry. At higher surfactant concentration however there is a significant deviation from this stoichiometry. This deviation is the consequence of the dynamic formation of surfactant aggregates and the interaction of these species with bromo- phenol blue. The spectrophotometric data treated in terms of monomer ~1dimer + trimer . . . +n-mer type association afford the calculation of the average aggregation number of polyoxyethylene(6) nonylphenol in benzene and the relative concentrations of each species.Significantly bromophenol blue assists the formation of higher aggregates. Apparent dissociation constants pKtPP values have been determined for malachite green bromophenol blue thymol blue and methyl orange in water pools solubilized by polyoxyethylene(6) nonylphenol in benzene by titration with HC104. The pKtpp values in this medium differ from the corresponding pK values in bulk water by as much as 7.0 units. For malachite green effects on pKtppincreases with increasing ratios of [polyoxyethylene(6) nonylphenol]/[H,O]. At given HClO concentrations ratios of unprotonated to protonated malachite green increase with increasing concentrations of the surfactant.In no case however is the deprotonation complete. The higher the HC104 concentration the less complete is the deprotonation. Water pools entrapped by polyoxyethylene(6) nonylphenol in benzene provide therefore a relatively basic medium. Dr. R. A. More O’Ferrall (University College Dublin) said In connection with the question of whether a concerted reaction can occur in conditions where it is not ‘‘enforced ” by the instability of an intermediate it has been claimed recently that in olefin forming eliminations of fluorene derivatives i.e. CH3O-/CH30H22 5’ RCHCH2X + RC = CH2+X-where RCH2 is fluorene reaction can occur by a concerted E2 mechanism even though the fluorenyl carbanion is quite stable. The principal evidence for a concerted reaction is the high bromide/chloride elimination rate ratio ; with two methyl substituents a to the leaving groups the ratio is 30.As calculated from the pK,’s of the fluorenes the free energies of the potential carbanion intermediates are some ten kcal lower than those of the E2 transition states. Prof. W. P. Jencks (Brandeis University) said If the leaving group becomes sufficiently good it will leave with no significant energy barrier the “intermediate ” will have too short a lifetime to exist and the reaction must become ‘‘concerted ” rather than step-wise. When it becomes concerted it is likely that there will be significant breaking of the bond to the leaving group in the transition state which in turn facilitates proton removal. Is it known whether there is a significant barrier for expulsion of good leaving groups such as bromide from a carbanion “inter- mediate ” in this reaction? Dr.R A. More O’Ferrall (University College Dublin) said :No that is not known. Prof. W. H. Saunders (University of Rochester) said I think it unlikely that the eliminations from fluorenylmethyl halides follow the E2 path solely because of instability of the carbanion intermediate that would be involved in a stepwise path. Taken to its logical conclusion this hypothesis would put the carbanion along the reaction path at some point after the transition state but make it unobservable as an intermediate because it would decompose to products without an activation R. A. More O’Ferrall and P. J. Warren Chem. Comm. 1975,484.GENERAL DISCUSSION energy. In such a case a sizable bromide/chloride rate ratio would not be observed because the bond to the leaving group would still be intact at the transition state. There may well be different requirements for concertedness for elimination reactions which involve proton transfers from carbon and the type of reactions Jencks deals with which usually involve proton transfers between oxygen and/or nitrogen atoms. In the latter case the proton transfers are diffusion-controlled so long as they are exothermic. Thus no energetic advantage is to be gained from concertedness unless it makes an otherwise endothermic proton transfer less endo- thermic. The proton transfers in elimination reactions are in contrast slow even when exothermic and a concerted process can reduce the activation energy under any circumstances except where a full carbanion on the P-carbon is required before expulsion of the leaving group can begin.This ready availability of a lower-energy path would seem a sufficient condition for concertedness without the necessity for any additional assumption about the instability of the carbanion.

ISSN:0301-5696    

DOI:10.1039/FS9751000050

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Energetic and Dynamical Aspects of Proton Transfer Reactions in Solution BY R. A. MARCUS Department of Chemistry University of Illinois Urbana Illinois 61 801 USA Received 19th May 1975 Several energetic and dynamical aspects of proton transfers are treated. The effect of intrinsic barrier asymmetry on BEBO calculated Bronsted plots is investigated and contributions to work terms are also considered. The dynamics of transfer of a light particle between two heavier ones is discussed for a particular potential energy surface making use of classical trajectories semidassical concepts and a previous quantum study. The question of nonequilibrium polarization of solvent is also considered. 1 INTRODUCTION It is a pleasure to participate in this symposium honouring Professor R.P. Bell whose work has illuminated so many parts of the proton transfer field. In this paper I would like to comment on several aspects of proton transfer both energetic and dynamic (1) effect of '' intrinsic barrier asymmetry " on Bronsted plots (2) dynamics revealed by recent classical and quantum mechanical studies for an H-atom transfer (3) contributions to the "work terms " and (4) the possibility as in electron transfers of nonequilibrium polarization of the solvents. 2 INTRINSIC BARRIER ASYMMETRY AND BRONSTED SLOPES Sometime ago we considered a model of a proton transfer reaction,' AiH+A2 +Ai+HAZ (2.1) (charges are omitted for notational brevity) in which the process occurred in three steps AIH+Az + AIH-* *A2 (2.2) AIK--A2-+ Al.* *HA2 (2.3) Al** *HA2-+ Al +HA2. (2.4) Of these only the middle one depended on the standard free energy of reaction AGO' for (2.1).* Step (2.2) involves a free energy change w' (called a " work term ") for bringing the reactants close together ; wr includes steric (orientation) effects and where necessary,2 partial desolvation. The next step (2.3) is the actual proton transfer and involves intramolecular and solvent reorganization to form the transition state followed by an intramolecular and solvent relaxation. Step (2.4) is a " disorienting " and resolvating one ; it contains a work term -wp wp being the analogue of w' for the reverse reaction. * AGO' is actually the " standard " Gibbs free energy change in the prevailing medium and at the prevailing temperature.60 R. A. MARCUS In an approximation which was quadratic for treating (2.3) and which at the same time neglected “ ),-asymmetry ” (3.f-A2 defined later) the rate constant was given by k = Z exp( -AG*/kT) (2.5) apart from the usual statistical factors.‘k 2 is the collision frequency in solution AG* is AG* = ~‘+(R/4)(1+AGg’/i)~ (IACi’I < i) (2.6) and AG;’ = AG” +W” -w‘. (2.7) i is the “ intrinsic ” free energy barrier,’ i.e. the barrier in (2.3) when AG;’ = 0. AG;’ is seen from eqn (2.7) to be the effective standard free energy of reaction for the proton transfer step itself. Similarly in a bond energy-bond order (BEBO) type of calculation the corres- ponding vaIue of AG* is given by eqn (2.8) when i-asymmetry is neglected and when the E’s in ref.(1) are replaced by free energies AG* = w‘ +044) +(AG;;’/2)+()./4y)ln cosh(2yAG;;’/l.) (2.8) where y is ln2. The difference between eqn (2.6) and (2.8) was typically relatively small.’ Implica-tions of the equations are apparent a small 1implies a large curvature of a Bronsted plot; a small 2 also implies a large limiting rate at large negative AGi’ when w‘ is small but the limiting rate is small when wr is large. A question which arises is the effect of ).-asymmetry. Specifically if a potential energy surface is varied by varying AGO‘ of reaction (2.1) holding constant the intrinsic barriers of the exchange reactions AiH+A + A,+HA, (i = 1 2) (2.9) do the preceding considerations prevail ?3 The intrinsic barrier E.J4 for the reaction in eqn (2.9) may depend on i and the difference in is called here the 2.-asymmetry Differences in 2 and L2 were neglected in deriving eqn (2.6) prompted in part by a finding that such effects were relatively minor in the quadratic case i.e.in eqn (2.6).‘ We consider now their effect on the BEBO derived formula eqn (2.8). The problem is how to calculate the effect of varying AGO‘ holding the intrinsic barriers constant and not assuming I, = Typically a potential energy surface is not automatically characterized in terms of ,il 2 and AG;’. For example in a BEBO model for the reaction in eqn (2.9) the potential energy of formation of an intermediate state can be written as AE = V,-Vln?‘-V2n4* (2.10) where along the reaction path bond order is conserved 12’ +?I = 1 (2.1 I) n and Viare the bond order and bond energy of the A,H’th bond pi is an exponent which reflects a property of that bond.For the exchange reaction in eqn (2.9),ni is 5 in the transition state and so AE for that reaction which we may call AEl or better yet Ai/4 is found from eqn (2.10) to be Vi[I -2(+)p’]. * If srand sp are the statistical factors for forward and reverse steps. it suffices to replace wr and WP by Wr-kTIn sr,wP-W In sP to include their effect.’ Further k is the k in footnote 3 Of ref (1) in the case of diffusion effects. DYNAMICS OF PROTON TRANSFER The potential energy change accompanying the reaction in eqn (2.3) is A V AV = V1-V2. (2.12) Thus the effect of AV on the potential energy barrier to reaction A€ can be investi- gated holding the intrinsic barriers l1and E.constant only by varying p1 and/or p2 simultaneously. (It is not clear that this precaution was followed previ~usly.)~ The value of til in the transition state is obtained by setting dAE/dnl = 0 and intro- ducing the resulting it1 and n2 into eqn (2.10). Investigation of the effect of A V on A€ holding R and Az constant is considerably simplified as in eqn (10) of ref. (l) by noting that p z 1 and expanding nf' in eqn (2.10) in a Taylor series retaining only the first two terms. The barrier A€ is found (eqn (12) of ref. (1)) to be AE = n$AV-(&4y)tii In nf (2.13) i= 1.2 where it! and ni are the solution of dAE/dnf = 0 i.e. of 0 = -AV-(Al/4y)(ln nfi + 1)+(i2/4y)(h nf,+1) (2.14) nt+nt = 1.(Eqn (2.8) can be obtained from eqn (2.13) and (2.14) by setting ;Cl = L2 = I. replacing AV by AG;' and adding to (2.13) the barrier wrof the first step (2.2).) Now at last AE depends only on AG;' and 12. The slope of a (AE AV) plot at a given A1 and 11 is obtained by observing that dAE/dAV is the sum of (dAE/JAV,,?) and of (aAE/an?)dY(anf/aAV).Since (dAE/dn\)Av is zero one finds from eqn (2.13) that dAE/dA V = n!. (2.15) The A€ in eqn (2.13) can be obtained by first introducing values for n? and ni into eqn (2.14) solving the latter for AV and introducing this result into eqn (2.13). In table 1 the results of such a calculation are given choosing a rather large asym- metry A1/4 = 12 and A2/4= 2.TABLE 1.-EFFECTSOF REACTANT ASYMMETRY ON (AE,A V)PLOTS n2 AV AE nz' AV AE 0.1 -19.2 0.4 0.6 0.0 7.2 0.2 -15.2 1.0 0.7 5.4 10.7 0.3 -11.7 1.8 0.8 12.8 16.3 0.4 -8.2 3.1 0.9 25.1 26.9 0.5 -4.4 4.8 From table 1 one sees that the l./4 in eqn (2.6) namely AE at AV = 0 is 7.2. The latter is close to (Al +A,)/2. The Bronsted slope dA€/dAV for the system is seen from eqn (2.15) to be n$. Thus when the true slope is 0.1 0.3 0.6 and 0.8 say one finds from the above A and the corresponding AV's in table 1 that the slope calculated from eqn (2.6) is 0.17 0.30 0.50 and 0.72 respectively values which are fairly close to the true slopes. 3 SOMEDYNAMICAL ASPECTS OF LIGHT PARTICLE TRANSFER Chemical kinetics has received additional insight from recent studies with mole- cular beams lasers and infra-red chemiluminescence.6 On the theoretical side the main method for interpreting these data has involved computer-calculated classical R.A. MARCUS trajectories of the atoms,6b* because of the difficulty of solving the fully three- dimensional reactive collision problem numerically and quantum mechanically. Numerical quantum mechanical studies have been almost entirely confined to collinear collisions.* In the case of proton transfers no trajectory or quantum mechanical numerical studies appear to have been made as yet. Some insight into the dynamics can be obtained by studying instead the transfer of a hydrogen atom between two heavier particles. The only quantum mechanical study which has appeared is that of a collinear collision between HBr and C1.C1+ HBr +-ClH + Br using a London-Eyring-Polanyi-Sat0 potential energy surface. This limitation of collinearity is perhaps not in itself too dismaying; the actual collisions in solution with major steric or solvation features can differ substantially from the usual three- dimensional gas phase collisions. The transmission probability was calculated for the reaction and more specifically for the formation of various vibrational states of the product HC1 of this exothermic proce~s.~To analyze the results of this study and to obtain implications for other light particle transfer Dr. Ellis of this laboratory has undertaken some classical trajectory studies on this and related systems. While the results will be described elsewhere,lO some features are summarized below.3 4 R,a u FIG. 1.-Skewed-axes plot of potential energy contours for reaction (3.1). R1 = RCI-H R2 = RH-B~/C,where C is the usual mass-scaling factor l1 (0.987 here). The dotted line denotes a transition state and a reactive trajectory is also indicated. A diagram of the surface used is given in fig. 1 in the usual skewed-axes form? (As is well known plots in rectangular-axes form while frequently used are misleading for purposes of analyzing the dynamics of individual trajectories.) The radial coordinate is essentially a scaled C1- Br distance while the angular coordinate is the protonic coordinate. In one definition the transition state is the line of steepest ascent from the saddle-point indicated in fig.1 by the dotted line. The latter is seen to be curved in the present highly exothermic instance. A typical trajectory for reactants with an initial zero-point vibrational energy and with a substantial initial translational energy (9 kcal/mol above the barrier height of 1 kcal/mol) is indicated in fig. 1. For most of the trajectories corresponding to these and lower energies the relevant part of the dotted line is effectively perpendicular to the horizontal axis. Thereby the reaction coordinate in this appreciably exother- mic system is essentially the C1 -Br distance. We found that the classical probabilities agreed approximately with the quantum mechanical values for the transitions which were classically allowed i.e.those for which the final vibrational states of products were attainable from the initial ones of reactants via real-valued classical mechanical trajectories. (Classically-forbidden transitions are those which require complex-valued trajectories. 2 A substantial fraction of the trajectories which passed through the transition state region (i.e. across the dotted line) did or did not recross it to reform reactants depending on the initial translational energy. The behaviour in the preliminary DYNAMICS OF PROTON TRANSFER studies appears to suggest that a proper phasing of the H-and C1-Br motions is needed for reaction. The recrossing itself " wastes " phase space. It implies that apart from tunnelling corrections the rate will be typically less than that predicted by transition state the~ry.'~ However even a factor of three as a discrepancy between transition state theory and the actual dynamics is a minor one considering the large variations in rate which can be studied by variation of factors such as AG;'.A second deduction can be made from the classical trajectories using semi- classical l4 arguments Because the zero-point energy of the vibrational motion (more precisely a vibrational " action variable " J)is roughly constant up to the transition state region in the above study the vibrational motion is substantially " adiabatic " l4 in this region of space. The "quantum number " of the vibration N is related to J by the well-known Bohr-Sommerfeld eqn (3.2) for a vibrational coordinate a formula later justified by the WKB solution of the Schrodinger equation.J = (N+t)h. While N can have any real value classically but only integer values quantum mechani- cally the same approximate adiabatic behaviour which led to a tendency to preserve J classically in the present case in the region up to the transition state will lead to a similar tendency to preserve N quantum mechanically in that spatial region. The vibrational energy is for a harmonic oscillator of frequency v equal to Jv both classically and quantum mechanically. Thus apart from minor variations of v in this region the vibrational energy is also roughly constant. Since isotopic effects on the rate constant in the absence of tunnelling are largely attributed to differences in zero-point energies of reactants and the transition state,' there should be essentially no isotopic effect on the rate constant in this appreciably exothermic system when H is substituted for D.Finally a type of Franck-Condon principle also operates in the region where the system moves from one channel to another the momentum of the " slow " coordinate C1-Br being substantially conserved in that region. Here the protonic motion is very nonadiabatic and a significant increase of its vibrational action (and energy) occurs. Thus in the reverse reaction vibrational energy should facilitate the proton transfer an effect which might be observable in a suitably stabilized (e.g. intra- molecularly hydrogen-bonded) system using short laser pulses. In the case of the corresponding thermoneutral system Cl+HC13 ClH+Cl (3.3) the potential energy surface is quite different from that depicted in fig.1. The surface is now symmetrical about the bisector of the acute angle and the dotted line repre- senting the transition state now lies along that bisector. The reaction coordinate is in the vicinity of the transition state perpendicular (as before) to the dotted line and so now is substantially a motion of the proton. The original zero-point energy of the protonic motion has thus been lost or really converted to motion along the reaction coordinate when the system passes across the dotted line region. The full effect of an H and D isotopic difference in zero-point energy is thereby felt yielding a maximum isotope effect (tunnelling corrections aside).These facts are well-known,I5 but it is interesting to see them borne out by the behaviour of the trajectories. The various dynamical results classical and semiclassical thus have implications for approximate dynamical treatments of light particle transfer but we shall omit here further discussion of them. The above remarks apply to potential energy surfaces such as that in fig. 1 and its analogues for less (or more) exothermic reactions. In the case of proton transfers R. A. MARCUS in solution the effective surface is more apt to have potential energy wells in the two channels rather than free escape channels out to infinity wells created by hydrogen bonding or by cage effects. Nevertheless from semiclassical considerations effects similar to those described above are expected to apply in this case also.4 WORK TERMS wr AND wp The work term can be a composite of several terms. In the case of carbon acids or bases which do not participate in hydrogen bonding some desolvation of an attacking nitrogen oxygen base or acid may be needed and not compensated for by a favourable AG;' and so contribute a term wiesto wr. Again in the large molecules which are usually involved and when the reactants are not joined by hydrogen bonding an appreciable steric restriction may occur and contribute a term wit. For example in the gas phase abstraction of a hydrogen atom from an alkane by a methyl group CH3+HR + CH4+R (4.1) a steric factor of the order of can be anticipated,16 and would correspond to a work term wr of about 4 kcal/mol.Such steric factors might be reduced somewhat by favourable AGi' but only a slight effect would be anticipated in the present case. If one assigns to the partial desolvation a contribution of the order of 6 kcal/mol and assumes a steric effect of the above magnitude the net wr for nonhydrogen bonded reactants would be about 10 kcal/mol which is of the same order as that needed to explain the data.2* Another contribution to the work term can also occur when the immediate product of the third step in the reaction eqn (2.4),is not the separated products but rather is a metastable intermediate which later ruptures (cf. eqn (5.1) later). Whenever this last step has an activation barrier Wiec which exceeds the barrier for the intermediate to reform the reactants this wiec should in effect be added to the previously computed free energy barrier.We then have W' = Wies +Wft +wiec. (4.2) Of these w' contributions only the first two contribute to the w' in eqn (2.7). 5 NONEQUILIBRIUM SOLVENT POLARIZATION In electron transfer reactions a charge transfer occurs between two reactants and the " charge centres " are usually some 5 to 10 A apart. In the transition state the electron cannot be in both places at the same time and the solvent orientation- vibrational polarization adopts a value which is some compromise. The solvent electronic polarization on the other hand can largely follow the motion of the electron being transferred. This situation where the nuclear part of the solvent polarization is not that dictated by either charge centre alone and where the electronic part is dictated by the instantaneous position of the transferred electron and by the field due to the nuclear part was termed " nonequilibrium polarization " and treated in some In the case of a simple proton transfer between two adjacent centres as in eqn (2.3) the charge is transferred only over a relatively short distance and an effect such as the above would be expected to be minor.In some cases however the assumed mechanism involves rearrangement of several bonds with a somewhat larger dis- placement of charge in the proton transfer step (5.1). One example might be AH+R1RzC=N+=N-+ A-+RIR2CH-N+ N (5.1) (followed by elimination of N2and by other processes).S 10-3 66 DYNAMICS OF PROTON TRANSFER To obtain the potential energy of the transition state for any given configuration of the nuclei of the reaction complex and of the surrounding solvent the Schrodinger equation is solved for the electronic wave function. When attention is focused on the electrons of the reactants and the electrons of the solvent are treated for reason of simplicity as forming a polarizable dielectric continum one obtains a nonlinear Schrodinger equation. The free energy of formation of a nonequilibrium polarization state with an arbitrary orientation-vibration solvent polarization is given by W,, = -[(I -1/D0,)/8n] J D2dr-J P . D dr+2nc JP2dr (5.2) neglecting dielectric image effects.D(r)is the field directly due to the charges on the reactants l/c is l/Dop-l/Ds r is any point in the solvent P(r) is a function of the arbitrary orientation-vibration polarization and Do and D are the optical and static dielectric constants of the solvent respectively. Ultimately eqn (5.2) can be replaced by a more rigorous statistical mechanical expression but it will suffice for purposes of the present discuqsion. The Schrodinger equation for the wave function $ of the electrons of the reactants for any nuclear configuration Y of the reactants and (positions) of solvent molecules is obtained by minimizing 2o the following functional p($) with respect to ,$ at a given P. (5.3) where r1 denotes the totality of coordinates for the reactants’ electrons and I V$ I2 really denotes a summation over such electrons a b .. . I Va$ l2 + I Vb$ l2+ . . . ; V(r,r,) includes the potential energy arising from interactions within the reactants and with the solvent molecules apart from that included in the relatively long-range polarization term W,,,. UItimately all values of the r are considered and a suitable quantum and statistical mechanical average is made over r,. The D appearing in eqn (5.2) is 1/1 r-r1 I being an abbreviation for a sum over reactants’ electrons 1 /I r-ra I + 1/1 r-ryb I+ . . . When the resulting (nonlinear) Schrodinger equation is solved for $ one obtains a $ which depends on P(r). p($) then becomes a function of P which can then be obtained by then minimizing p with respect to P.In the case of electron transfer reactions it was possible to introduce a simplifying approximation writing $ as a linear contribution of two terms with weak overlap between them one term being the same as for the reactants and the other being the same as for the products and both reactants treated as spherical.18 The results obtained from eqn (5.2)-(5.4) can be shown (Appendix 1) to be equivalent to those obtained l8 earlier by a different and in some respects less general method. To the extent that the electronic wave function for the transition state of the reaction in eqn (5.1) could be similarly approximated for this purpose,21 the previous 4* * results for electron transfers could be adapted to that for proton transfer and added to the contribution to AE in eqn (2.10).When $ cannot be written as a linear combination eqn (5.2)-(5.4) remain applicable but more formidable. Elect-ronic structure calculations for the transition state of reactions such as (5.1) would therefore be helpful. R. A. MARCUS When the electronic energy of the system has been obtained as a function of r and P,the latter remain to be treated statistically as in transition state theory or dynamically. Examples of dynamical treatments for other or related potential energy surfaces are given in ref. (21) and (22). 6 SUMMARY A substantial " reactant asymmetry '' does not have a large effect on the slope of Bronsted plots (Section 2). Possible contributions to the work terms are summarized in Section 4 and the relation of the nonequilibrium polarization study in electron transfers to a possible one in proton transfer is considered in Section 5.On the dynamics side some results and implications of a recent study of dynamics of light- particle transfer are described in Section 3. APPENDIX 1 RELATION OF EQN (5.3) TO THOSE IN REF. (18) If $l denotes the electronic wave function for the pair of reactants as in ref. (18) and t,b2 denotes that for the products a trial $ is This $ is introduced into eqn (5.3) and the variation 69 is calculated at fixed P,and set equal to zero. The 6cl and 6c2 are subject to Cl+C2 = 1. (A2) When the assumption of weak overlap of t+bl and $2 is imposed one can show that one obtains the result that the free energy of reactants with an arbitrary P equals that of the products in this same P environment.This condition is identical with that imposed in ref. (18) to satisfy the Franck-Condon principle for these weak overlap systems. One next finds P by minimizing 9subject to this new constraint obtaining a relation the same as that used in ref. (18). The results in that paper are then obtained when the approximation of spherical reactants is introduced. This work was supported in part by the Office of Naval Research. R. A. Marcus J. Phys. Chem. 1968 72,891. M. M. Kreevoy and D. E. Konasewich Ah. Chem. Phys. 1971 21 243. G. W. Koeppl and A. J. Kresge J.C.S. Chem. Comm. 1973 371. R. A. Marcus J. Chem. Phys. 1965 43 679. H. S. Johnston Adv. Chem. Phys. 1960 3 131. E.g.(a)J. P. Toennies Physical Chemistry an Advanced Treatise ed. H. Eyring D. Henderson and W. Jost (Academic Press New York 1974) vol. 6A chap. 5 ; (b)articles in Faraday Disc Client. SOC. 1973 55 and ref. cited therein ; (c)T. Carrington MTP International Review of Science Physical Chemistry Series One Chemical Kinetics ed. J. C. Polanyi (Butterworths London 1972) vol. 9 p. 135 ; J. L. Kinsey p. 173. 'D. L. Bunker Molecular Beams and Reaction Kinetics ed. Ch. Schlier (Academic Press New York 1970) p. 355 ff; M. Karplus p. 372 ff; J. C. Polanyi and J. L. Schreiber Physical Chemistry an Advanced Treatise ed. H. Eyring D. Henderson and W. Jost (Academic Press New York 1974) vol. 6A chap. 6. E.g. E. M. Mortensen J. Chem. Phys. 1968 48,4029; D. G. Truhlar A. Kuppermann and J.T. Adams J. Chem. Phys. 1973 59 395 and refs. cited therein ; for 3-D calculations see A. Kuppermann and G. C. Schatz J. Chem. Phys. 1975 62 2502 and A. B. Elkowitz and R. E. Wyatt J. Chem. Phys. 62 2504. M. Baer J. Chem. Phys. 1972 62 305. lo R. L. Ellis and R. A. Marcus to be published. S. Glasstone K. J. Laidler and H. Eyring The Theory of Rate Processes (McGraw-Hill New York 1941). DYNAMICS OF PROTON TRANSFER l2 J. Stine and R. A. Marcus Chem. Phys. Letters 1972,15,536 ; T. F. George and W. H. Miller J. Chem. Phys. 1972,56 5668. l3 E. Wigner Trans. Faraday SOC. 1938 34,29. l4 Cf. R. A. Marcus Techniques of Chemistry Investigation of Rates and Mechanisms of Reactions ed. E. S. Lewis (John Wiley and Sons New York 1974) vol.6 pt. 1 chap. 2 ; ref. (50) and (51) cited therein. l5 R. P. Bell The Proton in Chemistry (Cornell University Press Ithaca New York 2nd ed. 1973) ; R. P. Bell Chem. SOC. Rev. 1974 3 51 3. l6 E.g. J. A. Kerr Free Radicals ed. J. K. Kochi (John Wiley and Sons New York 1973) vol. 1 chap. 1 p. 15. A. J. Kresge S. G. Mylonakis Y. Sato and V. P. Vitullo J. Amer. Chem. SOC. 1971 93 6181 ; W. J. Albery A. N. Campbell-Crawford and J. S. Curran J.C.S. Perkin 11 1972 2206; M. M. Kreevoy and S. Oh J. Amer. Chem. SOC. 1973,95,4805 ; A. J. Kresge Chem. SOC. Rev. 1973 2,475. R. A. Marcus J. Chem. Phys. 1956 24,966 979. l9 R. A. Marcus J. Chem. Phys. 1965 43 3477 Appendix I. 2o Cf. S. I. Pekar Untersuchurlgen iiber die Elektronentheorie der Kristalle (Akademie Verlag Berlin 1954).21 For a weak overlap approach see also V. G. Levich R. R. Dogonadze and A. M. Kuznetsov Electrochim. Acta 1968 13 1025 ; Electrokhym 1967 3 739. 22 E.g. K. D. Godzik and A. Blumen Phys. stat. sol. 1974 66B 569 and ref. cited therein ; S. F. Fischer G. L. Hofacker and M. A. Ratner J. Chem. Phys. 1970 52 1934; S. F. Fischer and G. L. Hofacker Internat. Symp. Phys. of Ice Munich 1968 eds. N. Riehl B. Bullemer and H. Engelhardt (Plenum Press New York 1969) p. 369.

ISSN:0301-5696    

DOI:10.1039/FS9751000060

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

The Reaction Complex in Proton Transfer BY Awv I. HASSID,MAURICE AND TAI-MING LIANG M. KREEVOY" Chemical Dynamics Laboratory University of Minnesota Minneapolis Minnesota 55455 U.S.A. Received 22nd April 1975 Catalytic coefficients for hydrolysis of diphenyldiazomethane by a long range of neutral oxygen acids (mostly phenols and carboxylic acids) in 80% DMSO+20% water form a curved Bronsted plot which appears to reach a limiting rate of not more than 10 M-' s-'. This plot can be fitted to the Marcus theory of proton transfer yielding a W' value of about 70 kJ mo1-1 and a A/4 value of about 6 kJ mol-'. For 4 of the neutral acids Hammett p values were obtained by determining HA for three ring- substituted diphenyldiazomethanes in addition to the parent compound.From these the Hammett p associated with W,pw and that for A pn can be determined. The latter is close to zero while pw is -1.5. The primary hydrogen isotope effect for the neutral acids is around 4 and appears to be invariant under change of pKm up to a change of almost 7 units. Phenols other than nitrophenols form a homogeneous series with carboxylic acids as judged by km a PHA and ~HA/~DA values but H+is not a member of the series. These observations can be rationalized by a model in which a reaction complex with a strong hydrogen bond is formed with a very substantial net input of work. The proton transfer itself then takes place with very little additional activation energy beyond the work required to reach the level of the product complex.For a number of years evidence has been accumulating that proton transfer reactions are best described by a five-step mechanism similar to the Winstein mechanism for solvolysis reactions,l as shown in eqn (1)-(5). The first intermediate AH+B +AH 11 B (1) AH 11 B +AH*B (2) AH*B +A-*HB+ (3) A-*HB+ +A-11 HB+ (4) A-11 HB+ +A-+ HB+ (5) AH 11 B is the familiar encounter complex formed by the diffusion of the partners to the shortest range at which they are still non-interacting. It is analogous to the solvent-separated ion pair. The second intermediate AH-B is formed by rotational realignment of the partners interpenetration and reorganization of their solvent shells and in at least some cases substantial restructuring of their primary valence structures.It is analogous to Winstein's intimate ion pair but since the reactants may be of any charge type it is called the reaction complex. The actual proton transfer occurs within the reaction complex (eqn (3)) and the products separate in steps analogous to those leading to the reaction complex (eqn (4) and (5)). The formation of an encounter complex is intuitively required for a bimolecular reaction and there are a number of well-known cases where this is rate-determining.2 There is now a small number of cases where evidence suggests that a reorganization of the encounter complex without proton transfer may be rate-determining 39 thus requiring a second intermediate before the proton transfer stage. (This reaction 69 REACTION COMPLEX IN PROTON TRANSFER complex may actually be a virtual intermediate in at least some cases since there is neither theoretical nor experimental evidence that its reversion to the encounter complex has an activation energy).There is also a considerable number and variety of proton transfer reactions which have been shown to require the input of a very substantial free energy before proton transfer is likely to begix1.6'~ This is deduced from the observation that when extended Bronsted plots are analyzed in terms of the Marcus theory of proton transfer:. lo Wr,which should measure the standard free energy required to produce the configuration in which proton transfer occurs is substantially larger than would be expected for diffusion. Since by definition the formation of an encounter complex requires no substantial input of energy a second intermediate or virtual intermediate is required.In the present paper the reaction of diphenyldiazomethane with a series of oxygen acids is shown to be another such case. The primary isotope effect and the effect of substitution in the aromatic ring have also been studied. These effects suggest that a significant reorganization of the primary valence structure is involved in the process for whch W' is the standard free energy. RESULTS THE BRONSTED PLOT Rate constants kHA for reaction of diphenyldiazomethane (DDM) with a long series of neutral oxygen acids (mostly phenols and carboxylic acids) and with H+ have been measured by spectropliotometric techniques. The reactions involved are shown in eqn (6) and (7).R.D. AH +(C6H5),CN 3 A-=(C6H,),CHNl (ion pair) (6) fast via A-*(C,H,),CHN; --+ N2+ (C6H5),CHOH and/or (C6H&CHA (7) several paths The solvent contained 80% dimethylsulphoxide and 20% water by weight and had a constant ionic strength of 0.19. The ionic strength was maintained with tetraethyl- ammonium perchlorate and in most of the present work this salt constituted at least 90 % of the total electrolyte. The dissociation constants KHA of many ofthese acids at infinite dilution in this solvent were known." For those acids values suitable for the present ionic strength were estimated by subtracting 0.3 from their pKHA.12 Those dissociation constants which were not known were measured in the same medium in which rates were measured by electrochemically determining the pH of solutions of known composition.Rate and equilibrium constants are given in table 1. A Bronsted plot constructed from these data is shown in fig. 1. MARCUS THEORY The Bronsted plot shows the now-familiar curvature. This plot was fitted to the Marcus formulation which is summarized in eqn (8)-( 11):. lo AG* = W'+AG* AG* = (1 +AG"'/A)*3,/4 (1 > AG"'/A > -1) AG* = 0 (-1 > AG0'/2) AGf = AGO' (AGo'/A > 1). A. I. HASSID M. M. KREEVOY AND T,-M.LIANG TABLE 1.-DISSOCIATION CONSTANTS AND RATE CONSTANTS FOR REACTION WITH DDM no. acid K€lAIM a knAIM-ls-l 1 H+ 23.5 7.5k0.2~lo-' lo-' f* g 5_+3x lo-' f 2 CH3S020H 1.4~ 3 C4F7COOH 4.2~ c*f 1.7f 0.3 f 4 CF3COOH 2.7~ '*f 1.4k0.3f 5 2,6-(NO,),C,H3OH 3.5 x 10-4 c 2.7kO.l x lo-' 6 CC12HCOOH 2.7 x 10-4 d 4.920.1 x lo-' 1.4~ 7 ~,~-(NOZ)~C~H~OH 10-4 d 2.0k0.2 x 8 CICH2COOH 4.1 x 5.6k0.3x 9 CsF50CH2COOH 3.5x 5.5k0.2~ 10 4-(N02)CsHjCOOH 2.4~ 4.0f0.2~ 11 2-CI-4-(NO2)C6H3OH 9.1 x 10-7 c 1.2f0.1 x 10-3 12 C6FsOH 3.o~10-7 c 5.8k0.1 x 10-7 c 5.4k0.1 x 13 2,3,5,6-FaCbHOH 1.9~ 10-7 c s.~+o.~x 14 C6HsCOOH 1.2~ 10-3 15 CH3COOH 2.0x i.9ko.i x 10-3 10-9 c 4.150.2 x 16 4-(NO2)C6 H,O H 6.2~ 17 3,4,5-CI 3C 6 H 2OH 9.4x 10-'O 2.1f0.3~ 18 4-(CN)CsH2OH 3.0~ 9+ 1x 19 (CF3)2C(OH)2 1.6~ 5.2k0.1x a In a medium containing 80% DMSO+20% water by weight and an ionic strength of 0.19 made up for the most part of the tetraethylammonium perchlorate.b This value is dimensionally C homogeneous with the others. Its numerical value is of dubious significance. Measured in the present work. d Adapted from ref. (11). e Adapted from J.-P. Morel Buff. Soc. Chim. 1967 1405. f Because of the strength of the acid up to half of the inert anion was replaced with the active anion in measuring both the rate and equilibrium constants. g This equilibrium constant may be in error by as much as SO% but its general order of magnitude was verified by Ranian spectroscopy in more concentrated solutions. The notation is the same as that used previously except that a subscript R has been dropped and Fhas been replaced by G throughout in order to conform with Faraday Division practice. When the usual assumptions are made that W' A and Wp,the counterpart to W' for the reverse reaction are independent of the acid strength eqn (12) is obtained from eqn (8) and (9)6- lo AG" = W' +{ 1+(AGOHA +C)/A.}21./4.(1 2) Eqn (13) with the coefficients shown is exactly equivalent to eqn (12) log kHA = a +b log &A +C(l0g (13) A = -2.3 RT/4c Wr= 2.3 RT {log (kT/h)-a +b2/4c) C = (2b-1) (-2.3 RT/4c). Eqn (13) illustrates most clearly that if the Bronsted equation is regarded as a truncated power series the Marcus Formulation can be considered simply to add one more term. Eqn (13) was used to fit the data in table 1 by the method of least squares. Several sets of claculations were made. If all the data except the point for H+ were fitted W' was 73f2 3 was 32+9 C was -27+7 all in kJ mol-l.If the correlation was limited to the 12 carboxylic acids and phenols other than nitro phenols W' was 72.3 k0.1 1was 22 &2 C was -25 f2. (In this case it was necessary to switch from eqn (9) to eqn (11) for points numbered 17 and 18). The correlation coefficient was 0.962 in the first case and 0.998 in the second. The latter correlation is shown as the REACTION COMPLEX IN PROTON TRANSFER solid line in fig. 1 and the former as the dashed line. If additional terms were arbitrarily added to the power series (eqn (13)) the correlation coefficient did not improve. t 1 I I I I I I I I I I 1-0-1w: -1 -a -I I -I I I I I I I I I I -10 -9 -8 -7 -6 -5 -4 -3 -2 -1 0 1 log KHAqh FIG. 1.-The Bronsted plot.Unadorned open circles represent points generated by phenols ; barred open circles points generated by nitrophenols ; closed circles points generated by carboxylic acids; crossed open circles points generated by other acids. The numbers correspond to the numbers in table 1. The dashed line is that generated by eqn (12) or (13) using the data points for neutral acids. The solid line is generated using only the points representing carboxylic acids and phenols other than nitro phenols. HAMMETT CORRELATIONS For H+ and four of the neutral acids rates of reaction with three ring-substituted diphenyldiazomethanes have also been measured. These results are shown in table 2. The values of log kHA are linear functions of the Hammet 0’s l3 for the ring substituents as shown in fig.2. Eqn (14) an adaptation of the Hammett equation was fitted to these data by log kEA = pHAa+log kEA (14) inspection. (The superscript H here indicates a rate constant for reaction with DDM itself; the superscript R that for reaction with a ring-substituted DDM). The five values of PHA are shown in table 3. 2.-vALUES OF k~/M-l TABLE S-l AS A FUNCTION OF RING SUBSTITUTION acid ring substituents H+ HCClzCOOH CH3COOH c6F50H 3,4,4-c13C sH2OH p-C1 p-C1 0.23h0.04 0.112f0.003 4.3f0.2~10-4 1.49f0.13~ 10-3 5.7f0.1 x 10-6 p-c1 0.41 ~t0.02 0.21 fO.01 9.2f0.1 x 10-4 2.9f0.2~10-3 1.13&0.03x 10-5 none 0.75 *0.02 0.49 hO.01 1.88*0.06x 10-3 5.8*0.1 x 10-3 2.1 ho.3 x 10-5 P-CH~ 1.05 h0.03 0.86Zk0.06 3.2Zk0.1 x 10-3 1.01Zk0.04~ 10-2 3.8h0.1 x 10-5 The PHA values can be related to the Marcus parameters as follows.The quantity -2.3 RT log kRHA,which is AG is given by eqn (12). The subscript R is not now the Marcus subscript R but denotes quantities associated with ring-substituted A. I. HASSID M. M. KREEVOY AND T.-M. LIANG *R FIG.2.-Hammett plots for four neutral acids and H+. Slopes PHA are shown in table 3. TABLE 3.-HAMMETT PARAMETERS FOR THE VARIOUS ACIDS acid HA PITA AGA,/kJ mol-1 H+ -1.10 -8 HCClzCOOH -1.43 20 CH3COOH -1.34 44 CsF50H -1.36 37 3,4,5-C13CGHzOH -1.30 51 DDM's. All the other parameters which refer to the reactions of the substituted diphenyldiazomethane are similarly subscripted. A similar equation gives -2.3 RT log k:A in terms of the parameters for the unsubstituted diphenyldiazomethane which are subscripted with H.When these expressions are substituted in eqn (14) the multiplications carried out and terms collected eqn (1 5) results This equation contains too many hard-to-evaluate terms to be useful in itself but a useful simplification results if the assumption is made that all the II values are the same. This is a reasonable assumption since by definition A is the activation energy for a proton transfer in which AGO' is zero. Thus the base strengthening or weakening effect of the substituent is compensated for by pairing the substituted diphenyldiazo- methane with an arbitrarily stronger or weaker acid. There is no apparent reason why a p-substituent should have any effect on such a quantity.The simplified result is eqn (16) REACTION COMPLEX IN PROTON TRANSFER Both W' and C are standard free energies and can be fitted to the Hammett equation as shown in eqn (17) and (18). pw is formally the rho-value for the Wi-WA = -2.3 RTpWa (17) CR-C,,= -2.3 RTpCo (18) product K1K2 where these are the equilibrium constants for the reactions shown in eqn (1) and (2). Since K1 should not depend on the electronic properties of the p-substituents however in practice it should approximate the p for eqn (2). C is (AGo'-AG;IA) but since AGGA is independent of the p-substituents pc is the p for eqn (3). Use of eqn (17) and (18) simplifies eqn (16) to eqn (19) Eqn (19) formally predicts that PHA should not be a constant; that is the plots shown in fig.2 should be nonlinear because (CR+C,)/4% varies systematically with the electronic character of the p-substituents. However it will subsequently be shown that this dependence is very shallow so that the linear plots shown in fig. 3 are not actually inconsistent with the equation. Eqn 19 indicates that pHAshould be a linear function of AG~A with slope pc/23.. The plot is shown in fig. 3. It is linear for the four molecular acids insofar as can be determined from just four points and gives a value of 0.2 for pc when 22 kJ mol-' is used for A. From this and the intercept a value of -1.5 is obtained for pw. The point for H+ as HA falls quite far off the line which is not surprising since this point also falls substantially off the Bronsted plot.0 10 20 30 40 50 AGHA FIG. 3.-sHA as a function of AGGA. The numbers identifying the acids correspond to those in table I. Returning to the variation in pHAwith R predicted by eqn (19) because of the variation of (CR+cH)/41 with R; insertion of 0.2for pc in eqn (18) reveals a variation in cRthat leads to a trivial variation in (cR+cH)/4/! over the range of a values encompassed by the present study. ISOTOPE EFFECTS For the same four molecular acids and for H+ rates were also measured in a D,O+DMSO mixture with the composition chosen so as to have the same mole fraction DMSO as the H,O+DMSO mixture 0.479. Isotope effects kHA/kDA were A. I. HASSID M. M. KREEVOY AND T.-M. LIANG obtained from rates measured at about the same time in the same thermostat.(Although the k, values corresponded very well with those measured at other times). Before and after the kinetic experiments the hydrogen content of the D20component of the solvent was determined by comparing the intensity of its signal in the n.m.r. spectrum of the sample with that of the 13C satellite of the DMSO peak. A short extrapolation to the results for completely deuterated solution was made using eqn (20) for the molecular acids and eqn (21) for H+.14 The hydroxylic hydrogen content of these solutions was never more than a few percent and the difference between k,/k and kD/kHwas never as large as 20%. The resulting isotope effects are shown in table 4. TABLE 4.-kOTOPE EFFECTS AS A FUNCTION OF ACID STRUCTURE acid kHlkD a AGo’/kJ mol-1 H+ 2.05& 0.04 -33 3.52 HCCIZCOOH 4.0k0.I -5 C,F,OH 3.9+ 0.2 12 CH3COOH 3.8k0.3 19 (3.8) 3,4,5-C13C 6HzOH 4.8k0.7 26 UEach value is the average of 4-10 individual determinations.The cited uncertainties are probable errors. b Primary hydrogen isotope effects determined by competition assuming that the proton is transferred directly from the H30+unit to the substrate (J. M. Williams and M. M. Kreevoy Adv. Phys. Org. Chem. 1968 6 63). For acetic acid this should be equal to km/kDA determined directly. For H+ the kinetically determined value contains a secondary solvent isotope effect so the value determined competitively is more comparable with ~HA/~DA for a molecular acid. C This value is more uncertain than the others because of phenoxide-catalyzed exchange between the hydroxylic hydrogens and the methyl hydrogens of the DMSO and because of other experimental difficulties associated with the slowness of the reactions.DISCUSSION The results shown in fig. 1 and tables 1 3 and 4 clearly show that carboxylic acids and at least those phenols which are not nitro phenols form a single series of Bronsted acids. There is no discontinuity of rate slope of the Bronsted plot Hammett p or isotope effect on going from carboxylic acids to phenols of similar acid strength. The Bronsted plot is clearly nonlinear. As shown in fig. 1 it goes from a slope of about 1 .O to a slope only a little above zero. The systematic curvature is adequately represented by a quadratic equation.The coefficients of the quadratic are determined with reasonable reliability. They correlate the results with a high correlation coefficient. The inclusion of a cubic term in the power series does not improve the correlation and the uncertainly in the coefficient of the cubic term is as large as the coefficient itself. When the coefficients are interpreted in terms of Marcus theory W‘ is much larger than A/4. That is when AGO’ is zero or negative most of the free energy of activation is expended in forming the reaction complex from the separated starting materials. REACTION COMPLEX IN PROTON TRANSFER Qualitatively this conclusion emerges from the fact that the rate becomes almost independent of the acid strength at a level below the diffusion limit by a factor of nearly 1O'O.It does not depend on the specifics of the Marcus formulation and we cannot readily imagine a theory that would lead us to a different conclusion from this observation. The variation in rate with substitution in the aromatic rings also appears mostly as a variation in W'. The small value of pc is easily understandable. The trans- formation to which it pertains shown in fig. 4 in first approximation does not involve a net shift of electrons either toward or away from the carbon to which the aromatic rings are attached. Whatever small withdrawal of electrons occurs is apparently outweighed by the change in hybridization from sp2 to sp3. The latter orbitals are less electron-attracting than the former. Ar Ar 1 + -AGO' -I A-H * * -* C=N=N + A -* -H-C-N=N+ -I I Ar Ar FIG.4.-The reorganization of bonds and charges in the proton transfer step.By comparison with the structure sensitivity of hydrogen bond formation pw is not unreasonable if the reaction complex is regarded as having a fairly strong hydrogen bond. For example Clotman and Zeegers-Huyskens l6 have found Taft p values as high as 1.07 for the formation of hydrogen bonds between amines and phenol. Such a hydrogen bond would not be formed spontaneously to the central carbon of diphenyldiazomethane and the value of W' must reflect the energetic cost of a relatively small acid-base distance (with a separation of perhaps 2.5 or 2.6 nm between oxygen and carbon). Such a separation would permit the transfer of the proton with activation energy of the order of 3 when AGO' is zero.17 The only section of the present results not entirely well accommodated by this model are the isotope effects given in table 4.The effects for molecular acids would have been expected to go through a maximum between dichloroacetic acid and pentafluorophenol where AGO' is zero. The near-constancy of the isotope effect suggests that it all lies in W' and that 3 is relatively insensitive to isotope effects. This is contrary to previous ideas and experience.s* lo The equilibrium constant for the formation of the reaction complex might well have an isotope effect as large as a factor of about two. Such an isotope effect has recently been observed for the formation of the strongly hydrogen-bonded complex between trifluoroacetic acid and DMS0.12 Although ad hoc explanations can be constructed it is not clear why the balance of the isotope effect does not show the expected structure-sensitivity.It appears to us that the foregoing model is capable with appropriate adjustment of the various parameters of spanning the whole gamut of data for proton transfer to and from carbon. Two puzzling observations in particular can be rationalized. In terms of eqn (19) the positive p values for rates of reprotonation of nitronate anions l9 is probably due to a positive pw rather than to an anomaly in the proton transfer itself. Such a pw value could be understood as the result of a need to convert the nitronate ion with its negative charge concentrated on its oxygen atoms to a nitrocarbanion with its charge localized on carbon in the process of forming the reaction complex.The large scatter in many plots of isotope effect as a function of ApK can be attributed to the use of a variety of substrates. In terms of the Marcus theory in order for such plots to be free of scatter only C can be structure- sensitive; W' Wp,and 3 must be constants. It is now clear that the work terms ate A. I. HASSID M. M. KREEVOY AND T.-M. LIANG in fact quite sensitive to the structure of the carbon acid or carbon base. For various diazoalkanes reacting with neutral oxygen acids W'has now been shown to vary from 32 to 72 kJ mol-'.6-* This implies that reactions with the same ApK may have quite different Bronsted parameters and correspondingly different isotope effects.Such a conclusion is reinforced if W' is isotopically sensitive to varying extents. When a variation in isotope effect was produced by varying ApK without varying the substrate structure less scatter was observed.8 We thank the U.S. National Science Foundation for financial support of this work through grant GP-31360X to the University of Minnesota. D. J. Raber J. M. Harris and P. v. R. Schleyer Ions and Ion Pairs in Organic Reactions ed. M. Szwarc (John Wiley & Sons New York 1974) vol. 2 pp 253-256. M. Eigen and L. de Maeyer Technique of Organic Chemistry. Investigation of Rules and Mechanisms ed. E. S. Lewis and A. Weissberger (Interscience Publishers Division John Wiley New York 1961) vol.8 part 2 pp 1034 and 1035. M. M. Kreevoy J. Dolmar and J. T. Langland Abstracts of Papers 166th Natl. Meeting Amet-. Chem. SOC. Aug. 26-31 1973 PHYS 148. G. D. Burfoot E. F. Caldin and H. Goodman J. C. S. Faraday I 1974 70 105. L. Melander Arkiu Kemi 1961 17 291. M. M. Kreevoy and D. E. Konasewich Ado. Chem. Phys. 1971 21 243. 'W. J. Albery A. N. Campbell-Crawford and J. S. Curran J.C.S. Perkin ZI 1972 2206. M. M. Kreevoy and S. Oh J. Amer. Chem. SOC. 1973,95,4805. A. J. Kresge Chem. SOC. Rer;. 1973 2,475. lo R. A. Marcus J. Phys. Chem. 1968 72 891. E. H. Baughman and M. M. Kreevoy J. Phys. Chem. 1974 78,421. This change with ionic strength was derived from Debye-Huckel theory ; A. I. Hassid Ph.D. Thesis (University of Minnesota 1974) p. 93.However the theory can not be expected to apply to solutions as concentrated as our present ones. To check this estimate several PKHA values were remeasured in the same medium in which rates were measured. The agreement was satisfactory in all cases. I3 L. P. Hammet Physical Organic Chemistry (McGraw-Hill New York N.Y. 2nd ed. 1970) p. 356 ; The CJ relevant to the di-p-chloro compound was assumed to be twice that for a single p-chloro subs tit uent . l4 V. Gold Ado. Phys. Org. Cliem. 1969 7 259. l5 R. W. Taft Jr. and M. M. Kreevoy J. Amer. Chem. SOC. 1957 79,4011. I6 D. Clotman and Th. Zeegers-Huyskens Spectrochim. Acta 1967 23A 1627. I7 In symmetrical bicarboxylate ions such as bis-trifluoroacetate the barrier seems to disappear entirely at oxygen-oxygen distances of about 2.4 nm ; A. L. MacDonald J. C. Speakman and D. Hadii J.C.S. Perkin II 1972 1151. Is K.-C. Chang Ph.D. Thesis (University of Minnesota 1975) p. 61. F. G. Bordwell and W. J. Boyle Jr. J. Amer. Chem. SOC.,1972 94 3907.

ISSN:0301-5696    

DOI:10.1039/FS9751000069

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

Proton Transfer Reactions in Macrocyclic Complexes and in Metal-peptide Complexes BY CHARLESE. BANNISTER,DALEW. MARGERUM* AND JOHN M. T. RAYCHEBA LOUISF. WONG Department of Chemistry Purdue University West Lafayette Indiana 47907 Received 12th May 1975 The rates of proton transfer reactions at the ycarbon atom in macrocyclic tetraazadiene complexes and at the nitrogen atom in metal peptide complexes are several orders of magnitude slower than the reaction rates of typical oxygen and nitrogen acids with similar pKa values. The Bronsted plots for the macrocyclic complexes deviate little from a slope of 0.5 over a wide range of APKa values while the metal peptides undergo a relatively fast transition from 01 values of unity to zero without reaching the diffusion limiting rates.In order to fit the data to the Marcus theory it is necessary to include in addition to the reorganizational barrier (A/4) a term Wi which represents the solvent reorgan- ization necessary to initiate hydrogen bonding after the encounter acid-base species are formed. The Wi term is in addition to the work necessary to form the encounter species but it is independent of ApK,. The metal-macrocyclic complexes have large h/4 values and small Wi values while the metal-peptide complexes have small h/4 values and large Wi values. The proton transfer kinetics of two types of coordinated ligands are examined in this work. The 14-membered macrocyclic-tetraazadiene complexes have proton addition or loss at a carbon atom (eqn (I)) while the metal-peptide complexes have H H HB MLf MLH2+ proton addition or loss at the peptide nitrogen accompanied eventually by changes in metal-peptide bonding (eqn (2)).The two types of complexes share some common M' M features. In both cases their bases lack readily available electron pairs for hydrogen bonding. The copper(I1) and nickel(I1) complexes of both types have pKa values in the range of 6-9 but their proton transfer rates (table 1 and table 2) are several orders of magnitude slower than is the case for "normal " acids and bases. Normal 78 C. BANNISTER D. MARGERUM J. RAYCHEBA AND L. WONG acids and bases according to Eigen's cla~sification,~ would have H,O+ and OH-rate constants in the vicinity of 1O'O M-I s-l. On the other hand the behaviour of the macrocyclic complexes and the peptide complexes with general acids and bases are not at all similar.Fig. 1 presents the data for the macrocyclic diene carbon acids of Nirr and Cul*,where the Bronsted a and p values are 0.50 over almost the entire ApK range of 20 units. By contrast the (log k log KH) plots for the triglycine complexes Cu(H_,glyglygly)- and Ni(H-,glyglygly)-have an S-shape as seen in fig. 2 with a values which appear to change rapidly from 0 to 1 to 0. APKa FIG.1.-Bronsted plots for the reactions of metal macrocyclic complexes (eqn (1)) with OH- H20 nitrogen bases and their conjugate acids. 0 NiLH2+ (pK 6.21) and NiL+ A CuLHZ+ (pKa 9.35) and CuL+ ApK = PKa(MLH)-PKa(HB). The bases from left to right are OH- OH- glycinamide Tris Et3N Hen+ Tris H2dienz+ Hen+ 2,6-lutidine H20 H20.r -4 16 14 12 10 8 6 4 2 0 -2 PKHB FIG.2.-Kate constants for the protonation reaction of the triglycine complexes (298 K) A Cu-(H-,glyglygly)-with the acids HzO H3B03 HzEDTA2- HOAc Hoxalate- and H30+; 0 Ni(H-2glyglygly)- with the acids HzO H3B03 H2EDTA2- Hmaleate- Hsuccinate- HOAc Hfumarate- Hoxalate- H2gly+,H30+. The detailed nature of proton transfer reactions which occur in the vicinity of metal ions has had relatively little study. With the peptide complexes the weakening or cleavage of the metal-N(peptide) bond during the protonation and the possibility of initial protonation of the peptide oxygen must be considered. These complicating factors do not occur with the macrocyclic carbon acid complexes where the metal- nitrogen bonding remains intact and there is only one atom which can lose or gain PROTON TRANSFER REACTIONS IN METAL COMPLEXES the proton.Since the acid-base rate constants of both types of complexes are large it is useful to consider a general mechanism which includes the diffusion limiting situations. GENERAL MECHANISM A mechanism for proton transfer reactions is given in eqn (3) where k and k-3 are defined as diflusion-controlled rate constants for the k2 k2 k3 A-H+B + (A-H)(B) + (A)(H-B) + A+HB (3) k-I k-2 k-3 eiicounter of the acids and bases and k- and k3 are for the diffusion-controlled separation of the encounter pairs. The k and k- rate constants in this mechanism include all the reorganizational energies needed (1) to establish hydrogen bonding after the (AH)(B) encounter complex is formed (2) to transfer the proton from A to B and (3) to break the hydrogen bonding in the encounter complex (A)(HB).If the mechanism is divided into these steps with encounter complexes as intermediates then a steady-state approximation can be used to give the forward rate constant k in eqn (4). This equation can be rearranged k-kIk2k3 -k2k3+ k- ik3 +k- 1k-2 (4) using the definition of the overall equilibrium constant K = (klk2k3)/(k-lk-2k-3) to give eqn (5) kl k-l+(k- 1/k2)+(k1/k-3)(1/K) (5) * When there is very little reorganizational energy needed for proton transfer the rate constant k2is very large and the middle term in the denominator of eqn (5)drops out giving eqn (6).This is typical of the k-kl (6) -1+(W-3)(1W) behaviour of " normal " acids and bases where kf equals k-,K = (klk2k3/k-lk-2) for small values of K and k equals k for large values of K.3 When the reorganizational energy (due to solvation changes hydrogen bonding and internal reorganization) becomes appreciable then the middle term in eqn (5) must be considered. The value of k2equals kiK; where k,"is the rate constant for proton transfer in the encounter complex when AG," is zero and x depends in part on K2. It is convenient to define k,"as the observed rate constant when K = 1. Then the observed rate constant for any value of K is given by eqn (7) provided that the k2 step is helping to limit the rate The value for x is derived from the activation energy for kf[k,= (kT/h)exp( -AG*/ RT)]according to the Marcus theory 4* as modified by Kreevoy,6 AG* = WR+(;1/4)(l+AGi/;1)' I AG; I < A (8) where AG; is the standard free energy of reaction within the reaction complex and W' is the work required to form the reaction complex.In our case the work required to convert the encounter complex to the reaction complex is Wd = (WR-AGY) and C. BANNISTER D. MARGERUM J. RAYCHEBA AND L. WONG similarly for the products Wi = (W,+AG;). Then it can be shown that x is given by eqn (9) and x, the value of x when K = 1 is given by eqn (10) 1 W'-W' RT x = -+ -[In K-111 KlK3] 2 21 411 x-xo = -RT In K. 41 When k," is small as with carbon acids eqn (11) holds and if A is large x is approximately equal to 1/2 while (x-x,) approaches zero k = /c;K~/(K -xo 1~3)~ (11) so that k = k;KO.' the frequently observed Bronsted relati~nship.~ The proton transfer reactions of the macrocyclic complexes where k," is 102-7M-' s-' for amine bases appears to fit this relation over a wide range of K values (log K = ApK in fig.1). Several interesting possibilities arise using eqn (7) when k,"has intermediate values (i.e. lo4to lo8 M-I s-' ). If a sufficiently large ApK range were possible the values of dlnk,/dlnK (c&. or P0bs.j would change gradually from 1 at very low K values to approximately 0.5 at K = I and gradually approach zero at very high K values. However the equations also predict that it would be possible to have c&.values which vary from 1 to 0 over a relatively narrow ApK range without k approaching the diffusion-limiting value of k,. Thus when 3 is small the value of x can change rapidly from 0 to 1 as K changes and yet the middle term in the denominator of eqn (7) can still predominate if WF; is large. The predicted &bs. value in eqn (12) is obtained by taking dlnk,/dlnK from eqn (1 1). It can be seen that a small value of A will cause a,bs. to change rapidly with InK. Naturally when 1-is small the Wd values must be relatively large to keep the k values below the diffusion limit. Kreevoy and Oh have reported this type of rapid change in c(obs. for the reactions of diazoacetate ion with R3NH+ acids. We appear to see the same effect for ctobs.in the reaction of acids with the metal peptide complexes. PROTON TRANSFER REACTIONS OF THE MACROCYCLIC COMPLEXES The NiLH2+(pK 6.21) and CuLH2-+(pK 9.35) complexes undergo large changes in their molar absorptivities (at 330-360 nm) when they react with base to form NIL+ and CuL+ (eqn (1)). The reactions can be observed by stopped-flow methods without the need to resort to coupled indicators and as a result relatively accurate rate constants are obtained. As seen in table 1 the H30+ rate constants and the OH- rate constants are lo3 to lo5 times smaller than is the case for " normal " proton transfer rate constants from oxygen or nitrogen acids and bases. In this respect the reactions are not unlike those involving the keto and enolate anion of acetylacetone (eqn (13)).H I HH PROTON TRANSFER REACTIONS IN METAL COMPLEXES However the macrocyclic diene ligand complexes offer some advantages for the study of carbon acids. (1) They do not have the second proton-transfer cycle involving the enolate acid (proton transfer to the oxygens) which complicates the acetylacetone TABLE 1 .-SECOND-ORDER RATE CONSTANTS FOR THE REACTIONS METAL MACROCYCLIC ACIDS AND BASES (EQN (I)) 298 K 0.1 M NaC104 complex a I.eactant b k/M-1 s-1 NiLH2+ OH-4.0~ lo6 105 CuLH2-OH-4.9~ NiLH2+ H20 2.6~ CuLH2f H2O 4.5 x 10-4 NiL+ H30f 2.3~ lo6 CuL+ H30+ 5.5~ 107 NiL+ H20 2.0~ 10-3 CuL+ H2O 3.3x lo-' * The acid dissociation constants for NiLH2+ and CuLHZ+ are 6.21 and 9.35 respectively in 0.1 M NaCIO at 298 K.b The acid dissociation constants for H30+ and H20are -1.74 and 15.52 respectively in 0.10 M NaC104 at 298 K. reactions.8 (2) The pK value of the acids can be changed by varying the metal ion without affecting the nature or geometry of the groups adjacent to the reaction site. (3) The large absorbance changes permit the reactions to be monitored directly. Table 3 gives the resolved rate constants for the rezctions of a number of acids and bases TABLE2.-RATE CONSTANTS 298 K a FOR REACTIONS OF M(H-2glyglygly)-WITH ACIDS WHERE M = Ni2+OR Cu2+ AND GLYGLYGLY = TRIGLYCINE complex b HB pKa(KB) kiiB/M-lS-Ni(H-2L)-H3B03 9.00 1 .ox lo-' 10' H2EDTA2-6.00 1.8~ Hmaleate-5.70 1.ox lo2 Hsuccinate-5.28 3.1 x lo2 HOAc 4.64 6.7~ lo2 Hfumarate-4.39 4.7x lo2 Hoxalate-3.51 2.5~ 104 H2!3lY+ 2.36 5.5~ 103 H30+ -1.74 7.3 x 104 Cu(H-2L)-H3B03 9.00 2.2 103 H2EDTA2-6.00 2.6~ HOAc 4.64 3.4~ 104 105 Hoxalate-3.51 3.9~ H30+ -1.74 1.3 x 107 For Ni(H-,L)- in 0.30 M NaC104 and for CU(H-~L)- in 0.10 M NaC104 solutions.b The acid dissociation constant for Ni(H-'L)- and Cu(HdIL)-are 7.7 and 6.7 respectively. with the nickel and copper macrocyclic complexes. All the points in the Bronsted plots in fig. 1 are for amine bases H,O and OH-reacting with MLH2+. The fact that the data for the NiI1and Cull complexes can be superimposed suggests that the reorganizational energies necessary for the proton transfer are the same for both complexes and that the pK differences adequately reflect the effect which the change of metal ions has on the rate constants.The k; value for the acetylacetone (acac) reaction in eqn (13) is lo3*'M-' s-' at C. BANNISTER D. MARGERUM J. RAYCHEBA AND L. WONG TABLE 3.-sECOND-ORDER RATE CONSTANTS FOR THE REACTION OF ACIDS AND BASES WITH THE METAL MACROCYCLIC COMPLEXES 298 K 0.10 M NaC104 acid base PaKWB) kBIM-1 s-~ kan/M-l S1 CuLH2+ 2,g-lutidine Hen+ 6.84 7.09 2.5~10' 8.1 x 10' 8.o~103 1.5 x 104 Tris 8.OO 1.6~lo2 3.6~103 EtSN 10.77 1.6~103 6.0~lo1 OAc- 4.64 2.3 x 10' 1.2x lo6 NiLH2+ gly-H2dien2+ 9.62 4.22 3.4x 10' 4.2~103 2.3 x 103 3.3 x103 Hen+ 7.09 2.3 x103 3.1 x lo2 Tris 8.OO 3.4~103 5.6~lo1 gl ycinamide gb-HPOZ- 8.04 9.62 6.70 8.4~103 1.1x105 6.4~104 1.3 x lo2 4.3x lo1 1.4~104 malonate2- 5.27 4.1 x 103 3.5~104 citrate3- 5.65 1.3 x104 4.6~104 aThe acid dissociation constants for NiLH and CuLH are 6.21 and 9.35 respectively.TABLE 4.-DEVIATIONS OF THE SECOND-ORDER RATE CONSTANTS FOR THE REACTION OF CHARGED BASES REACTING WITH MLH2+ acid base z Alog ka CuLH2+ Hen+ +1 +0.3 OAc-1 +1.0 glycinate -1 +0.8 NiLH2+ H2dien2+ +2 -0.2 Hen+ +1 +0.2 glycinate -1 +0.6 malona te -2 +1.4 HPOZ-2 +1.7 citrate -3 +1.7 aAlog k~ is the difference in log kB between the charged base and the Bronsted line given in fig. 1. The linear least square line using OH- H20and nzutral bases gives log k~ = 2.7-0.5 ApK where APKa = pKa(MLH)-pKa(HB). ApKa FIG.3.-Bronsted plot for the proton transfer rate constants of the metal macrocyclic complexes with OH- carboxylate bases and H20.0 NiLH2+,A CuLH2+ with the bases (left to right) OH- OH- gly- gly- OAC- HzO H20.PROTON TRANSFER REACTIONS IN METAL COMPLEXES 300.5 K which is very similar to the k,"value of lo2.' M-' s-' at 298 K; for the metal macrocyclic complexes with amine bases (eqn (I)). However the Bronsted plots for acac have significantly more curvature. In fact there is so little curvature in fig. 1 that in order for the data (with or without the OH-points) to fit eqn (8) (where AGi = AGO+ W,-W,) very large ;1/4 values are needed and the best fits give negative valuse for W and W,. Electrostatic attraction and repulsion have some influence as shown by table 4 where Alog k is the difference in log k for bases of charge 2 and the values expected (using the 0.5 Bronsted slope) for neutral bases of the same basicity.The anionic bases show much larger deviations than the cationic bases which is reasonable since the reaction site is not at the charge centres. The positive and negative bases should have different orientations relative to the metal ion with a greater distance between the metal centre and positive charged bases. TABLE5.-PARAMETERS (kJ mOl-') OF MARCUS THEORY acid base WR W'r 114 MLH2+ OH- RCO,' HzO 25 19 31 MLH2+ acac b OH- amines HzO OH- amines H20 -29(8) 39 -33(8) a 22 88(50) 26 acac b OH- RCO, H20 36 25 23 acac OH- RCO, H2O 44 38 14 RCO; amines C~(H-~glyglyhis)- 33 c= -1ld 5.9 The values in parenthesis are arbitrarily chosen to force WR and Wp to be slightly positive.The resulting fit has more curvature than the experimental data. b The Marcus parameters were calculated from the data of M. L. Ahrens M. Eigen W. Kruse and G. Maass Ber. Brmsenges. Phys. Chein. 1970 74 380. C M. M. Kreevoy and S. W. Oh J. Amer. Chem. SOC.,1973 95 4805. d C = (Wp-WR-AG&,). Fig. 3 gives the Bronsted plot for MLH2-'-reacting with OH- RCOO- and H20. This curvature can be fitted with the R/4 and WR values given in table 5. The 3,/4 values for the MLH2+ complexes are larger than for acac and the WR values are smaller. Several reasons can be suggested for larger A values including (1) the necessity of reorganizing more bonds when the metal is present (2) the need for greater electronic redistribution and (3) additional effects due to changes in the degree of axial solvation of the metal ion.Similarly several reasons can be suggested for lower W values including (1) the fact that MLH2+is already planar and needs less geometrical rearrangement than acac in order to initiate hydrogen bonding (2) the metal ion will tend to disrupt the solvent structure which might lead to less solvent reorganization as the hydrogen bonding is initiated. The relatively large rate constants for OH-with MLH2+ suggests the possibility that a metal hydroxide species could form prior to the proton transfer. We could find no evidence for a M(OH)LH+ species but the high OH-concentrations necessary to form it would cause rapid formation of MLf. However the trans diene complexes of Nil1 where proton loss from carbon atoms does not occur fail to form OH-comple~es.~A coordinated hydroxide ion would have reduced basicity and would require an intervening water molecule in order to accept a proton from the carbon atom.Hence the effectiveness of an M(OH)LH+ pathway is uncertain. Some acids and bases could have axial coordination to the metal ion with one atom and could transfer the proton with another atom. This may explain the particular effectiveness of HPOi-as a base and of H,PO as an acid. C. BANNISTER D. MARGERUM J. RAYCHEBA AND L. WONG PROTON TRANSFER REACTIONS OF METAL PEPTIDE COMPLEXES The lower part of the S-shaped Bronsted curves for Cu(H-,glyglygly)- and for Ni(H-,glyglygly)-is an anomaly which arises from considering H20 as a Bronsted acid when a second more favourable path is available if it acts as a Lewis base.Thus the solvent dissociation pathway for these species involves H,O replacement of the deprotonated nitrogen group from the metal followed by rapid solvent proton- ation of the -CON-'-) group. This pathway is observed for other metal-peptide complexes 2* lo and the general equation for the rate constant is given by eqn (14). kobs. = kd +kHBIHBl (14) Hence in the Bronsted plots we need to consider only why the slope changes from 1 to 0 before the rate reaches the diffusion limit. However a second difficulty arises as seen in the mechanism given in eqn (15) and (16) where M(Hm1G3)* is a reactive intermediate in which the peptide N has not yet moved away from the metal ion.This intermediate can react in ka MH-,G3 +HB + (MH-iGj)* +B k-s kb rapid (MH-1G3)*+B + MH-lG3B + products two ways with B as a Bronsted base (k-a)and with B as a Lewis base (kb). The resulting value of kHB is given in eqn (17) and we see that kHBis the actual proton-transfer rate constant only when k % kFa. In fig. 2 a11 the conjugate bases of HB can act as Lewis bases. When acids of non-complexing conjugate bases were tested they failed to give general acid behaviour with the triglycine complexes. This H FIG.4.-The copper(I1) complex of glycylclycyl-L-histidine,Cu(H-2glyglyliis)-. was the case for the Good buffers MES and PIPES as well as for 2,6-lutidine and 2-picoline with both Ni(H-,glyglygly)- and Cu( H-,glyglygly)-.With these bases k should be very small due to steric effects and an alternate mechanism is needed with H20 reacting as the displacing ligand so that only the H30+ rate constants are observed. All these buffers have pK values greater than 6. It is possible that more acidic HB species could give general acid behaviour. We cannot be certain from the trigiycine data alone if any of the rate constants in fig. 2 are actually due to proton- transfer steps although the HB species with pK < 6 might react in this manner. PROTON TRANSFER REACTIONS IN METAL COMPLEXES In order to avoid the above problem we examined the reactions of another peptide complex Cu(H-,glygly-L-his)- whose structure is given in fig. 4. Recent studies in this laboratory have shown that proton-transfer reactions can be made to be the rate- determining step when this complex is reacted in the presence of a high concentration of triethylenetetramine (trien).O Parallel conditions cannot be used for the triglycine 123456 lo2x [trien]~ FIG.5.-Observed first-order rate constants for the reaction of Cu(H-,glyglyhis)-with excess trien showing the approach to the limiting rates at high trien. Curve 1 no added general acid pH 7.5 calculated from resolved data. Curve 2 same conditions as curve 1 with 4.8 x loe3M H(HEPES)* added as HB pH 7.5. 7 6 5 m 24 M 23 2 I ' ' 11 8 7 6 I 5 I 4 I 3 I 2 I I I 0 I -I -2 P&(HB) FIG.6.-Bronsted plot for the proton-transfer dependent rate constants of Cu(H-,glyglyhis)-reacting with acids in the presence of excess trien.HB from left to right are Htris+ H(HEPES)* H(MES)* HOAc and H30+(p = 1.0 M 298 K). complexes because direct attack by trien as a nucleophile is too fast. However with Cu(H-,glyglyhis)- the nucleophilic reaction by trien is important only after a proton adds to the peptide group. Fig. 5 shows how the observed rate constants tend to level off as the trien concentration is increased. The effect of general acids is marked as seen by the higher plateau in the curve 2 run in the presence of 4.8 x lo3M H(HEPES)*. When reactions are run at high trien concentrations HEPES MES and Tris all accelerate the reaction as well as acetic acid and H30+. C. BANNISTER D. MARGERUM J. RAYCHEBA AND L.WONG Fig. 6 shows the variation of log kHB with the pKa(HB) value for these reactants. The mechanism where CuH-,L- is the glyglyhis complex is given in eqn (18)-(21) and (CuH-lL)* is a reactive intermediate with a proton present on one of the peptide nitrogens. (k,[H '1 +k;[H,]) k2[H2 trien '1 [CuH-,L-] rate = k- +k'_,[B-]+k2[Hztrien2+] At very high trien concentrations eqn (22) holds kobs. = kl[H+] +k',[HB]. Thus the presence of a large excess of a nucleophile removes the uncertainty about HB reacting in the proton-transfer step (H2trien2+ has a pK of 9.3 and does not contribute significantly as an acid). The H,O+ rate constant is 1.1 x lo7 M-l s-' f or Cu(H-,glyglyhis)- which agrees well with the value of 1.3 x lo7 M-l s-l for Cu(H-,glyglygly)- and the values close 1 to 107 ~-s-1 f or many other Cu(H-,tripeptide)- complexes.2 Therefore we conclude that all these reactions with H30+ are actually limited by the proton-transfer step.We propose that the curvature seen in fig. 6 is due to the situation described earlier where 2 has a small value and WR is relatively large. The small 2 causes gobs. to change rapidly as K becomes greater than unity and the kf value (kHB)levels off with large values of K(i.e. pKa(HB) = -1.74). The constants used to fit the curve in fig. 6 are given in table 5 A/4 = 5.9 kJ mol-l and W = 32.6 kJ mol-I. The reasons why the A/4 values are much smaller than the WRvalues are not clear but the degree of solvation of the reaction site may be involved. The behaviour of the deprotonated metal-peptides in intermediate between that of "normal " bases and the bases of carbon acids.CONCLUSION The rates of proton transfer reactions at the carbon atom of metal macrocyclic complexes and at the nitrogen atom in metal-peptide complexes are very different in their dependence on the ApK values of the reactants. In order to fit the data for either system the activation energy of the encounter species must have a term (Wi) which is independent of ApK as well as a term [1+ (AGi/A)I22/4 which depends on ApK,. The Wi term can be attributed to the solvent reorganization necessary to initiate hydrogen bonding after the work has been done to bring the reactants together in the encounter species. In the metal-peptide complexes this term is larger than the reorganizational barrier (A/4) and accounts for the limiting rates which are less than diffusion controlled.In the metal macrocyclic complexes A/4 is larger than W = (W{+AGY) but both terms change drastically with the type of base used. With the amine bases the MLH2+ Bronsted plots appear to be too linear. (The lack of curvature requires very large A/4 values which in turn means WRmust be negative to give the correct AG* values). The compensating nature of the 2/4 and Witerms makes predictions difficult. PROTON TRANSFER REACTIONS IN METAL COMPLEXES The work was supported by a National Science Foundation Grant and by Public Health Service Grant from the National Institute of General Medical Sciences. S. C. Cummings and J.G. Martin Inorg. Chem. 1973,95 1477. A review of the kinetics and mechanisms of metal peptide complexes is given by D. W. Margerum and G. R. Dukes in Metal Ions in Biological Systems vol. 7. ed. H. Sigel (Marcel Dekker New York N.Y. 1974) p. 157. M. Eigen Angew. Chem. Int. Ed. 1973,3 1. A. 0.Cohen and R. A. Marcus J. Phys. Chem. 1968,72,4249. R. A. Marcus J.Phys. Chem. 1968,72,891. M. M. Kreevoy and S. W. Oh J. Amer. Chem. Suc. 1973,95,4805. 'R. P. Bell The Proton in Chemistry (Cornell University Press Ithaca N.Y. 2nd edn. 1973 chap. 10 p. 195. * M. L. Ahrens M. Eigen W. Kruse and G. Maass Ber. Bunsenges.phys. Chem.,1970,74,380. F. P. Hinz Ph.D. Thesis 1973 (Purdue University). lo J. C. Cooper L. F. Wong D. L. Venezky and D. W. Margerum J. Amer. Chem. SOC.,1974 96,7560. N. E. Good G. D. Winget W. Winter T. N. Connolly S. Izawa and R. M. M. Singh Biochem. 1966 5 467.

ISSN:0301-5696    

DOI:10.1039/FS9751000078

出版商:RSC    

年代:1975

数据来源: RSC

摘要:

GENERAL DISCUSSION Prof. A. J. Kresge (University of Toronto) said In our BEBO calculations we varied AV both by changing p and holding V constant and by changing V while holding p constant. We find however that more important than changes in either of these parameters is the inclusion or neglect of end-atom repulsion. With this effect included we obtain sigmoid dependences of a upon AV which are fairly linear over their mid-portions (a= 0.2 to 0.8) and the (calculated) data for these regions are easily fitted to the simple Marcus formulation. The intrinsic barriers which these fits produce however are invariably less than calculated values of AE at AV = 0 often by factors as much as two and the work terms obtained are therefore corres- pondingly too large.I would like to add that although Marcus keeps both d1 and A constant in his model this does not necessarily correspond to empirical practice. For example along a series of proton transfers from carbon acids such as carbonyl or nitro compounds to a single base or even to a series of bases changes in AGO are commonly obtained by making structural variations in the substrate usually by changing the R. P. BELL oo 2 n \ z -Y El I c -30 -20 -10 0 10 20 30 AGO FIG.1.-Correlation of kinetic isotope effects for carbonyl compound ionization according to eqn (1) ; data from ref (1). extent of charge delocalization in the carbanionic product. Charge delocalization however is one of the more important factors governing the rate of proton transfer to or from carbon and these structural changes are therefore likely to alter d for the identity reaction between the carbon acid and its conjugate base.89 GENERAL DISCUSSION There is some experimental evidence to support this idea from isotope effect correlations using simple Marcus theory. The theory predicts that kinetic hydrogen isotope effects will be largest at AGO = 0 and will decrease symmetrically on either I BORDWELL AND BOYLE 0 0 -1 0 0 +I0 +20 AGO FIG.2.-Correlation of kinetic isotope effects for nitroalkane ionization according to eqn (1) ; data from ref. (2). side of this maximum at a rate controlled by AG the intrinsic barrier for the system being correlated small values of AGZ will give rapid decrease of kH/kD with AGO whereas large values of AG will lead to more gradual changes (eqn (1)).ln(k,/kD) = In(k,/kD),,,[ 1-(AG0/4AGZ)*]. (1) Experimental data for a group of carbonyl compound ionization reactions (fig. 1)' and a series of nitroalkane ionizations (fig. 2)2 do fitthis relationship reasonably well. But there are systematic deviations from the best (least squares) fits to all of the data in each case the points to the left of AGO = 0 cluster above the correlation line and those to the right of AGO = 0 concentrate below it. This suggests that AGZ changes along these reaction series and the direction of the change decreasing AG with increasing AGO is just that expected on the basis of more delocalization in the carbanions derived from the stronger acids required to make AGO < 0 than those produced by the weaker acids for whose reactions AGO > 0.Dr. W. J. Albery (Oxford University) said This morning Bell using Morse curves showed that in certain cases one could obtain maxima and minima in the variation of a with ApK (see fig. 3). Is there a simple explanation of the shape of these curves? R. P. Bell The Proton in Chemistry (Chapman and Hall London 1973) p. 265. F. G. Bordwell and W. J. Boyle Jr. J. Amer. Chern. SOC.,1975 97 3447. GENERAL DISCUSSION Prof. L. Melander (Giiteborgs Universitet) said It is pleasing to find that Marcus represents the longitudinal motions in the C1. . . H . . . Br system by means of the skewed potential-energy surface which is a very useful device but seldom used.In the present energetically very unsymmetric reaction the reaction coordinate in the transition state is almost parallel to one of the interatomic-distance axes. This means that the motion along that coordinate comes close to a mere approach between C1 and HBr i.e. a relative motion between two heavy species. In such a case the effective mass must be considerably heavier than the proton and tunnelling corres- pondingly negligible. (Another reason for tunnelling being negligible is the low potential-energy level of the saddle point relative to the reactants of course). Appreciable tunnelling could arise only when the reaction coordinate in the transition state has such a direction that the heavy neighbours of the hydrogen stay at almost constant distance from one another i.e.heavy-atom motion is negligible. I would be glad to know Marcus’ answers to the following questions Jn general could energetically very unsymmetric reactions be expected to have potential-energy surfaces as unsymmetric as the one presented ? In discussions of the variation of the “ classical ” isotope effect with symmetry it has been argued frequently that the (valence-bond force field) force constants between hydrogen and its next neighbours could hardly be expected to differ from each other by more than a factor of about ten. From fig. 1 it seems rather likely that the C1-H force constant could be even negative. Are such force constants likely to be applicable to strongly unsymmetric systems in general? Prof.R. A. Marcus (University of Illinois) said In response to Melander’s questions the answer is probably “ yes ” to the first one the exothermicity itself implies a large asymmetry about the bisector of the acute angle in fig. 1 of my paper. The experimental observation of high vibrational excitation of the reaction product HI in the gas phase reaction implies a large potential energy drop along the angular coordinate (the protonic coordinate) in the transfer region. Because of this large drop the answer to the second question should also be “yes” for a sufficiently exothermic reaction. Prof. R. P. Bell (University of Stirling) said The form of the Marcus relation frequently employed in treating proton transfers is based on a picture of intersecting parabolas or other types of potential energy curves.While this is reasonable for electron transfers hydrogen atom transfers (3-electron systems) certainly require a different approach such as BEBO. It is not clear however how far this approach can be transferred to transfers of protons or hydride ions which are 4-electron and 2-electron systems respectively and involve a net transfer of charge possibly the intersecting curve model is more appropriate here ? The form of the curves relating Bronsted exponents to AG” shown in fig. 3 of my introductory lecture depends on properties of the Morse energy curves on which they are based. These have a point of inflection at an energy of tD,where D is the dissociation energy and the sign of dp/dAG” at AGO = 0 depends upon whether the two curves intersect above or below this point of inflection.These results are not to be taken too seriously but serve to emphasize the fact that the theoreticaldependence of /3 upon AGO (and hence also values of AG; w and w deduced from experiment) do depend critically upon the model adopted for the energy profile. The consequences K. G. Anlauf P. E. Charters D. S. Horne R. G. McDonald D. H. Maylotte J. C. Polanyi W. J. Skrlac D. C. Tardy and K. B. Woodall J. Chem. Phys. 1970 53 4091. GENERAL DISCUSSION in this respect of an electrostatic model of proton transfer are at present being investigated. For a reaction of the type X+HY + XH+Y (charges not specified) the accepted procedure is to plot the energy surface with skewed co-ordinates as in fig.1 of Marcus’ paper and when X and Y are heavy entities the angle between the axes becomes very small. This procedure is undoubtedly correct when the species concerned are moving freely in one dimension but is questionable if the proton is being transferred within a reaction complex in which the reactants are already in a suitable position orientation and state of solvation. Under these circumstances the situation may resemble the transfer of a proton or hydrogen atom between two fixed centres in a solid lattice for which the energy surface becomes a function of a single linear co-ordinate. Melander has suggested that the validity of this last approach depends on how “stiff” the reaction complex is with respect to motions of the reactants and solvent molecules it would be valuable to have Marcus’ views on this question.Prof. M. M. Kreevoy (University of Minnesota) said The values of W‘ which have been obtained certainly have substantial uncertainties. It’s also easy to conclude that W‘ may be a weak function of pKHA. Nevertheless the general conclusion that W’ is much larger than the free energy required to assemble an encounter pair seems quite secure. That conclusion does not depend on the detailed form of the function used to fit the Bronsted plot. In the work I described today and in several earlier studies acids of increasing strength have been studied until a limiting rate has been reached or closely approached. The existence of such limiting rates far short of the diffusion rate is sufficient to establish that large values of Wrare required.Dr. R. A. More O’Ferrall (University College Dublin) said It is perhaps worth noting that Marcus’s expression must underestimate the “intrinsic ” barrier to reaction AG*’ when AGi’ = 0. Plotting AG* against AGif for the proton transfer step of an ideal reaction family / I FIG. 1.-Full line (ad) idealised plot of AG*’ versus AG;’ for proton transfer. Dashed lines best quadratic approximation to ad at AG;‘ =O (bb’); Marcus’s quadratic expression (cc’). M. M. Kreevoy and S. Oh J. Amer. Chem. Soc. 1973,95,4805. GENERAL DISCUSSION for which AG*' varies continuously with AG;' and -$ 0 and AG;' as AG;' + -and +infinity respectively gives the line labelled aa' in the figure which might be repre- sented by a hyperbola (E.S. Lewis unpublished) or by Marcus's BEBO expression. Near AG;' = 0 the best quadratic approximation to the line is given by the parabola bb'. However bb' differs from the quadratic expression used by Marcus AG"' = (1/4)(1 +AG;'/A)' which carries the implication that AG*' = 0 when dAG*'/dAGi' = 0 with the result that the parabola is shifted along the AG*' axis to cc'. Marcus's expression does not lead to an inferior correlation of results because what is plotted in practice is not AG*' but AG* = AG*'+wr so that low values of AG*' are compensated by a high value of w'. Nevertheless it is clear that A/4 must underestimate the barrier to the thermoneutral reaction and indeed if aa' is a hyperbola A/4 is low by a factor of 2 while if aa' is given by the BEBO expression it is low by 1.4.On the other hand this does not alter the qualitative conclusion emphasised by Kreevoy that a high Bronsted curvature associated with large activation energies implies a major contribution from resolvation to the reaction barrier. Prof. R. A. Marcus (University of Illinois) (communicated) In response to O'Ferrall and Kresge I would note first that the results of Kreevoy in finding large W"S from studies of the limiting rate at very negative AGO'S (fig. 1 of his paper) substantiate his earlier finding of large w' using a quadratic plot and measurements of AG" against AGO over a more restricted AGO range. Incidentally if one takes the limiting kHA/p there to be -3 M-l s-l and writes it as Zexp (-w'/RT) taking 2 the collision frequency in solution to be ca.10" M-I s-l then w' is about 62 kJ mol-l. This is only about 10 kJ 11101-' less than that deduced from the quadratic fitting to the data in fig. 1. As O'Ferrall and Kresge note evaluation of AG* from a restricted AGO range by fitting could lead to a model-dependent w'. However in a region of Bronsted slopes of 0.2 to 0.8 the In cosh BEBO expression (eqn (2.8) of my paper) is reasonably well approximated by the quadratic one (eqn (2.6)). For example when A/4 for eqn (2.8) is x kJ mol-' use of eqn (2.6) for AG* yields a new A/4 of 0.8~ kJ mol-l and a w' that is correspondingly 2x kJ mol-' larger than before. This difference in the wr's of these models is relatively small and helps explain the agreement of w"s in Kreevoy's results.Typically when obtaining w' from data over a restricted AGO range one might consider using both eqn (2.6) and (2.8) to explore the sensitivity of the deduced w' to the model equation. The extent to which any arbitrary curvilinear (AG* AGO) plot however can be approximated by a quadratic one depends on the shape of the former Use of an arbitrarily large (d3AG*/dAGo3)at AGO = 0 causes the fitting to differ by an arbitrarily large amount. Thus it is best to use experimental or theoretically-derived plots when comparing them with a quadratic one since otherwise any comparison could become quite meaningless. The quadratic plot appears to agree better with the In cosh one than with the plot in the figure in O'Ferrall's query.Prof. R. A. Marcus (University of Illinois) said Calculations of potential energy barriers have now been performed using harmonic BEBO and/or Morse-like potential energy surfaces by Bell,' Koeppl and Kresge and by me.3 Yet our conclusions R. P. Bell this Symposium Spiers Memorial Lecture. 'G. W. Koeppl and A. J. Kresge J.C.S. Chem. Cornrn. 1973 371. R. A. Marcus paper at this Symposium. GENERAL DISCUSSION as indicated in Bell’s comment are apparently so very different The first two sets of authors find that the Bronsted slope can differ considerably from 0.5 at AGO = 0 while I find that when A is held constant the slope is close to 0.5 at AGO = 0. Even for a highly asymmetric reaction in the sense that the A’s of the two reactants are very different (6 l) the slope at AGO = 0 is still close to 0.5 namely 0.6 in the present paper.In response to Bell’s comment I believe therefore that the difference in these findings lies in my use of a constraint on A rather than in choice of the form of potential energy surfaces themselves. For example a BEBO surface was used throughout the calculations in my paper rather than intersecting parabolas. In a previous paper it was shown that if A varies in a series of reactants for which AGO is varied the Bronsted slope should indeed be very different from 0.5 at AGO = 0 even negative or greater than unity in some cases. Unless special precautions are taken ,JL will vary when potential energy surface parameters are varied to vary the net potential energy change for the reaction AU.In particular variation of AU in a BEBO model by varying onZy the dissociation energies D1causes A to change. Accordingly both D,and the bond order coefficients pi were altered simultaneously in the present paper to alter AU at fixed A. The question arises therefore as to whether one should perform the calculations in a constrained way or not. The answer depends on the problem one wishes to examine. For example in proton transfers between an acid (base) having a large A and a series of bases (acids) having a small A the net ;1 for the reaction which is the mean of the two l’s should not vary appreciably from one member of the series to the next. In such cases e.g. for certain carbon acids reacting with oxygen or nitrogen bases which do not suffer extensive electronic rearrangement the conditions are best simulated by surfaces in which AU is varied holding A fixed.The fact that the experimentally observed Bronsted slope is indeed close to 0.5 at AGO = 0 for such highly asymmetric systems also supports this conclusion. Dr. D. M. Goodall (York University) said Marcus alluded to the possibility of facilitating proton-t ransfer reactions in hydrogen bonded systems by supplying vibrational energy using short laser pulses. Greenhow and I have already reported such an e~periment,~ in which liquid water was photoionized using a Q-switched neodymium glass laser. The quantum yield for this process was evaluated from the transient conductivity increase due to the excess hydroxyl and hydrogen ions produced.With Knight we have recently extended these experiments to excitation at other wavelengths. Quantum yields at 328 K are 1 x lo-’ (0.69 pm) 5 x lo-’ (1.06 pm) < 2 x lo-* (1.41 pm). Only an upper limit can be given at 1.41 pm where the transient conductivity profile is indistinguishable from that for microwave- heated water.4 We have shown that photoionization is a one photon process and anticipate that a theoretical account of this data will provide valuable information concerning the role of vibrational excitation in the activation process for proton- transfer reactions in solution. Prof. R. A. Marcus (University of Illinois) (commimicated) It was good to hear of these experiments of Goodall and of the new results.Of particular interest too would be measurements of the quantum yield 4 at still shorter wavelengths the R. A. Marcus J. Amer. Chem. SOC.,1969 91 7224. R. A. Marcus paper at this Symposium. D. M. Goodall and R. C. Greenhow Chem. Phys. Letters 1971,9 583. G. Ertl and H. Gerischer Z. Elektrochem. 1961 65,629. GENERAL DISCUSSION pre-exponential factor of the thermal liquid phase proton transfer rate constant of 2H,O -+ H30++OH- lo7 s-l corresponds to a AS* of -28 cal mol-’ K-I. If the deactivation frequency of the vibrationally-excited sys tem described by Goodall is of the order of 10l3s-l a collision frequency the maximum 4 would be 107/1013 i.e. at sufficiently high energies if the range of 0.. . 0 distances and solvent orientations over which the proton can transfer is the same as in the lower energy thermal system.If instead the range of these coordinates permitting proton transfer is greater in the photochemical case then the maximum 4 would exceed even with a collision deactivation frequency of lOI3 s-l. The experiments do not directly distinguish between statistical and non-statistical regimes of the unimolecular process whereas the lower energy picosecond experiments proposed in my paper would probably be in the non-statistical regime. Both types of experiments thereby complement each other. Dr. W. J. Albery (Oxford University) said I would like to present a somewhat different analysis of Kreevoy’s results which reinforces the point he has made that his conclusions do not depend upon the detailed algebraic formulation of the Marcus theory.In comparing the effects of changing the catalyst (HA) while keeping the diazo compound (S) the same with changing the diazo component while keeping the catalyst the same it would be helpful if we knew the dissociation constant Kf of the protonated diazo compound Unfortunately the values of K$ cannot be measured because SH+ decomposes. However we can assume that KF will obey a Hammett Relation and since Q depends on the dissociation constants of Ar COOH we may expect that ps > 1 log KF = log K:+pp. Now we relate the Bronsted coefficient as for changing S and keeping HA constant to the Hammett p values by writing where AGFD is the standard Gibbs free energy charge for the whole reaction and AG* kLA,and pHAare as defined in Kreevoy’s paper.Ignoring terms connected with Wp and assuming J. constant we may also write and as is the usual Bronsted coefficient for keeping S constant and varying HA. If awr equals zero then as = as and the system would be “ well behaved ”. Comparison of eqn (1) and (2) shows that -pHAshould vary linearly with aB. This is the equivalent of Kreevoy’s eqn (19) which does indeed lead to the linear variation shown in his fig. 3. However -PHA moves in the opposite direction to aB as shown by the following values -PHA aB H C Clz COOH 1.43 0.3 3,4,5-C13C6HzOH 1.30 1 .O cf. D. M. Goodall and R. G. Greenhow Chem. Phys. Letters 1971 9 583. W. J. Albery A. N. Campbell-Crawford and J.S. Curran J.C.S.Perkin 11 1972 2206. GENERAL DISCUSSION This implies that aWyris greater than unity which in its turn from eqn (2) means that for small values of aB,as is greater than unity. In terms of Kreevoy's pw and pc we can show that Ps = -(Pw+Pc) = 1.3 awr = I+pC/ps = 1.15 and for HCC1,COOH as = 1.10. As expected the value of ps is greater than unity but perhaps is somewhat small when compared to values for similar ionisations e.g. p = 2.8 for the ionisation of Ar N+H3;'the neglect of Wp terms may have led to the apparent low value. FIG.1.-Free energy profile for aliphatic diazo compounds showing the change in Wr plotted along the S co-ordinate going into the paper and the proton transfer going across the paper.The various compounds are of the type RIRzCNz where Rl and R2 are 1 H,[COO-; 2 COO- COO-; 3 Me COOEt ; 4 Me COMe; 5 COOEt COO-. The nitro-compounds studied by Bordwell amongst others are a similar system to the diazo compounds and it is interesting that Kreevoy's work demonstrates that like the nitro compounds the diazo compounds also have a value of cis greater than unity. The fact that aWr is close to unity confirms the variation of free energy with solvation and with the degree of proton transfer deduced by us for other aliphatic diazo compounds from an analysis of the degree of proton transfer in the H30+ catalysed reaction. Fig. I shows the pattern for five different diazo compounds which span a range of over 6 orders of magnitude in their rate constants.Similar " Marcusian " behaviour is found for the proton transfer for each compound across the diagram while the position of each compound on the solvation co-ordinate running into the diagram corresponds to awr 2i 1. F. G. Bordwell and G. D. Cooper J. Amer. Chem. SOC.,1952,74 1058. F. G. Bordwell paper at this Symposium. W. J Albery A. N. Campbell-Crawford and J. S. Curran J.C.S.Perkin ZZ 1972 2206. W. J. Albery C. W. Conway and J. A. Hall J.C.S.Perkin IZ 1976 473. GENERAL DISCUSSION Prof. W. H. Saunders (University of Rochester) said Would Kreevoy please comment on whether the encounter complex and reaction complex should be regarded as real or virtual intermediates? Could one be real and the other virtual? Bell and Goodall suggested that kH/kDwould reach a maximum value at ApK (or AGO) = 0.According to Marcus theory the quantity that should be zero is AGO within the reaction complex rather than the overall AGO. But observed kH/kD maxima usually occur within one or two pM units of ApK = 0 insofar as one can locate the maximum withreasonable precision. This fact suggests Wp-Wr < 2-3kcal/ mol. Although Wp and Wr might meet this requirement by being large but nearly equal another possibility is that both are fairly small. Prof. M. M. Kreevoy (University of Minnesota) said 1.The partners in an encounter complex are in no way different from completely separated molecules except that they happen to be close to one another. The process of reseparating them is simply their diffusion apart.Since diffusion does have a free energy of activation (albeit small) the encounter complex is a real intermediate. The reaction complex has a structure partially determined by the requirements of the transition state. That is the reaction complex must have a structure such that the transition state can have a minimum energy in all its normal modes except the reaction coordinate (including those of the solvsnt). Depending on the specifics of the solvent and the structure this may make the reaction complex correspond to a region of potential energy hyperspace which is not a local minimum but one in which the character of the atomic motions leading toward the products changes. The minimum energy path calculated for the transfer of a proton from NH; to NH3 shows such a reaction complex with the relevant coordinate changing from f" to rNH.' Such a reaction complex would be a virtual rather than a real intermediate.There may be other cases where the reaction complex corresponds to a local minimum in potential energy hyperspace. Even in those cases however it will probably be a very shallow minimum. As soon as the reaction complex begins to revert to the encounter complex normal solvation forces-and probably structural forces as well-will begin to work to lower its energy so that the free energy of activation for that reversion will probably always be much less than that for diffusion. 2. I think the kH/kDmaxima do not occur as close to ApK = 0 as Saunders implies. The two shown by Kresge in this symposium seen to have maxima around -3 and -4 kcal mol-l respectively.Caldin has referred to other results leading to the same conclusion. Further W' and Wp have several factors in common and in many cases would not be expected to be too different. I believe these results are entirely consis tent with the postulated large values of W'. Prof. E. F. Caldin (University of Kent) said (a) Although in a series of proton- transfer reactions a maximum isotope effect kH/kDis often found when AGO or ApK is zero (cf. Professor Bell's introductory lecture above) this is not necessarily to be expected when the series includes different types of base if tunnelling corrections are large. Variations in AGO (or equilibrium constant K) are usually influenced by the enthalpy of reaction AH"; but variations of kH/kDare greatly affected by the shape of the energy-barrier which does not depend solely on AHo but also on the barrier width and appears to be sensitive to the type of base.It is therefore not surprising that in our results (fig. 2 p. 127) the largest value of kH/kD is not at log K = 0. Nor is there a smooth relation between kH/kDand K;unpublished work by Dr. Rogne has P. Merlet S. D. Peyerimhoff and R. J. Buenker J. Amer. Chem.Soc. 1972 94 8301. S 10-4 GENERAL DISCUSSION shown that n-butyldiethylamidine has nearly the same equilibrium constant as tetramethylguanidine in its reaction with 4-nitrophenylnitromethane but a markedly higher value of kH/kD. These two bases both contain the imine grouping N€i=C/ -\’ the larger effect for n-butyldiethylamidine appears to be due to a higher barrier.(b) The effects of tunnelling on the parameters of the Marcus theory have not yet been worked out. Formally the theory can be expressed in terms of a potential- energy diagram composed of intersecting parabolae. For reactions where the tunnelling correction is important the rate will be affected not only by the energy terms for proton transfer and solvent reorganisation but very markedly by the width (or strictly the curvature) of the energy-barrier for proton transfer. This barrier- width will undergo changes if the positions of the parabolae are shifted either vertically or laterally by changing substituents or the solvent. The effects on the tunnelling correction and hence on the rate have not yet been examined; it seems likely that they will be significant.Dr. W. J. Albery (Oxford University) said Would Margerum like to comment on the fact that he has calculated Marcus parameters from Bronsted plots (fig. 1 2 3 and 6 of his paper) which include points at each extreme for catalysis by H20 and H30+? These two catalysts often deviate from the Bronsted plots for other catalysts ; for instance see point 1 in fig. 1 of Kreevoy’s paper.l Thus it may be unwise to calculate Marcus parameters from Bronsted plots including these catalysts. On the other hand if these catalysts are excluded from Margerum’s plots then it would appear that the remaining points do not cover a sufficient pK range for the Marcus parameters to be calculated at all.Prof. D. W. Margerum (Purdue University)said :Insofar as fig. 1and 3are concerned there is no reason to omit the data points for H30+,H20and OH-from the Bronsted plots since the points fall on the plots. The same is true for acetylacetone reactions where the Marcus parameters have been calculated by Kreevoy (see table 5 of our paper). It is amusing that one now needs to defend the fact that these points fail to fall below the Bronsted plot that is the deviant behaviour is now considered normal. However I would suggest that it may be unwise to omit these points in calculations of the Marcus parameters. The failure of H30+ H20 and OH-to fall on linear Bronsted plots is expected from the Marcus theory and it is not clear that one should arbitrarily drop these points.In the reactions of acids of differing strength with a given base the necessity to desolvate the acid can be interpreted in terms of a large WR term as suggested by Kresge.2 Our data with the macrocyclic complexes fail to show an anomalous desolvation effect for H30+ and H20 compared to the other acids used. In these reactions the WR values are of moderate magnitude and are comparable to the A/4 values. Hence it is difficult to justify the suggestion that the H30+points with the metal-peptide reactions are anomalous due to specific solvation effects for the reactions in fig. 2 and 6. In these reactions we find that WRvalues are the same order of magnitude as those found for acetylacetone and the macrocycles but the 2 values are much smaller.Kreevoy’s reactions were carried out in 80 %DMSO where the W values are much larger and where specific solvation of H30+ may be more important. For the metal-peptide reactions it is unfortunate that a wider pK range A. I. Hassid M. M. Krevoy and T. Laing paper at this Symposium. A. J. Kresge Chem. SOC.Rev. 1973,2,475. GENERAL DISCUSSION of other acids cannot be used but the experimental system precludes this. Our preference is to explain the observed behaviour in terms of the relative magnitude of Wkand A/4 rather than in terms of an exceptional behaviour for H30+which is not found for other reactions with nitrogen bases. Prof. D. W. Margerum (Purdue Uniuersity) said In our paper we point out the possible compensating nature of the A and WRvalues where large values of WRtend to occur with small values of L and vice versa.This is not unlike the compensating nature of AH" and ASo values and of AH* and AS* values which Bell has discussed and has indicated a preference for correlations of AGO (equilibrium constants) and of AG* (rate constants). The question which arises is whether or not it is possible to predict the relative magnitudes of WR and A for various systems. Prof. M. M. Kreevoy (University of Minnesota) said Large values of W' can be obtained by an appropriate choice of solvent and/or by steric obstruction of the transfer site. Solvents which interact strongly with one of the reagents in such a way that the interaction must be substantially disrupted in order to form the transition state should lead to large values of IT".The DMSO component of the solvent mixture used for our work on diphenyldiazomethane was chosen partly because it was thought that it would increase the strength of the hydrogen bonds from the acids to the solvent. I believe this is one of the reasons for the unusually high value of W' observed. For reactions involving Hf anhydrous DMSO as solvent should give particularly large values of W' because in that solvent H+ probably has the structure \s-0 . . M . . . 0-s /+ * ,2 / \ and one of the strongly bound DMSO molecules has to be removed in order to allow a reaction to occur. In support of this idea the rate constant for protonation of tribenzylamine by H+ in anhydrous DMSO has been found to be only about 1 104 ~-s-1.3 R. P. Bell The Proton in Chemistry (Cornell University Press Ithaca New York 2nd Edn. 1973). p. 79. M. M. Kreevoy and J. M. Williams J. Arner. Chern. Soc. 1967 89 5499. Y. Wang unpublished results.

ISSN:0301-5696    

DOI:10.1039/FS9751000089

出版商:RSC    

年代:1975

数据来源: RSC