摘要:

DISCUSSIONS OF THE FARADAY SOCIETYNO. 39 1965The Kinetics ofProton Transfer ProcessesTHE FARADAY SOCIETYLONDONDistribution arrangements overleaTHE SOCIETY’S PUBLICATIONSDiscussions of the Faraday Society Normally published twice a yearMEMBERSof the Faraday Society receive current issues of bothTransactions and Discussions free on publication.Enquiries regarding membership of the Societyshould be addressed to: The Secretary,The Faraday Society, 6 Gray’s Inn Square,London WCl (Telephone: Chancery 8101)NON MEMBERSmay obtain the Society’s publicationseither through their own bookselleror by making application as follows :FORAnnual Subscriptions: to current issues of E IT H E R Transactions and DiscussionsOR Transactions onlyBack Issues: Complete Volumes (comprising Transactions and Discussions)from Vol.41 (1945) onwardsAPPLY TOThe Aberdeen University Press LtdFarmers Hall, Aberdeen, ScotlandAND FORDiscussions: All current issues and back numbersBack Issues: Complete Volumes (comprising Transactions and Discussions)from Vol. 1 (1905) to Vol. 40 (1944)APPLY TOButterworth & Co. (Publishers) Ltd88 Kingsway, London WC2OVERSEAS ADDRESSES..Qustralia: Butterworth & Co. (Australia) Ltd.Sydney: 20 Loftus StreetCanada: Butterworth & Co. (Canada) Ltd.Toronto: 1367 Danfoth Avenue, 6New Zealand: Butterworth & Co. (New Zealand) Ltd.South AJrica: Butterworth 6( Co. (South Africa) Ltd.Cl/ellington: 49!5 1 Ba I I aiice StreetDurban: 33!35 Deach GroveU.S.A.: Discussions-current and back issues :Butterworth Inc.Washington, D.C.20014: 7300 Pearl StreetComplete Volumes 1-40:Johnson Reprint CorporationNew York: 111 Fifth Avenue. A GENERAL DISCUSSIONONThe Kinetics ofProton Transfer Processes12th, 13th and 14th April 1965A GENERAL DISCUSSION on the Kinetics of Proton Transfer Processeswas held at the University of Newcastle-upon-Tyne on the 12th, 13th and 14thApril, 1965. The President, Prof. F. S. Dainton, M.A., Sc.D., F.R.S., was inthe Chair and 200 members and others were present. Among the visitors from over-seas were :Mrs. M. L. Ahrens W. GermanyDr. P. Ausloos USADr. D. M. Brouwer HollandProf. S. G. Christov BulgariaProf. B. G. Conway CanadaDr. P. Courvoisier FranceDr.and Mrs. G. P. Cunning-ham GermanyDr. M. Eigen GermanyDr. H. Eisenberg IsraelProf. E. U. Franck GermanyDr. K. H. Grellman GermanyDr. S. Gordon USADr. P. Goudmand FranceProf. E. Grunwald USADr. Y . Haven NetherlandsProf. Dr. E. Havinga Nether-Dr. X. de Hemptinne, BelgiumProf. W. Jaenicke GermanyMiss M. Kasparian LebanonDr. M. M. Kreevoy USAProf. and Mrs. A. J. Kresge USADr. K. J. Laidler CanadaMrs. S. G. Lias USAMr. C . U. Linderstrom-LanglandsDenmarkProf. F. A. Long USADr. W. Van der Lugt Nether-Dr. C. MacLean NetherlandsDr. L. De Maeyer GermanyDr. S. J. Magram GermanyProf. M. Mandel NetherlandsDr. D. B. Matthews USADr. G. M . Meaburn FranceProf. J. Metzger FranceMr. S. 0. Nielsen DenmarkMr. H. Nord DenmarkProf.and Mrs. R. Noyes USADr. H. W. Nurnberg GermanyDr. G. Olofson SwedenDr. W. L. Reynolds USADr. and Mrs. E. B. RobertsonDr. P. M. Sorgo SwitzerlandMr. and Mrs. K. SteigenvaldDr. P. Stonehart USAProf. H. Strehlow GermanyDr. G. Szasz SwitzerlandDr. W. J. Wallace CanadaProf. A. Weller NetherlandslandsGermanyGerman0 The Faraday Society and Contributors 1966Printed in Great Britain at the University Press, AberdeeCONTENTSPage 71625364567758494105112121130136149General Introduction : Kinetics of Proton Transfer Processesby M. EigenIsotope Effects and the Nature of Proton-Transfer Transition Statesby R. P. BellQuantum- Mechanical Tunnelling and the Dimensions of Energy-barriersin Proton-transfer Reactions in Solutionby E.F. Caldin and (Miss) M. KasparianProton Transfer Reactions Occurring in the Gas-Phase Radiolysisby P. Ausloos and (Mrs.) S. G. LiasGENERAL DIscussIoN.-Prof. J. J. Weiss, Prof. F. A. Long, Prof. A. J.Kresge, Prof. M. Eigen, Prof. B. E. Conway, Mr. R. P. Bell, Dr. D. B.Matthews, Dr. N. A. J. Rogers, Dr. M. Fleischmann, Dr. H. W. Nurnberg,Dr. Roger Parsons, Prof. M. C . R. Symons, Prof. M. M. Kreevoy,Dr. J. R. Jones, Dr. J. R. Hulett, Prof. G. J. Hills, Prof. S. G. Christov,Dr. E. F. Caldin, Dr. M. J. Henchman, Dr. R. A. ROSS, Dr. P. Ausloos.Kinetics of Proton Transfer to Weak Aromatic Basesby B. C. Challis and F. A. LongGeneral Acid Catalysis in Moderately Concentrated Aqueous SulphuricAcid by A.J. Kresge, L. E. Hakka, S. Mylonakis and Y. Sat0Proton Transfer to Olefins by V. Gold and M. A. KessickGENERAL DIscussIoN..Mr. R. P. Bell, Prof. F. A. Long, Prof. V. Gold,Prof. A. J. Kresge, Dr. M. Spiro, Prof. M. M. Kreevoy, Mr. B. Case,Dr. R. Parsons, Prof. M. Eigen, Prof. H. Zollinger, Dr. A. Gandini,Prof. P. H. PleschSolvent Particiyation in Proton Transfer Reactions of Amines and theirConjugate Acids by E. Grunwald and M. CociveraAcid Catalyzed Hydration of Acetaldehydeby M.-L. Ahrens (Mrs.) and H. StrehlowKinetics of I ,2-Hydrogen Shifts in Carbonium Ionsby D. M. Brouwer, C. MacLean and E. L. MackorGENERAL DIscussIoN.-Prof. M. Eigen, Prof. R. M. Noyes, Prof. B. E.Conway, Mr. R. P. Bell, Prof. M. M. Kreevoy, Prof. A. Long, Dr. C. F.Wells, Prof.Dr. H. Strehlow, Prof. R. J. Gillespie.Influence of Electric Field in the Double Layer on Rate Constants Deter-mined with Electrochemical Relaxation Techniques for Fast HomogeneousProton Transfer Reactions in Solution by H. W. NiirnbergExamination of Proton Transfer Reactions by Temperature Jump andElectrochemical Methodsby A. Bewick, M. Fleischmann, J. N. Hiddleston and Lord Wynne-Jones159146172176183194200207216223239253278CONTENTSGENERAL DISCUSSION.-D~. W. J. Albery, Dr. H. W. Nurnberg, Dr.M. Fleischmann, Prof. M. Eigen, Dr. A. BewickEstimation of Very Fast Reaction Rates porn the Broadening of’ Vibra-tional Spectral Lines by M. M. Kreevoy and C. A. MeadRate of Proton Transfer in Strong Acids and Raman Line Broadeningby A.K. Covington, M. J. Tait and Lord Wynne-JonesGENERAL DIscussroN.-Dr. A. K. Covington, Dr. R. E. Weston, Jr.,Prof. M. M. Kreevoy, Mr. T. H. Lilley, Prof. R. J. Gillespie, Prof.B. E. ConwayEffect of Solvent and Temperature on Proton Transfer Reactions ofExcited Moleculesby H. Beens, K. H. Grellmann, M. Gurr and A. H. WellerProton Transfer during Reactions in the Excited Stateby T. S . Godfrey, G. Porter and P. SuppanProton Mobility in Water at High Temperatures and Pressuresby E. U. Franck, D. Hartmann and F. HenselProton Migration in Aqueous Solutionby G. J. Hills, P. J. Ovenden and D. R. WhitehouseGENERAL DIscussIoN.-Prof. M. C . R. Symons, Prof. B. E. Conway,Prof. B. E. Conway, Prof. R. M. Noyes, Prof. M. Eigen, Dr. K. E.Johnson, Dr. G. Kohnstam, Dr. H. Eisenberg, Prof. E. C . Baughan,Prof. Dr. M. Mandel, Prof. G. J. Hills, P. J. Ovenden, D. R. Whitehouse,Dr. P. A. H. Wyatt, Prof. R. J. GillespieClassical and Q uan t um- Mechanical Efec ts in Elect ro c h em ica 1 Pro tonDischarge and the Kinetics at Low Temperaturesby M. Salomon and B. E. ConwayProton Transfer Across Double Layers. Mechanism Evaluation fromIsotopic Effectsby J. O’M. Bockris, S. Srinivasan and D. B. MatthewsGENERAL DIscussIoN.--Prof. V. Gold, Prof. K. J. Laidler, Prof. B. E.Conway, Prof. S. G. Christov, Dr. M. Salomon, Prof. J. O’M. Bockris,Dr. S. Srinivasan, Dr. D. B. Matthews, Dr. H. W. Nurnberg, Prof.G. J. Hills, Dr. R. ParsonsAUTHOR INDE

ISSN:0366-9033    

DOI:10.1039/DF9653900001

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Kinetics of Proton Transfer ProcessesGeneral IntroductionBY M. EIGENMax-Planck-Institut fur physikalische Chemie, GottingenReceived 4th May, 1965Proton transfer processes-although they have never appeared in the titles-have repeatedly been the subject of Discussions of the Faraday Society. Theconference on Homogeneous Catalysis held at Cambridge University in 1928 wasalmost entirely dedicated to the catalytic action of proton donors or acceptors,i.e., acids and bases according to Bronsted's definition. One of the main concernsof the speakers at that meeting was the quantitative establishment of the fact thatthe catalytic effect of acids and bases is not an exclusive property of the hydrogenor hydroxyl ion-a view which was first adopted by Ostwald. Thus, the highlightof the meeting was Bronsted's paper on The Theory of Acid and Basic Catalysisin which he gave a quantitative account of the catalytic effects of acids and basesin relation to their dissociation constants.On reading these papers, it is inter-esting to note how close our modern views come to the original concepts. In themeantime these were partly obscured because most data were specifically selectedas a consequence of practical limitations. Experimental limitations may alsoexplain the fact that the chemists at that time turned their interest more towardsan investigation of complex reaction mechanisms rather than towards the study ofthe more simple elementary steps involved in acid-base catalysis.The success in establishing and categorizing such more complex mechanismswas reflected in another famous meeting, the discussion of Mechanisms andChemical Kinetics of Organic Reactions in Liquid Systems held at London in 1941.The illuminating papers given by Ingold, Hughes and others certainly represent alandmark in establishing the discipline of " physical organic chemistry ".The efforts of physical chemists to penetrate the mechanism of catalysis by study-ing the elementary steps were thwarted until suitable methods for the study of rapidreactions were found.Again it was the Faraday Society who opened the discussionof such possibilities at their Birmingham meeting on The Study of Fast Reactionsin 1954. Almost no experimental data were available at that time, but the expecta-tions expressed by Bell in his introductory remarks that chemists as a whole maytake advantage of the new methods are now apparently being fulfilled, especiallyin the field of protolytic reactions.Almost all simple proton transfer processesare rapid. Thus, many of the experimental results which will be reported at thismeeting could only be obtained by use of the techniques which were introduced atthe 1954 meeting. These include :THE SAMPLING TECHNIQUES such as isotope exchange [papers by Bell, Ausloos andLias, Challis and Long, Kresge et al., Gold and Kessick] and nuclear mag-netic resonance [papers by Grunwald and Cocivera, Ahrens and Strehlow,Mackor and co-workers],8 KINETICS OF PROTON TRANSFER PROCESSESRELAXATION OR PERTURBATION TECHNIQUES such as temperature, pressure and electricfield pulse, as well as sound and dielectric absorption techniques [cf.belowand papers by Bewick, Fleischmann, Hiddleston and Wynne-Jones] andflash photolysis or radiolysis Cpapers by Ausloos and Lias, Porter and co-workers],FLUORESCENCE TRANSFORMATION AND OPTICAL LINE BROADENING Cpapers by Wellerand co-workers, Kreevoy and Mead, Wynne-Jones and co-workers], andTHE ELECTROCHEMICAL METHODS [papers by Nurnberg, Bewick et al., Salomon andOther papers in this discussion deal with the behaviour of protons under extremeconditions such as :Conway Bockris, et aZ.1.HIGH TEMPERATURE AND PRESSURE [papers by Franck and co-workers, Hillis et al.]andLOW TEMPERATURE [paper by Caldin and Kasparian].Measurements of this type, especially when carried out under such a wide rangeof external conditions as reported by Franck and co-workers, give a much broaderbasis for theoretical considerations (energy surfaces, tunnel effect, etc.) than theclassical data, which were usually restricted to the liquid state under normal tem-perature and pressure conditions.However, there is no longer any justification for categorizing these papers onlyaccording to their technical approach.The techniques are sufficiently well estab-lished that we no longer merely have to look for " applications ". Rather, we maynow start from the chemical problem itself, and use at will any technique that isdesirable. It is this approach to the problem which will be at the centre of ourdiscussion, and to which I shall here add a few general remarks.*NATURE OF PROTON TRANSFER PROCESSESThe rate of proton transfer is decisively determined by the distance betweenthe donor and acceptor group at the moment of transfer.In the classical picturethis distance will have a great influence on the activation energy since it will deter-mine how well the potential curves along the reaction co-ordinate overlap. Fora tunnelling mechanism the rate will be even more sensitive to this distance. There-fore the formation of an H-bond, providing an optimal overlap of the potentials,is a most important prerequisite for a fast proton transfer. Furthermore it minim-izes any rearrangements of the heavier atoms. Restrictions of the Franck-Condontype are therefore much less important for proton- than for electron-transfer re-actions.Experimental data 1 9 2. show that proton transfer processes can becomequite slow when interference with the H-bond formation between the donor andacceptor group occurs. One would expect that following the tendency to H-bondformation, the rate of proton transfer parallels the series :O H . . . O > O H . . . N, N H . . . O > N H . . . N > S H . . . X, X H . . . S >P H . . . X, X H . . . P > C H . . . X, X H . . . CIn aqueous solutions the rates of the reactions with H3O+ and HO- have been* For a more detailed discussion reference is made to some recent reviews.ls 2(charge and steric effects may lead to certain distortions)M. BIGEN 9studied for a great variety of compounds including almost all well-known typesof inorganic and organic acids and bases.As is shown by comprehensive tablesin ref. (1) and (2), in most cases the rate constants of recombination (H3O+ + base,or HO- + acid) approach the limiting values for diffusion-controlled reactions (1010-1011 M-1 sec-1). They show certain influences of charge, steric requirements, solvent-H-bonding, etc., which have been discussed in detail.17 2 Where larger deviationsfrom this theoretically well understood behaviour occur, they can usually be relatedto more complex reaction mechanisms involving tautomeric or isomeric changes.Since H30+ is a very strong acid and HO- a very strong base, in most cases an ap-preciable gain of free energy is connected with the proton transfer : pK~,o+<pK~xor p K ~ , o % ~ K H x .If this condition is not fulfilled, i.e., if the pK of the acid formed-3 - 2 - I 0 I 2 3-3 - 2 -1 0 I 2 3FIG. 1 .-Theoretical dependence of log k (rate constant for proton transfer) on ApK (pK-differencefor donor and acceptor) assuming maximum rates (i.e., no potential barrier in favourable direction).rate constants are normalized by the rate Constan-kD for a diffusion-controlled process.Reactions including no charge neutralization :(e.g., XH+Y-+X-+YH,or XH++Y+X+YH+).Reactions including charge neutralization :(e.g., XH++Y-+X+HY: ZD>ZD).in the reaction is not higher than that of H3O+ (example : protonation of carbonyl-groups of ketones and aldehydes), the reaction can no longer be diffusion controlled,but-at least when oxygen is the acceptor-the reverse reaction might be.Aslong as the reactions are diffusion controlled-as most reactions of H3Of and HO-are-they provide little information about the nature of the proton transfer, sincethis step is not the rate-limiting one10109654KINETICS OF PROTON TRANSFER PROCESSES- 4 - 3 - 2 - 1 0 I 2 3 4 5 65* M0 -0-5APK(b)FIG. 2.-Measured dependences of log k (rate constant for proton transfer) on ApK. (pK-differ-(a) Upper curves : phenol with different acceptors (around pK 10) ; lower curves : thyoglycole(b) Acetylaceton (ketonic) with different acceptors (PK values between -2 and 16) ; cf. alsoence for donor and acceptor.)with different acceptors (around pK 10)M.EIGEN 11More information about this step may be obtained from studies of proton transferreactions in those acid-base systems which show only small pK-differences for thedonor and acceptor. Fig. 1 represents what is to be expected if the rates are maximal.For a pK-difference of 0 the rates in both directions are equal and half of themaximum value. For a positive pK-difference (increase of pK in direction ofthe transfer : pKm > pKm for XH + Y -tX + HY) the rate approaches the maximumvalue and hence becomes independent of the pK-difference (for acids and bases otherthan H30+ and HO- k is usually 109-1010 M-1 sec-1) whereas for the reverse reactionlog k depends linearly on the pK-difference. Again, charge and steric effects (alsointernal H-bonding) might lead to certain distortions (cf.fig. 121).Some representative examples of experimentally determined dependences forthe different types of acid-base systems (studied with a great variety of differentdonors and acceptors) are shown in fig. 2. Here the rates show the above-mentionedbehaviour. For certain transfer processes of the type OH . . . 0 the rates approachthe ideal behaviour according to fig. 1 throughout the whole ApK-range. ForNH . . . N-systems deviations around ApK=O become perceptible. For SH-compounds these deviations become quite pronounced and for CH-compoundsthey extend over a very large ApK-range. However, even for these systems thegeneral character of the curves is still maintained, i.e., log k becomes independentof or linearly dependent on ApK for extreme values of ApK.According to somemeasurements of Grunwald et al. and Luz and Gill 3 one may expect the curvefor phosphine-compounds to come below that of sulf hydril-compounds. Further-more, recent studies of the properties of the solvated electron4 in water allows anestimation for the behaviour of its conjugate acid: the H-atom (which is an acidcomparable in strength to phenols). This curve would fall in between the -SHand -CH-compounds. There is no single curve describing all the different CH-compounds and the larger the deviation from the ideal curve in fig. 1 the larger therange of scatter (the curve in fig. 2 refers to the group of aliphatic ketones). Evenfor OH-compounds one can still distinguish several classes of curves.The idealbehaviour is only shown for " hard " acids and bases where the charges are con-centrated at the donor or acceptor site. The presence of resonance stabilization ofthe acid or base with respect to its conjugate compound introduces small butperceptible deviations from the ideal curve.ACID-BASE CATALYSISThe behaviour of the different classes of acids and bases as shown in fig. 2 shouldgive us a key for an understanding of the mechanisms of acid-base catalysis. TheBronsted relation, which describes the dependence of the log of the rate constantof catalysis on the pK of the catalyzing acid or base respectively, is nowadaysusually known as a linear relationship of the form log k = a (pK)i-c, with a and cbeing constants (c might include an individual statistical correction for multifunctionalgroups).It should be emphasized, however, that Bronsted originally expected abehaviour as shown in fig. 2, is., a continuous variation of a between 0 and 1.The linear relation (resulting from a McLaurin expansion) should then hold only fora limited pK-range. The limitations of the time-range of classical techniques ledto the fact that mostly only those rates were measured which are sufficiently farfrom diffusion controlled and where a remains constant over a wide pK-range.There is no reason to assume that a really changes continuously with pK. Fromsimple models of potential curves one might even expect that a remains almostconstant in a relatively wide pK-range. However, this can only be true fo12 KINETICS OF PROTON TRANSFER PROCESSESprocesses in which there are high barriers of free energy of activation in both directions.In this range the overlapping branches of the potentials can be approximated bylinear functions which is only possible if the region of overlapping is sufficientlyfar from the bottom of the potential curves.Examples of a continuous variationof a with ApK are provided by certain keto-enol changes which have been studiedby Bell and his school 5 using the bromination technique. It was found that for aseries of related compounds the observed value of a changes with the substrate pK,whereas for each substrate there was a constant a in the limited pK-range of thecatalyst which was studied. Relaxation studies in a wider pK-range show thatmost curves for different substrates and catalysts if plotted as a function of ApK,can be combined in a single curve with an a continuously varying from 0 to 1 (cf.acetylacetone in fig.2). More constant U-values in a wider pK-range have beenfound by isotope exchange studies, as reported at this Discussion by Challis andLong, and Kresge et al. However, the absolute values of the rate constants are muchlower than in the cases mentioned above, so the range of validity of the linear ap-proximation in the McLaurin expansion might here include the whole pK-rangestudied.There is a great number of thoroughly studied6 reactions which require botha catalyzing acid and a base (examples: mutarotation of glucose, hydration-dehydration reactions, certain hydrolysis reactions and other prototropic changes).Their general mechanismHS+HA +B+SH+ B+HAis often written in the form of a stepwise mechanism :HS + HAgHSH + BB + HSH+HA + SH(HA being the acid catalyst and HS the substrate).In such a formulation one has to assume that one of the protonation or de-protonation processes is the rate-limiting step.Thus the overall rate constant kmight consist of the product of an equilibrium constant and a rate constant k*.If one plots this rate constant k* as a function of pK one often encounters a constanta< 1 extending over a large range up to k*-values of 109-1010 M-1 sec-1, or even rateconstants exceeding the limiting values for diffusion-controlled reactions which is incontradiction to what has been found for other simple proton transfer processes.The only way out of these difficulties is the formulation of a concerted or co-operativemechanism where both the acidity of the catalyst HA and the basicity of its con-jugate base B come into play at the same encounter. In aqueous solution one canformulate this easily with the help of one or several H20 molecules (having thebifunctional effect of an acid and a base). Lowry in his introduction to the 1928meeting, has already emphasized the importance of H20, I quote :rapid ++“ In aqueous solution the action will generally be :base + HS + HOH +acid + SH + OH(-)OH2 + HS + acid+OH(i) + SH + base orwhere the water plays the part of an acid when the catalyst is a base, or of a basewhen the catalyst is an acid.”With our present knowledge we can now say that the mechanism must involveboth processes at the same encounter, e.g.13H - - - - ‘ H be A\o’- - - .- HA -./ - / # - - - - O - - - - HA __c s’ S‘ - - - c- -H’ ‘H. . . _ . ”possibly including several water molecules in the ring.Such a co-operative mechanism would be favoured since it avoids the solvationand desolvation of any charged intermediate. Restrictions of the Franck-Condontype impose a certain structure on the substrate-solvent-catalyst complex. Thisstructure (and therefore the whole concerted mechanism) may be disfavoured ifone (or both) of the donor-acceptor sites include a non-H-bonding group (e.g.,CH).Here the stepwise mechanism might be more effective.16 I4 12 10 8 6 4 2 0 -2- log10 (qK4d.P)FIG. 3.-Bronsted plot of the catalytic rate constants [M-1 sec-11 of various acids in the dehydrationof acetaldehyde hydrate according to Bell and Higginson.8 Solid line: measured rate constantagainst PKA of catalyst (including statistical correction factors p and q). This linear relationshipis obeyed by more than 50 substances showing a mean deviation of 0.15 log units throughout thewhole pK-range (note measured point for H20 which shows positive deviation). Broken line:rate constant k* for ratelimiting step (deprotonation) assuming a stepwise mechanismrapidHS + HA+HS+H+BHS+H+B +SH+HA.For p K ~ s @IS = hydrated acetaldehyde) a value of -2 has been chosen arbitrarily.(Note thatthe rate constant k* for HO- as a base would be above diffusion controlled and that no chang ofIX according to fig. 2 is present.14 KINETICS OF PROTON TRANSFER PROCESSESIn an H-bonded substrate-solvent-catalyst complex the potential surfaces alongthe path of proton transfer have very similar forms for a given substrate and a varietyof different catalysts, resulting in a persistently constant a over a wide pK-range.(The rate constants for the concerted transfer are still far from diffusion controlled ;in other words, o! must tend to 0 if the rate constants approach these limits.)In the concerted mechanism the catalyst merely triggers the reaction, and itselfremains essentially unchanged. This does not exclude an intermediary protontransfer from or to the catalyst. However, the lifetime for the intermediate shouldnot exceed the time necessary for orientation or disorientation of the solvent mole-cules, i.e., the time for the solvation of the intermediates.Similarly, the term“ concerted ” does not necessarily require a concerted motion of all the protonsinvolved, but means a correspondence between these motions within times of< 10-10 sec. This correspondence might be quite strong since the intermediatestates are not solvated. Further experimental and theoretical studies are requiredto resolve all the couplings and normal modes in the transition complex whichare effective for the co-operative proton transfer.In aprotic solvents the reaction requires bifunctional catalysis as Swain andBrown have clearly demonstrated. If the catalyst itself is not bifunctional bothan acid and a base in a ternary complex with the substrate are required.(Athorough discussion of these experiments is given in ref. (5).) The bifunctionalrole of water is also expressed in Grunwald’s experiments on proton transfer be-tween an acid and its conjugate base via H20 by the finding of a pK dependencecorresponding to an “ a-value ” between 0 and 1. A more detailed descriptionof such concerted proton transfer processes, which are of great importance for anunderstanding of the catalytic properties of enzymes, is given in a forthcomingpaper?CONCLUSIONSWe have seen at least the possibility of an understanding of the nature of acid-base catalysis in terms of the elementary steps of the reactions involved.A numberof such steps has already been identified and categorized and we might expect thatthis procedure is applicable to any type of reaction. It should then be possiblenot only to predict a certain catalytic mechanism but also to define its optimalconditions. This might also bring us close to a solution of the problem of enzymeswhose turnovers often include concerted internal acid-base catalysis (hydrolases,etc.). At present we can only give upper limits for a most favourable mechanismof this kind and-surprisingly enough-the measured turnover-numbers comefairly close to these figures. Other applications include a revision of certain ruleswhich have been proposed for the kinetics and mechanisms of organic reactionsincluding prototropic changes (such as keto-enol tautomerism and others29 7.On the other hand, while we now have some understanding of the relations be-tween rates and structure, many of the details are still obscure.What influencesdetermine the potential energy surface in a proton transfer reaction? What isthe nature of the activated state? Can we find any correlation of the isotope effectswith the variation of the Bronsted coefficients? Both quantities should be relatedto the form and height of the potential barrier. As Bell points out in his paper,there is no doubt that systematic experiments with isotopes will bring us moreinformation about these questions. The same may be true for the study of protontransfer processes at electrodes, where one can change the variables continuouslyin a defined manner. However, any progress in this field depends largely on abetter understanding of the metal/solution interface propertiesM. EIGEN 15I might close by expressing the hope that this conference will live up to the standardso excellently defined by Ingold in his introductory remarks to the 1941 discussions :“The Council of the Faraday Society has always shown a remarkable facility forselecting subjects for discussion in so timely a manner that the discussions them-selves do not merely record chemical history but make it. ’’la Weller, Prog. React. Kin., 1961, 1, 189.1b Eigen, Kruse, Maass, De Maeyer, Prog. React. Kin., 1964, 2,285.ZEigen, Angew. Chern., 1963, 75, 489; int. edn., 1964, 3, 1 ; (part TI on “ Mechanisms of3 For references cf. Grundwald‘s paper in this Discussion.4 Matheson and Rabani, in press ; cf. also Rad. Research, 1963, 19, 180 ; 1964, 4, 1.5 Bell, The Proton in Chemistry (Cornell University Press, New York, 1959).6 Bell, Acid-Base Catalysis (Oxford University Press, Oxford, 1941).7 Ingold, Structure and Mechanism in Organic Chemistry (Cornell University Press, New York,8 Bell and Higginson, Proc. Roy. SOC. A , 1949, 141, 197.Catalysis ” in preparation).1953)

ISSN:0366-9033    

DOI:10.1039/DF9653900007

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Isotope Effects and the Nature of Proton-TransferTransition StatesBY R. P. BELLPhysical Chemistry Laboratory, OxfordReceived 12th January, 1965The magnitude and variability of hydrogen isotope effects in proton-transfer reactions indicatesthat some of the modes of the transition state must involve considerable motion of the proton.These may include stretching vibrations (especially when the process involves the movement ofother atoms in the system), bending vibrations, or the " vibration " of imaginary frequency whichleads to reaction : in the last case the tunnel correction is isotope-dependent. There will be somecancellation between the contributions to the isotope effect of transition state bending and of thetunnel correction. It is suggested that non-equilibrium solvation of the transition state may alsocontribute.As shown by the contributions to this discussion, there is now ample evidencethat many physical and chemical processes involve the transfer of a proton betweentwo atoms.The present paper attempts to speculate about what information maybe obtained about the transition state, especially by the study of kinetic hydrogenisotope effects. The following questions may reasonably be asked. (i) What isthe nuclear configuration of the transition state, and in particular where does theproton lie? (ii) What are the forces acting in the transition state; in other words,what are its frequencies, or the shape of its energy surface? (iii) Can the answersto (i) and (ii) be predicted or interpreted by theories of interaction between bondedor non-bonded atoms? (iv) What part is played by solvation, or by change ofsolvation, in the formation of the transition state?The short life of the transition state precludes the use of the usual methods ofstructural investigation, such as spectroscopy, and the only useful experimentalevidence is that derived from the kinetic measurements themselves.A purelytheoretical approach is difficult, since it necessarily involves non-classical structureswith partly broken bonds, and for this reason an electrostatic model has often beenused.Kinetic investigations of a single process give only general information aboutthe transition state, such as its enthalpy or entropy of formation, and it is only inthe simplest systems that this information can be given a structural interpretation.A more hopeful approach is to investigate the variation of velocity with the structureof the reactants.In proton-transfer reactions this principle was first applied toacid-base catalysis, where the Bronsted relation between acid-base strength andcatalytic power, and deviations from this relation, can be given a molecular inter-pretation.1 More recently, the development of techniques for studying fast reactionshas made it possible to extend the same ideas to a wide range of acid-base reactions,and an excellent summary of this general field has been given recently by Eigen.2Isotopic substitution is a particular type of structural variation, and in proton-transfer reactions particularly valuable information may be obtained by replacingthe proton being transferred by deuterium or tritium (the primary isotope effect),1R.P. BELL 17though it may also be useful to substitute other hydrogen atoms in the reactants(secondary isotope effect) or to change the solvent from H20 to D20. These sub-stitutions will not affect the shape of the energy surface, but will modify the energylevels in a way depending upon its shape, thus providing some hope of answeringquestion (ii) above.Swain and Thornton 3 recently stated " If the transition state isotope effectof every atom were determined, valence force constants could presumably be deter-mined," and it will be shown later that this includes the negative force constantassociated with motion along the reaction co-ordinate.In practice this would demandan accurate study of deuterium and tritium isotope effects over a wide temperaturerange, but it is of interest to see what can be deduced from the more limited informationactually available.Considering a proton transfer represented by AH + B+A+ HB (charges beingomitted) the most important factor determining the primary isotope effect is thezero-point energy of the A-H bond, and the simplest expression for the isotopeeffect is k ~ / k ~ = exp (Aiso/kT), where A E ~ represents the difference in zero-pointenergies between A-H and A-D, and is derivable from spectroscopic data. Inthe common case of proton transfer from carbon, Aeo is about 1.15 kcal/mole ifonly the stretching vibration is considered, and the predicted value of k ~ / k ~ is 6.9at 25".In the transition state the stretching frequency of A-H changes into motionalong the reaction co-ordinate, A . . . H . . . B, which has no real frequencyand hence no zero-point energy. The simple picture therefore predicts that k ~ / k ~should be independent of the nature of B : moreover, since the C-H frequencyin a series of similar organic compounds is almost constant, k ~ / k ~ should vary littlein such a series.The rather fragmentary data available for proton abstraction from carbon(mostly obtained from the rates of zero-order base-catalyzed halogenation re-actions) 4 show that this simple expectation is not realized. For a given organicspecies there is definite evidence that k ~ / k ~ increases with the strength of the basewhich is abstracting the proton.This is shown particularly clearly by recent workin this laboratory 5 on the abstraction of protons or deuterons from ethyl a-methyl-acetoacetate by the anions of carboxylic acids of varying strength. The resultsare summarized in table 1, and show a smooth variation of k ~ / k ~ with the basicc -+ 4-TABLE 1 . - h T E OF IONIZATION OF ETHYL a-METHYLACETOACETATE AT 25°Ccatalyst Kl kH (M-* sec-1) kHlkDCHCl2CO 2 3 . 3 2 ~ 10-2 3 . 5 7 ~ 10-5 3 . 8 5CH2ClCO 2 1 . 3 8 ~ 10-3 3 . 0 0 ~ 10-4 5 . 1 8CH2ClCH2CO; 1 . 0 4 ~ 10-4 1 . 7 6 ~ 10-3 5.72CH3C0; 1 . 7 5 ~ 10-5 3 . 2 3 ~ 10-3 5-92Me3CCO; 9.35x 10-6 5-75 x 10-3 6.45strength of the catalyst.It is not practicable to study proton abstraction byhydroxide ion in the reaction of table 1, but we have recently found 6 k ~ / k ~ = 9.6by direct measurement of the reaction between nitromethane and hydroxide ions,which may be compared with the values 6.5, 4.3 and 3.8 previously found 7 in thezero-order bromination of nitromethane catalyzed by CH3CO 2, CH2CICO; andH20 respectively; there is again a large increase of isotope effect with increasingbase strength18 ISOTOPE EFFGCTSThere is also good evidence 4 of variation of isotope effect in a series of similarcompounds AH when 3 remains the same in the reaction AH + B+A + HB. Table 2contains results for proton transfer to water at 25", mainly from work in thislaboratory.TABLE 2.-RATES OF IONIZATION IN WATER AT 25°C(p = Bronsted exponent for base catalysis)ref.kH(sec-1) ~ H I ~ D P substance5 2.45~ 10-5 2-0 0.79CHMe(C02Me)2 5 3-53 x 10-7 2.4CHBr( CO2Et)2 5 2-15 x 10-4 2.7 0.73- CH2(C02Et)2 ;yy2,TflC"'"' 20 2 . 3 0 ~ 10-3 3.4 0.64\- (MeC0)2CHMe 8 9.7 x 10-5 3.5MeCOCH2C02Et 5 1 . 1 6 ~ 10-3 3.5 0.59MeCOCHMeC02Et 5 1 . 1 4 ~ 10-5 3.8 0.60MeN02 7 6 . 5 ~ 10-8 3.8 0.67(MeC0)zCHBr 5 3.35x 10-2 3.9 0.42MeCOCHBrC02Et 5 1 . 5 6 ~ 10-2 4.3 0.42(MeC0)2CH2 5 1 . 3 2 ~ 10-2 4.5 0.48The isotope effect varies between 2.0 and 4.5, but bears no relation to the reactivity.It does, however, correlate fairly well with the exponent of the Bronsted relation be-tween catalytic constant and base strength, probably within the experimental errorsof both quantities.Values so far obtained for the reaction between nitroparaffinsand hydroxide ions 6 show a smaller variation : MeN02, k13[ = 28 M-1 sec-1,k ~ / k ~ = 9.6 ; MeCH2N02, k~ = 5.2 M-1 sec-1, k ~ / k ~ = 9-3 ; Me2CHN02, k H =0.31 M-1 sec-1, k ~ / k ~ = 7.5. Another series of interest was investigated byStewart and Lee,g who measured k ~ / k ~ for the oxidation of five ring-substitutedphenyl-trifiuoromethyl-carbinols by CrVr in 77 % aqueous acetic acid, and founda regular increase from 7-4 to 12.9 as the reactivity of the carbinol decreased by afactor of 50. This last reaction probably involves the transfer of a hydride ion ratherthan a proton, but the principles involved should be similar.The facts described in the last two paragraphs show that isotopic substitutionmust affect not only the initial states but also the transition states of proton-transferreactions, and the obvious suggestion is that the latter involve vibrations whosefrequencies (and hence zero-point energies) are affected by the mass of hydrogen.*Such vibrations would decrease the kinetic isotope effect, and since the observedeffects are usually (though not always) smaller than those predicted in terms of theinitial state only, the problem is sometimes referred to as " small isotope effects ".I1It is, however, by no means clear what kinds of transition state vibrations are in-volved, and it is necessary to clarify this point if we are to use observed isotope effectsfor obtaining information about the transition state.If the transition state is regarded as a linear, tri-atomic species, its normal modesof vibration can be represented as follows :* Because of the short life of the transition state, doubt has sometimes been expressed 10 whetherit is legitimate to consider quantization of vibrational levels.This objection should be less validfor the zero-point energy, which is a direct consequence of the uncertainty principleR. P. BELL 19UNSYMMETRICAL STRETCHc -b cA . . . H . . . B (imaginary frequency iv3, reaction co-ordinate).“ SYMMETRICAL ” STRETCH4- -bA . . . & . . . B (v1, motion of H indeterminate).BENDINGt tIA . . . H . . . B (v2, doubly degenerate).Early explanations of small and variable isotope effects supposed that the A-Hbond was “ incompletely broken ” in the transition state, so-that some of its vibra-tional zero-point energy is retained.This is equivalent to assuming a real finite valuefor v3, and it was pointed out by Westheimer 12 that this is inconsistent with thedefinition of the transition state, which must have one normal mode correspondingto a maximum in the energy surface, and hence lacking a real vibrational frequency.Westheimer further pointed out that when the force constants for A . . . H andB . . . H are unequal, the “ symmetrical ” stretching mode v1 will in fact involvea considerable motion of the proton and can be a source of mass-dependent zero-point energy. This view has been generally adopted, and leads to the predictionthat in a series of similar reactions k ~ / k ~ will have a maximum value when the transi-tion state is symmetrical, since v 1 will then not involve any movement of the proton.This prediction is consistent with some of the experimental results mentianed above.Thus the anion of ethyl a-methylacetoacetate is certainly a stronger base than acarboxylate anion, so that in the transition states of the reactions listed in table 1the proton will be closer to the carboxylate; increase in the basic strength of thecarboxylate catalyst will therefore render the transition state more symmetricaland should increase k ~ / k ~ , as actually observed.This idea receives support fromthe correlation between k ~ / k ~ and p shown in table 2, since the latter may also bea function of the position of the proton in the transition state.34Westheimer’s treatment is certainly qualitatively sound, but closer examinationshows that the “ symmetrical ” vibrations of a three-centre system can hardlyaccount for the observed effects.If the distances A . . . H and H . . . B arerespectively r1 and r2, he writes for motion along the line of centresand the usual treatment gives a quadratic equation for v1 and v3 in terms of themasses rnl and m2 and the force constants. Westheimer now makes the simplifyingassumption k: -klk2 = 0, which corresponds to v3 = 0, and obtains a simpleexpression for v1 which shows a large isotopic dependence. For example, if kl =10 kz, ml = 12, m2 = 16, then vy/vF = 1.32, which is not much less than themaximum value of 2+. However, this result is dependent on the assumption v3 = 0 :actually we must have k: > klk2, when the equation has an imaginary root iv3, where-v$ is ameasure of the curvature of the energy barrier along the reaction co-ordinate.The assumption v3 = 0 is thus an artificial one, suggesting an activation energyof zero, and it would be more natural to suppose that the energy surface has curvaturesof similar magnitude in different directions.It is therefore of interest to calculatethe effect of varying kI2 on the magnitude and isotopic dependence of V I inWest heimer’s treatment20 ISOTOPE BFFBCTSTable 3 shows the results of calculations on the basis of eqn. (1) for variousvalues of k12, assuming as before kl = 10 kZ, ml = 12, m2 = 16.The first rowcorresponds to Westheimer's assumption k& -klk2 = 0. It is clear that an increaseof k12 produces a decrease both in v1 and in vT/vy, both of which will diminish theeffect of the " symmetrical " vibration in decreasing the isotope effect. It is im-possible to predict the value of k12 corresponding to real cases, but it would haveto be at least 2(klk2)* if the barrier curvature (measured by 23) is to be similar toTABLE 3.-bNGlTUDINAL MOTIONS OF TRANSITION STATEA = 4712~2, kl = 10 k2, mi = 12. m2 = 16k d k i kd* q - + y I k i "H 1 7 lki V l H I V 11 0.558 0 1 -321 *27 0-415 - 0.024 1.281.74 0.233 -0.136 1.143.16 0.138 - 0.952 1 a08the other curvatures of the energy surface. A similar conclusion follows from anelectrostatic treatment of the motion of a proton between two unequal negativecharges.13 If we take rn = 10 for the repulsive exponent in eqn.(10) of ref. (13),and z = 1 (corresponding to a total charge of -1 on the transition state), then inorder to make kl = 10 k2 we must put y = 0.535, which leads to k1223-5 (klkz)*.A similar result is obtained for any reasonable choice of m and z. All the abovecalculations have involved the rather extreme assumption kl = 10 k2: a smallerratio between the two force constants will diminish both the frequency and theisotopic dependence of the " symmetrical " stretching vibration. It seems doubtful,therefore, whether this vibration can account for the low values of kH/kD oftenobserved and for the variation in a series (cf.table l), at least in terms of a three-centre model.The three-centre model may, however, be a misleading one for many of theproton-transfer reactions studied in practice. This is particularly clear when thetransfer of the proton is concerted with the making or breaking of another bond inthe system, for example in the E-2 mechanism for base-catalyzed fl-eliminationreactions, for which there is good evidence.14 The transition state can be writtenasII!C . . . H . . . BIx . . . c -and it seems intuitively obvious that the degree of breaking of the C-H bond inthe transition state may vary greatly from one system to another, although thesystem as a whole must be passing through a maximum of potential energy: thisinvolves a return to the organic chemist's concept of an '' incompletely broken "C-H bond, rightly criticized by Westheimer on the basis of a three-centre model.A more precise picture is obtained by considering an idealized model in which thetransition state is linear.There will now be four normal stretching modes whichmay be represented as follows, a query denoting a small displacement which maybe in either direction. Mode (i) represents the reaction co-ordinate, in which thetransfer of the proton to the base is accompanied by a lengthening of the C-X bondand a shortening of C-C: it will correspond to a maximum in potential energR. P. BELL 21and will have no real frequency. Mode (ii) is analogous to the “ symmetrical”vibration of the three-centre model, and according to the arguments of the lastsection will have a low frequency and a small isotope effect.The new feature isrepresented by mode (iii) in which the motion of the B-H-C part of the systemresembles that in the reaction co-ordinate (i), but the motion of the C-C-X unitis no longer concerted with it, and reaction does not result: instead, the modeTABLE 4.-NORMAL MODES OF A TRANSITION STATE( i ) + t -+ t +B-H-C-C-Xf ? + c 4(ii) B- H4-C-X(iii) -+ +- ? -+ -+B- H- C- C- Xt f ? + +(iv) B-H- C- C- Xwill have a real frequency which is clearly highly dependent upon the mass of thehydrogen, even when it is symmetrically situated in the transition state. Mode(iv) will also have a real frequency, but it will be lower and may be less dependentupon the hydrogen mass.The general result of these considerations is that the concerted nature of themechanism leads to a real vibration of the transition state which is “ unsymmetrical ”as far as the system B .. . H . . . C is concerned, and therefore permits largevariations in the magnitude of the isotope effect. In fact, in the reaction C&5CH2X+C~HSCH = CH2 in presence of sodium ethoxide, the effect of replacing the CH2-groupby CD2 varies from 3.0 to 7-1 according to the nature of the group X.14 Moreover,the same arguments apply in reactions where the proton-transfer is concerned notwith the cleavage of another bond, but only with a change in its multiplicity. Thissituation is a common one in slow proton transfers; for example, in reactions ofI Ithe type B+H-C-C = O-+BH+C = G O - such as those in tables I and 2.I I I 1The transition state modes for this system will be entirely analogous to those intable 4, and the same conclusions may be drawn.Similar considerations may applyin quite simple systems : for example, the loss of a proton from RC02H or HCX3will involve changes in the C-0 distance and XCX angle respectively, and onlyone of the normal modes will lead to reaction.In practice, transition states will usually be non-linear, so that the vibrationsinvolved cannot be classified in terms of stretching and bending. However, thepart of the system represented by A . . . H . . . B is probably close to linearin many cases, and it is of interest to consider the bending vibrations of a linearthree-centre model, represented by the doubly degenerate v2.As usual, we haveno direct evidence for the magnitude of this frequency, but the bending frequencyof the ion HF; is 1225 cm-1,15 and calculation on the basis of an electrostatic model 13gives a simiiar value for a transition state.*This degenerate vibration will contribute doubly to the zero-point energy of thetransition state, and its frequency clearly depends upon the isotopic mass even ina symmetrical transition state. It therefore provides another reason for low andvariable values of k ~ / k ~ : it seems intuitively likely that the frequency of such a%-** It is sometimes stated 16 that the assumption of central forces predicts a zero frequency forthe bending vibration of linear triatomic molecules, but several authors 179 18 have shown thatthis is not the case22 ISOTOPE EFFECTSvibration will be at a minimum when the proton is symmetrically placed, and thisagrees with the results of the electrostatic treatment (eqn.(7) and (14), ref. (13)).There is, moreover, an additional reason why transverse frequencies may be quitehigh and may vary in a series of transition states. The central-force treatmentaccounts for the bending force constant in terms of non-directional forces betweenthe proton and the centres A and B, and the same picture will apply to the bifluorideion. An initial state such as -C-H will possess bending frequencies around1400 cm-1 which can be attributed to changes of hybridization with bond angle, orto interactions between non-bonded atoms.These factors will diminish as the bondis stretched, but they will still be present in the transition state, especially if it isclose to either the initial or the final state. Their result will be to increase thefrequency of the bending vibration, and also its sensitivity to the configuration ofthe transition state.So far nothing has been said about the tunnel correction for motion along thereaction co-ordinate. This is a quantum correction having exactly the same logicalstatus as the correction for zero-point energy in the transition state, both beingdirect manifestations of the uncertainty principle. In fact, it has been shown 19that under most conditions the tunnel contribution to the isotope effect can becalculated merely by inserting the imaginary frequency iv3 in the exact expressionfor a vibrational contribution (rather than the approximate expression exp (A&o/kT)).There are now several investigations of proton-transfer reactions in solution, mainlyinvolving deuterium or tritium isotope effects, which appear definitely to establishthe part played by the tunnel effect,20-24 and there is no logical justification foromitting its consideration. Its effect will always be to increase k ~ / k ~ , and treat-ment of a rather general central-force model 13 predicted that this increase wouldalways be at least as great as the decrease due to the bending vibration v2, cancella-tion being exact for a model in which the proton moves between two non-polarizablenegative charges.However, this conclusion should not be taken too seriously.The predictionmentioned in the last sentence is valid only if the bending force constant is due tocentral forces between the proton and the two basic centres, and any contributionfrom valency bending forces will not be cancelled out by the tunnel correction.Moreover, there are two reasons why the tunnel correction may be less importantthan previous calculations suggest. In the first place, the usual treatment 1 9 ~ 2 5expresses the correction solely in terms of the curvature of the energy surface alongthe reaction co-ordinate, which implies that its value is the same for symmetricaland unsymmetrical barriers. This is very nearly true when the correction is small,but it is physically obvious that a large tunnel correction will be reduced when thereis a large energy difference between the initial and final states (since no tunnellingcan take place from systems with energies lower than the higher of these states),and this is confirmed by calculations for unsymmetrical parabolic 26 and Eckhart 27barriers.* In the second place, there is no real justification for applying tunnelcorrections based on a one-dimensional barrier to systems where the correctionsare not small, and there are indications that the one-dimensional treatment mayconsiderably over-estimate the correction289 29 Nevertheless, the contribution ofthe tunnel correction to the isotope effect will certainly oppose that of the bending\/* This effect might provide an additional contribution to the variation of the isotope effect withthe configuration of the transition stateR.P. BELL 23vibrations, and is probably of the same order of magnitude : thus it may not be abad approximation to ignore both these contributions.Finally, there is one more factor which may influence the magnitude of the isotopeeffect in proton-transfer reactions. It has recently been pointed out 309 31 thatthe dielectric relaxation time for water, about 10-10 to 10-11 sec,32 is probablygreater than the time involved in a proton transfer between two properly orientatedand activated molecules, estimated at 10-12-10-13 sec. This implies that the re-orientation of water molecules cannot keep pace with the transfer of the proton,so that the free energy of the transition state will be greater than that correspondingto equilibrium orientation.Since the deuteron moves more slowly than the protonthe departure from equilibrium will be greater in proton transfer than in deuterontransfer, leading to a decrease in k ~ l k ~ . Moreover, in a series of similar reactionsthe amount of reorientation involved in passing from the reactants to the transitionstate will depend upon the distance through which the proton has to move, and thenon-equilibrium effect may therefore contribute to the variation of k ~ / k ~ in sucha series. Unfortunately, it seems impossible to make any estimate of the importanceof non-equilibrium behaviour, since this would demand a knowledge of the timeinvolved in the proton transfer and the relaxation time of water molecules in thevicinity of the reacting system.It is of interest here that many proton or deuterontransfers take place 20-40 % more slowly in D2O than in H20, even when no solventspecies are formally involved in the reaction and no isotopic change is made in thereactants. This is the case for proton or deuteron transfers from acetone to acetateions,33 from nitromethane to acetate or monochloroacetate ions,7 from methyl-acetylacetone to acetate ions,8 and from 2-carbethoxycyclopentanone to monochloro-acetate ions.20 Since the relaxation time for D20 is greater than that for HzO,32these observations might be accounted for by non-equilibrium solvation of thetransition state.There are thus many factors which may affect the magnitude of hydrogen isotopeeffects in proton-transfers, and more systematic experimental work is needed todecide their relative importance.Ultimately, however, the study of isotope effectsshould prove a most useful means of elucidating the detailed structure of transitionstates.This paper was written while the author was on Sabbatical leave at Brown Uni-versity, Providence, Rhode Island, U.S.A., and he is grateful to the Brown ChemistryDepartment for hospitality and to the National Science Foundation for a Fellowship.1 Bell, The Proton in Chemistry (Cornell University Press, New York, 1959), chap. 10.2 Eigen, Angew. Chem., 1963,75,489.3 Swain and Thornton, J. Amer. Chem. SOC., 1962, 84,817.4 ref. (l), table 24, p. 201.5 Bell and Crooks, to be published.6 Bell and Goodall, to be published.7 Reitz, 2. physik. Chem., A , 1936, 176, 363.8 Long and Watson, J. Chem. Soc., 1958, 2019.9 Stewart and Lee, Can. J. Chem., 1964,42,439.10 Kassel, J. Chem. Physics, 1935, 3, 399.11 Thornton, J. Org. Chem., 1962, 27, 1943.12 Westheimer, Chem. Rev., 1961, 61, 265 ; see also Bigeleisen, Pure Appl. Chem., 1964, 8, 217.13 Bell, Trans. Faraday SOC., 1961, 57,961.14Bunnett, Angew. Chem. (Int. Edit.), 1962, 1, 225.15 Cot6 and Thompson, Proc. Roy. SOC. A, 1951,210,206.16 Herzberg, Infa-red and Raman Spectra of Polyatomic Molecules (Van Nostrand, N.Y., 1945),p. 16124 ISOTOPE EFFECTS17 Longuet-Higgins, Phil. Mag., 1955,46,98.18 Pearson, J . Chem. Physics, 1959, 30, 1537.19 Bell, Trans. Faraday SOC., 1959, 55, 1.20 Bell, Fendley and Hulett, Proc. Roy. SOC. A , 1956, 235,453.21 Hulett, Proc. Roy. SOC. V, 1959, 251, 274.22 Caldin and Harbron, J. Chem. SOC., 1962, 3454.23 Lewis, J. Amer. Chem. SOC., 1964, 86, 2531.24 Shiner and Martin, Pure Appl. Chem., 1964,8, 371.25 Wigner, 2. physik. Chem. B, 1932, 19,203.26Bel1, Proc. Roy. SOC. A, 1935, 148, 241.27 Johnston and Heicklen, J. Physic. Chem., 1962, 66, 532.28 Johnston, Adu. Chem. Physics, 1961, 3, 131.29 Johnston and Rapp, J. Amer. Chem. SOC., 1961,83, 1.30 Kreevoy and Kretchmer, J. Amer. Chem. SOC., 1964, 86, 2435.31 Grunwald and Price, J. Amer. Chem. SOC., 1964,86,2965,2970.32 Hasted, Prog. in Dielectrics, 1961, 3, 103.33 Reitz and Kopp, 2. physik. Chem. A , 1939,184,429.34 Lefflex and Grunwald, Rates and Equilibria of Organic Reactions (Wiley, New York, 1963),p. 158

ISSN:0366-9033    

DOI:10.1039/DF9653900016

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Quantum-mechanical Tunnelling and the Dimensions of Energy-barriers in Pro ton-transfer Reactions in SolutionBY E. F. CALDIN AND (MISS) M. ISASPAWAN"Physical Chemistry Dept., The University, Leeds 2Received 15th January, 1965Quantum-mechanical tunnelling in proton-transfer reactions should lead to non-linear Arrheniusplots and to anomalous isotope effects; both these phenomena have been observed. By fittingtheoretical equations to the data, it is possible to derive values for the dimensions of the energybarriers in these reactions. The results to date are compared. The width of the energy barrierappears to depend on the atoms between which the proton is transferred, the charges on the atoms,the structures of the reacting molecules, and the solvent. An instance of a non-linear Arrheniusplot is reported, for the reaction between hydrofluoric acid and the 2,4,6-trinitrobenzyl anion inethanol over the range -90 to +25".It is well known1 that if quantum-mechanical tunnelling occurs in a proton-transfer reaction, it will lead to deviations from linearity in the Arrhenius plot atsufficiently low temperatures, and also to anomalous isotope' effects, in particularto an A-factor for H+-transfer less than that for D+-transfer.The magnitude ofthese effects will depend markedly on the dimensions of the energy-barrier, andespecially on its width. Both effects have been observed, first in Bell's laboratory 2, 3and later elsewhere. By assuming the energy-barrier to have a particular shape(a parabola is the simplest to handle), it is possible to deduce its width and height.As very little is known about the widths of energy-barriers, the results of such in-vestigations are of great interest.The largest effects observed in the work of Bell, Fendley and Hulett 2, 3 were fora proton-transfer between carbon and fluorine (CH .. . F). A reaction in whichthe rate-determining step is a proton-transfer from fluorine to carbon (C . . . HF)might be expected to have an energy-barrier of similar shape. Such a reaction isthat between hydrofluoric acid and the 2,4,6-trinitrobenzyl anion [C6H2(N02)3CH2]-,derived from 2,4,6-trinitrotoluene :The rate of this reaction was studied aver the temperature range -90 to - 50" in anearlier investigation ; 4 the Arrhenius plot proved to be linear within experimentalerror.We have now extended the temperature-range upwards to +25" by usinga stopped-flow apparatus.5 A curvature of the Arrhenius plot is now apparent,and we interpret this in terms of quantum-mechanical tunnelling. The resultscan be compared with those already obtained for the reaction of acetic acid withthe same anion.6EXPERIMENTALThe stopped-flow apparatus and its operation have already been described? The changesof optical density with time are recorded by photographing the trace produced on an oscillo-* present address : Department of Chemistry, American University of Beirut, Beirut, Lebanon.226 TUNNELLING I N PROTON-TRANSFERscope using a time-base calibrated against 50-cycle a.c. Several traces (3-6) were photo-graphed in each run; sometimes two different time-bases were used.The times for half-change varied from 0.04 to about 0.4 sec. An Ilford 623 filter was used ; this has a maxi-mum transmission at 495m,u, corresponding to the absorption maximum of the 2,4,6-trinitrobenzyl anion at about 500 mp. Temperatures were measured by platinum resistancethermometry to within fO.05 deg.The materials were purified as in previous work.4 A stock solution of T.N.T. was pre-pared fresh each day. The solvent was the same as in Caldin and Jackson's work,4 viz.,ethanol containing 0.4 % by volume of toluene (not 0.8 %, as erroneously stated in theirpaper). The ionic strength was made up to 0.002 M by addition of lithium iodide.Kinetic runs were carried out at 0, 7 and 25°C. At each temperature, a series of runsat nearly constant buffer ratio (around 16) was performed.As HF was always in large excess(not less than 30-fold), linear first-order plots were obtained.RESULTSThe results of the kinetic runs are summarized in table 1, along with the resultsat -50 to -90" from earlier work.4 The symbols used are the same as in the pre-vious work, and are as follows :-d = conc. of 2,4,6-trinitrotoluene (T.N.T.) ;b = formal conc. of ethoxide = conc. of F- in reaction mixture ;c = formal initial conc. of HF ;r = (c - b)/b = buffer ratio ;c- b = initial. conc. of HF in reaction mixture ;s" = slope of first-order plot (decadic logarithms), sec-1.TABLE 1 .--RESULTS OF KINETIC EXPERIMENTS ON HF+ C~H~(NOZ)~CH~concentrations in mole 1.-1; s" in sec-1 (decadic logs)s = slower time-base ; f = faster time-basetemp.O C0.00f0-057-00f0.0525-00i 0.05no oftrices lWd6(f) 1-544 (s) 1.545 1-544(s) 1.544(f) 1-543 1.545 1-244 (s) 1.243 (f) 1.244(f) 1-244(s) 1.245 1.244 1.244(s) 1.245 ( f ) 1.24104b5.685-6811-0016-3016-302.978.3014.7014.705-085.083.797.352.0 12-0110%0.9600.9601.8902.8302.8300.4711.4202.5602.5600.8540-8540.6251 -2503.1303.1301 O*(C - b)0-9030.9031.7802-6672.6670.4411.3372.4132-41 30.8030.8035.8711-772.932.93r15.915.916.216.416.414-916.116-416.415-815-S15.516.014.614.6s* (obs.)1.53 f0.151.44 f0-032.45 f0.063-73 f0.053.57 50.230.84 50.042-78 f0.234.96 f0.065.21 10-651.93 50.111-76 f0.144.26 f0.308-21 50.512.16 f0.072.42 A0.06s" (corr.)1.54 f0.151.54 rt0-032.42 f0.063.68 50.053.52 f0-230.99 f0-092.77 ~k0.234-90 50.075.15 f0-651-96f0.111.79 f0.144-33 f0.308.21 310.512.35 50.1 12.61 f0.11sa (Calc.)1 *491 -492.523-553.550.932-885.005.001-831.834.378.202.472.47All concentrations are in mole 1.-1, corrected for the change in volume onchanging the temperature. Where two time-bases have been used, they are desig-nated as f and s, meaning faster and slower.The values of s" derived from eachset of oscillograms have been averaged and the mean is denoted by s"(obs.); theerror indicated is the standard deviation from the mean.So that all results shallrefer to a buffer ratio of 16, small corrections based on the results of earlier worE . F . CALDIN AND M. RASPARIAN 27at different buffer ratios4 have been applied to the first-order constants when thebuffer ratio differs froni 16 ; these corrections are usually within the limits of experi-mental error (see table 1). The corrected values are denoted by s” (corr.). Thevalues of s” calculated from the “ best ” values of the rate constants (below) areshown as s” (calc.) for comparison.Rate constants for the reactions of HF and H-t with the anion C~H~(NOZ)~CH;were determined as before4 from plots of the first-order constants s” against theconcentration of HF, which is ( c - b ) .Taking account of the reactions of solvent,HF and H+, the equation for s’’ isHere kY refers to reaction with solvent ; k2 is the second-order constant for HF ;and k; is a first-order rate constant for hydrogen ions at the buffer ratio Y. Theslope of a plot of sff against (c- b) gives k2 ; the intercept (since kYl is known fromearlier work) gives ki. Straight lines were fitted to the plots by least-squares. Thebest values of k2 and kj, with the standard deviations, are shown in table 2.2.303s” = kE 1 + k2(~ - b) + kkr.temp. “C+ 25.00 f0.05 + 7.00k0.050.00 k 0.05- 49.94 f0.03- 59.92 f0.03- 69.88 f0.03- 79.86 f0.03- 89.82 f0-03TABLE 2.-RATE CONSTANTS FOR HF+ C6H2(NO2)3CHzk2 in 1. mole-1 sec-1; k j in sec-1ki k2 log k2 (obs.) log k2 (Arrh.) -10-2 (7-8 f0.07)10-2 (6.2 f0.7)10-4 (4.1 3t0.5)lO-S(3-1f1.2)lO-5(1-4f0-3)10-6 (4.8 f3-2)10-2 (3.4 50.4)10-5 (7.5 *o.s)103 (1.49 f0.04) + 3.173 f0.012102 (4.54 f0.21) + 2,657 *0.020102 (2.70f0.09) +2-431f0-0153-25 f0.18 +0.512 f0-0241.11 kO.10 +0.045&0-038- 0.445 h0.026- 1.005 f0.020- 1.565 f0.05910-1 (3.59 k0.21)10-2 (9.88 f0.44)10-2 (2.72 f0.35)+3*173 + 2.653 + 2432 + 0.448- 0.06 1- 0.61 5- 1.230- 1.910log kz (obs.)-log k2 (Arrh.)o*ooo + 0.004-0.001 + 0.064$0.106+0.170 + 0.225 + 0.345In the earlier work covering the range from -50 to -go”, the temperaturesin the reaction cell differ slightly from the measured temperature of the thermostat,the maximum difference being 0.18”.The correction has been found to dependon the liquid in the cell. Caldin and Jackson used values determined with petroleumether instead of ethanol ; we have therefore corrected their values, using the correc-tions reported by Caldin and Harbron 6 for ethanol. The uncertainty in these tem-peratures is about 0.03 deg.6REACTION WITH HYDROGEN IONSThe Arrhenius plot for the reaction with hydrogen ions (logk; against l/T),which is subject to fairly large experimental errors, shows no clear deviation fromlinearity. The best straight line through the points gives the Arrhenius parametersas E; = 9.96+0-51 kcal mole-1, log A j = 6.35+0-50. The rate constant k;, how-ever, is composite. If the rate constant for the reaction of ethoxoniurn ions EtOHzis k 3 ~ , and that for hydroxonium ions (due to traces of water in the solvent) is k 3 ~ ,then k; is given by 4wherek i = k3,Ki + ~ ~ H K H [ H ~ O ]Kg = [EtOH:][F-]/[HF], and KH = [H30’][F-]/[HF][H20].To obtain approximate values for the Arrhenius parameters for the reaction ofEtOHi, we assume, as in the earlier work,4 (a) that Kg has at all temperatures th28 TUNNELLING I N PROTON-TRANSFERsame value as at 25", which may be estimated by assuming that the pK of hydro-fluoric acid increases by 5.56 on passing from water to ethanol, like that of carboxylicacids; and (b) that the term k 3 ~ K ~ [ H 2 0 ] contributes about 10 % of kj at alltemperatures. We then obtain E3E = 9.96 kcal mole-1, loglo A3E" 15.1 (A1. mole-1 sec-I), for the reaction EtOH: + C6H2(N02)3CH;.REACTION WITH HF MOLECULESThe Arrhenius plot for the reaction with HF molecules (log k2 against l/T),shown in fig.1, is not linear over the whole temperature range now available. It isof the same general shape as that for the reaction with HOAc, but the deviationfrom linearity begins at a higher temperature. The points from 25 to 0°C lie on1 0 3 / ~FIG. 1.-Arrhenius plot for reaction of 2,4,6-trinitrobenzyl anion with HF: loglo k2 against lO3/T.a straight line within 5 1 %; the best line calculated by least squares yields theArrhenius parameters E2 = 11.1+0.1 kcal mole-1, logloA2 = 11*3+0-1 (A in1. mole-1 sec-1). From this line are calculated the values of log k2 (Arrh.) in table 2.The deviations from this line at lower temperatures are shown in fig.2. At -80and -go", the observed rate constants are about 70 :(, and 120 % faster than thecalculated rate constants. These deviations are far outside the standard deviationsin the rate constants, which are 5 % and 14 % respectively. The temperature errorrequired to account for them, if the Arrhenius plot were linear, would be more than4", which is many times the actual uncertainty of around 0.03".Bell's equations for quantum-mechanical tunnelling through a symmetricalparabolic energy barrier 7 have been fitted to the data. Only the first term in Bell'seqn. (12) was used, i.e., Q = (ncr/P)/sin (nn/P). The calculated values (kqmt) arevery sensitive to the assumed width of the barrier (2a) and depend also on its height(Eqmt).From fig. 2 it may be seen that the calculated values reproduce the observedvalues to within about 10 % over the whole range of temperature when Eqmt =11.91 kcal mole-1 and a = 0-73 A. The results when Eqmt = 12.00 kcal mole-1and a = 0-72 A also represent the observed values within about 11 o/o, the deviationsat each temperature being slightly larger than before. The best value for the half-width of the barrier is probably 0-73 A within 0.01 A. With Eqmt = 12-20 kcal mole-1and a = 0.73 A, the calculated points lie well outside the standard deviations.The values for Eqmt and a may be too high because we have assumed that thE. F. CALDIN AND M. KASPARIAN 29energy of activation is determined by the motion of the proton alone. Contribu-tions from changes in solvation, and from the changes in the bond-lengths C-CH3,G-N and N-0 due to electronic reorganization, have been ignored. Thesecontributions could be quite considerable; in that case the energy of activationassociated with the motion of the proton would be less than the overall energy ofactivation derived from the Arrhenius plot, and to fit the results we should needsmaller values of Eqmt and of the barrier-width.Some evidence that these contribu-tions are important comes from the fact that when various acids are used in place3.7 4') 4.5 4.9 5 - 3 . 5-71 @/TFIG. 2.-Deviations from the Arrhenius line : plots of [loglo k2-loglo k2 (Arrh.)] against lO3/T.Vertical lines : experimental values of k2, with standard deviations. -.-.-.- calc.curve withEqmt = 11.91 kcal mole-1, a = 0.73 8, ; - - - - talc. curve with Eqmt = 12-00 kcal mole-1,of HF the values of EA lie between 8.5 and 10 kcal/mole,4 whereas if they werewholly attributable to the motion of the proton they should, on the simplest assump-tions, cover a range of about 5 kcal/mole. This constancy suggests that the activa-tion energy may be in large part associated with motions in the substrate molecule,other than those of the proton transferred.Alternative explanations for the curvature of the Arrhenius plot 8b have beeninvestigated, with negative results, in the same way as for the reaction betweenthe trinitrobenzyl anion and acetic acid.6 (a) A change in the mechanism of thereaction a.t low temperatures is unlikely, because the absorption spectra give noindication of any new species at low temperatures,6 and the original T.N.T.isquantitatively regenerated when HF is added to a solution of the anion, as maybe shown by determining the optical density of such a solution with a Unicamspectrophotometer, adding HF and then ethoxide at low temperature, and redeter-mining the optical density. (b) An equation derived from the assumption ofvariations in AH* due to a constant AC:, such as might arise from changes ofa = 0.72 8, ; . . . . calc. curve with Eqmt = 12.20 kcal mole-1, a = 0.73 A30 TUNNELLING IN PROTON-TRANSFERsolvation on forming the transition state, was tested but could not be fitted to thedata. (c) The results cannot be attributed to changes in the structure of the solvent,since deviations occur at widely different temperatures for hydrofluoric, acetic 6and monochloroacetic acids.6z I n2RzIR2 n.. . . . . - . . . . E. F. CALDIN AND M. KASPARIAN 31OTHER REACTIONSIn table 3 are collected the results so far available on the dimensions of energy-barriers in proton-transfer reactions. They have all been obtained by applyingBell's 1959 equations for tunnelling through symmetrical parabolic barriers 1, 7 tovarious reactions which show either (a) a non-linear Arrhenius plot or (b) an isotopeeffect in which AD/& is significantly greater than unity. These two criteria aredesignated in the column headed " method " by " Arrh." and " D-H " respec-tively (or " T-D-H " in one instance where tritium exchange has also been studied).The width of the parabolic energy barrier at the base is 2a (column 6), and the heightfor the proton is E,,t. Under the heading " type " are indicated the atoms betweenwhich the proton is transferred and its initial position, e.g., " CH .. . 0 " indicatesthat the transfer is from C to 0.Reactions (l), (2) and (3) in table 3 are those studied by Bell, Fendley andHulett ; 2 the values of a were recalculated by Hulett 3a using Bell's 1959 equations 7in the first approximation, i.e., only the first term in Bell's eqn. (12) was used. Forreaction (4) we fitted the equations to Hulett's results,3a and obtained 2a = 1-42 A,the calculated and observed rate constants agreeing within 10 "/o over the wholetemperature range ; on including the second term of Bell's eqn.(12), Hulett 3b ob-tained nearly the same value, 1.40 A. Hulett 8, 3b has also found that for reaction( 5 ) the inclusion of the second term does not appreciably alter the best value of a,though it improves the fit of the equations to the experimental data.Reaction (9) has been found 12 to give a linear Arrhenius plot over the rangefrom + 19 to -78" ; taking the maximum curvature compatible with the experi-mental accuracy, we have computed that 2a must be at least 1.92A. Reaction(1 1) gives a linear Arrhenius plot from +20" to - 32" ; 13 here we find the minimumvalue of 2a comes out as 1-22 A. Reaction (12) shows clear evidence of tunnelling 11but the value of 2a has not been published.DISCUSSIONGENERAL CONSIDERATIONS.-The values collected in table 3 for the barrier-width2a are of the same order as those expected from the bond lengths and van derWaals distances, viz., 1.35 A for type CH .. . 0 or C . . . HO, and 1.34 A for typeCH . . . F or C . . . HF. No great emphasis can be placed on the absolute valueof 2a, however. On the one hand, the actual energy-barrier is probably bell-shaped,rather than parabolic, and the effective width at the base will be greater than 2a.On the other hand, the use of a two-dimensional model has been criticized, and itappears that Bell's equations may over-estimate the tunnelling correction at lowtemperatures; 14 the value of a obtained will then be larger than the true value.The relation between the value of 2a and the true barrier width is thus not knownwith precision.The values of 2a derived from experiment are themselves subject to some un-certainties.It has already been mentioned that the values will be too large ifsolvation changes, or electronic reorganization and consequent changes of con-figuration, contribute appreciably to the energy of activation. Moreover, in allthe calculations of a it has been assumed that the energy-barrier is symmetrical,so that the reaction has zero AH. If AH is not zero, the effect of tunnelling onthe rate is smaller ; 15 consequently, if the rate measurements are analyzed by meansof the equations for symmetrical barriers, the value obtained for a will be largerthan if the correct equations had been used.Unfortunately, the values of A32 TUNNELLING I N PROTON-TRANSFERare not known for most of the reactions listed in table 3. For reaction (lo), AHis + 1.4 kcal/mole ; 10 for reaction (9), it is +3.6 kcal/mole.l2 For reactions (l),(2), (3) and (4) the known pK together with estimated values of AS" give AH asaround 17 kcal/mole ; use of the correct equations would therefore give values ofa smaller than those in table 3, but no calculations have been made. For reactions(7) and (8), which go effectively to completion, and would be expected to have AS"around -20 cal deg.-1 mole-1, AH is probably more negative than -9 kcal/mole.Values of AH could be experimentally determined and the calculations extendedto take account of them.It seems worth while, however, in spite of these uncertainties, to attempt somecomparisons between reactions (1)-( 1 l), on the provisional assumption that the rela-tive values of 2a reflect the variations in barrier width, though the results may haveto be revised when values of AH and more accurate computations become avail-able.Valid comparisons might be expected particularly for closely-related pairsof reactions such as (2) and (3), which differ only in the base (types CH . . . 0 andCH . . . F); or (7) and (8), which differ only in the acid (types C . . . HO andC . . . HF) ; or (3) and (4), which differ only in the medium and method of analysis.The factors that might be expected to be important are (i) the atoms between whichthe proton is transferred ; (ii) the charges on these atoms ; (iii) the groups attachedto these atoms, i.e., the structures of the reactant molecules; and (iv) the medium.We consider these in turn.THE ATOMS BETWEEN WHICH THE PROTON IS TRANSFERRED.-We might expect,on a simple potential-energy picture of proton-transfer, that these atoms wouldbe the most important influence on the barrier-width, since their repulsions willdetermine the minimum van der Waals distance of approach in a non-reactivecollision (AH.. . B) and the bond lengths A-H and H-B+ will then fix thethe distance that the proton has to travel in the reaction. Examination of table 3suggests, however, that other factors are also involved. The following comparisonsare relevant.(i) The barrier width 2a varies considerably for a given type ofreaction. Thus, for reactions of type CH . . . 0 and C . . . HO (reactions (l), (2),(5), (6), (7), (lo), (9) and (11)) the values are 1.26, 1.17, 1-13, 1.59, 1.66, 1.64, <1.92,and +1.22k For reactions of type CH . . . F and C . . . HF (reactions (3), (4)and (8)), they are 1-17, 1-40 and 1.46A. (ii) The barrier width is not the same forreaction (7) as for reaction (9) which is closely related to the reverse of reaction (7).The difference is at least 0.26A. This would probably not be reduced by using theequations for an unsymmetrical barrier, since AH is numerically larger for reaction(7) (cf. above), so that on revision the value of 2a would probably be reduced morethan for reaction (9).(iii) The difference between barrier widths for reactions oftypes CH . . . 0 and CH . . . F varies considerably, even when the groups attachedto the carbon atom are the same and the conditions similar; thus for reactions(2) and (3) the two values of 2a are nearly the same, but for reactions (7) and (8)they differ by 0.2A. The values calculated from bond lengths and van der Waalsdistances agree closely (1 -35 and 1.34 A) ; apparently some additional factor isinvolved.THE CHARGES ON THE ATOMS.-A comparison of reactions (1) and (2) may indicatethe effect of the charge on the oxygen atom when the other conditions are as similaras possible. The value of 2a is smaller by 0.09 A for reaction (2) than for reaction(l), possibly because of the extra attraction of the negatively-charged oxygen of theanion for the protonE .F . CALDIN AND M. KASPARIAN 33THE STRUCTURES OF THE REACTANT MomcuLw-The groups attached to theatoms between which the proton is transferred might affect the barrier width in severalways. (i) As donors or acceptors of electrons, they will decrease or increase thestrength of any hydrogen bond that may be formed before the proton-transfer occurs,and will also affect the repulsion between the two reactants. (ii) If conjugationcan occur, it will alter the distribution of charge and the effective charge on thecarbon atom, and so affect the repulsive forces. However, if we compare the pairof reactions (2) and (7) (type CH . . . 0) with the relatedpair(3) and(8) (typeCH .. . F),the predicted effects are the reverse of those observed; in reactions (7) and (8)the carbon atom is conjugated with C ~ H ~ C H ~ ( N O ~ Z , whereas in reactions (2)and (3) it is conjugated with C 4 whose effect should be weaker, yet the values of2a are markedly shorter (by 0.5 and 0.3 A) for reactions (2) and (3).THE SIZES OF THE SUBSTITUENT GROUPS may be important; they may lead todifferent degrees of steric hindrance and thus affect the barrier width. Examinationof models does not, however, suggest that steric hindrance will be much greaterfor reactions (7) and (8) than for (2) and (3). Lewis has recently suggested 11 thatsteric hindrance is a major factor for reaction (1) ; the values of k ~ / k ~ increasemarkedly when B is changed from pyridine (9.84) through 2-picoline and 2,6-lutidine to 2,4,6-collidine (24.2).Lewis points out that in a sterically-hinderedtransition state much of the energy results not from the stretching or bending ofbonds, but from compression (repulsion) and therefore depends on a high powerof the distance ; consequently the barrier is high and thin, and tunnelling is favoured.Lewis' suggestion is clearly an important one, and should be followed up.THE mDIuM.-Table 3 shows that there is a group of reactions ((l), (2), (3), (5))with values of 2a around 1.1-1.2A; another group ((6), (7), (10)) with markedlyhigher values around 1.6 A ; one reaction (9) with the exceptionally high value1.92 A ; and two reactions ((4), (8)) with intermediate values.The high and lowvalues do not correlate with the type of reaction (CH . . . 0 or CH . . . F), with themethod of analysis (isotope effect or Arrhenius plot), or with temperature of observation(for example, reaction (6) was studied at 25-65", reactions (7) and (10) down tobelow -looo). They are correlated, however, with a difference of solvent; thelower values refer to reactions in water and the higher to reactions in ethanolicsolutions. It is also clear from a comparison of reactions (3) and (4) that themedium can be of considerable importance ; the value of 2a is higher by 0.23 Afor reaction (4), carried out in 5.2 M aqueous sodium bromide solution, than forreaction (3), which was identical except that the solvent was D20 containing 0.2 MKBr.It is therefore important to consider how the solvent might influence thebarrier width (apart from any consequences due, as mentioned earlier, to the effectof a change of solvation on the height of the energy barrier).The anions concerned in these reactions are no doubt solvated by one or moresolvent molecules, with the hydroxylic H (or D) atoms of the solvent adjacent tothe anion. The initial state in a reaction of the general type CH . . . X- in a solventROH or ROD (R = H or Et) may thus be represented asR/0IH\ I-C-H . . . X-34 TUNNELLING I N PROTON-TRANSFERA solvating ethanol molecule would be expected to exert a stronger repulsion onthe carbon atom, and so give rise to a wider energy barrier, than a solvating HzOor D20 molecule, both because it is more bulky and because the oxygen atom inethanol is more negative than the oxygen in water.This effect may explain thehigher values for reactions (6), (7), (8) and (10) compared with those for reactionsThe difference between the barrier widths for C . . . HO and C . . . HFreactions in ethanol ((7) and (8)) may be due to the hydrogen-bonding propertiesof fluorine. There is no direct evidence for hydrogen bonds involving carbon andfluorine, but HF is known to form strong hydrogen bonds, and CH-0 and CH-Nhydrogen bonds are well established, in chloroform solutions and HCN respectively.16Hydrogen bonding of -CH to the anions may thus occur in ethanolic soluticn,and more strongly with fluorine than with oxygen, with a consequent shorteningof the CH-F distance relative to the CH-0 distance. In water, with its higherdielectric constant, this hydrogen bonding would be expected to be weaker and itseffect on the barrier width less marked.17).It is larger by 0 - 2 3 A than for reaction (3), from which reaction (4) differs only inthe high concentration of sodium bromide in the medium; the ratio [Na+]/[H20]is about 0.1.This implies that a F- ion will generally be near a Na+ ion, which willreduce the attraction of F- for the partial charge on the proton in -CH(6+).This will lead to a greater minimum distance of non-reactive collision, and so toa wider energy-barrier. (On this interpretation the approximate agreement betweenthe values of 2a for reactions (4) and (8) is largely fortuitous.)(21, (3) and (5).\/The barrier-width for reaction (4) finally requires comment (cf.Hulett\/SUMMARY OF THE PRESENT POSITIONQuantum-mechanical tunnelling in proton-transfer reactions is now past thestage of discovery and entering that of systematic study. If we take the values of2a in table 3 as representing the barrier widths, we can see clear instances wherethis width is affected, as we should expect, by the atoms between which the protonis transferred (compare reaction (7) with (8)), by the charges on the atoms (compare(1) with (2), by the groups attached to them (compare (7) and (8) with (2) and (3)),and-to a surprising extent-by the medium (compare (3) with (4)). There are,however, some considerable difficulties in determining barrier widths, and in theinterpretation of the values; the differences between the highest and lowest valuesof 2a do not seem to be satisfactorily explained. Systematic experimental workis needed, especially on closely-related series of reactions ; and the computationsshould be extended to take into account the effect of non-zero heats of reaction.We are grateful to Mr. R. P. Bell and Dr. J. Hulett for helpful discussions.One of us (M. K.) acknowledges a British Council maintenance grant.1 Bell, The Proton in Chemistry (Methuen, London, 1959), chap. 11.2 Bell, Fendley and Hulett, Proc. Roy. Soc. A, 1956,235,453.3 Hulett, (a) Proc. Roy. Soc. A , 1959, 251, 274 ; (b) personal communication.4 Caldin and Jackson, J. Chem. Soc., 1960, 2413E. F. CALDIN AND M. KASPARIAN 355 Allen, Brook and Caldin, Trans. Faraaky SOC., 1960,56, 789.6 Caldin and Harbron, J. Chem. SOC., 1962,3454.7 Bell, Trans. Faraday SOC., 1959, 55, 1.8 Hulett, (a) J. Chem. SOC., 1965,430 ; (b) Quart. Reo., 1964, 18,227.9 Shiner and Martin, Pure Appl. Chem., 1964, 8, 371.10 Caldin and Kasparian, in preparation.11 Funderburk and Lewis, J. Amer. Chem. SOC., 1964,86, 2531.12 Caldin and Long, Proc. Roy. SOC. A, 1955, 228, 263.13 Bell and Norris, J. Chem. SOC., 1941, 854.14 Johnston and Rapp, J. Amer. Chem. SOC., 1961, 83,l. Sharp and Johnston, J. Chem. Physics,15 Bell, Proc. Roy. SOC. A, 1936, 154,423.16 Pimentel and McClellan, The Hydrogen Bond (Reinhold, New York, 1960), chap. 6.17 Hulett, Trans. Faraday SOC., 1963, 59, 1815.1962,37,1541

ISSN:0366-9033    

DOI:10.1039/DF9653900025

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Proton Transfer Reactions Occurring in Gas-PhaseRadioly sisBY P. AUSLOOS AND (MRs.) S . G. LIASNational Bureau of Standards, Washington, D.C. 20234Received 4th January, 1965From the products formed in the gas-phase radiolysis of various reaction mixtures containingdeuterium labelled compounds, it is inferred that: (a) H:, ArH+, KrH+, and probably XeH+transfer a proton to n-pentane. In all cases, the protonated pentane ion decomposes to methane,ethane, propane, and the corresponding butyl, propyl, and ethyl ions. The relative probabilityof these three fragmentation processes does not vary with the nature of the proton donor, withinexperimental error. (b) H i , ArHf, KrH+ and CHS effectively transfer a proton to cyclopropane.On the basis of isotopic labelling experiments, it is deduced that the protonated cyclopropane re-arranges to the sec-propyl ion structure prior to or during reaction.Similarly, the protonatedcyclobutane rearranges to the sec-butyl ion structure. (c) H:, ArH+ and KrH+ transfer their protonto ethylene, propylene and butene, to form mainly ethyl, sec-propyl, sec-butyl ions, respectively.Carbonium ions such as C2HS can also transfer a proton to olefins. In the latter case, however,alternative modes of reaction such as addition and hydride transfer reactions occur as well.Gas-phase proton transfer reactions have been reported to occur in the massspectrometer, but they have seldom been observed at atmospheric pressures underthe action of ionizing radiation. The major reason for this has been that thepresence of ionic reactants or intermediates in such systems can only be inferredfrom a product analysis after irradiation.However, in recent radiolytic studieswhich were carried out in the presence and in the absence of an electrical field,and in which extensive use was made of deuterium-labelled compounds, a distinctioncould be made between products of ion-molecule reactions and those of neutralexcited molecule decompositions.1 It has also been demonstrated that it is feasibleto trace the course of ionic reactions in a radiolytic system in much the same wayas free radical reactions are studied in photolytic or pyrolytic systems. Becauseone sees only the neutral product of the ion-molecule reaction in radiolysis, theinformation obtained from the radiolytic system is complementary to that gainedfrom mass spectrometry where only the charged species are observed.Conventionalanalysis of the neutral products formed in the radiolysis of suitable deuterium-labelled compounds has the advantage that it is possible to establish, from thepositioning of the deuterium atoms in any particular product, the structure of thereacting ion or of the intermediate complex. These techniques were applied instudies 29 3 on proton-transfer reactions occurring in the gas-phase radiolysis ofhydrocarbons. The results presented here are an extension of that work.EXPERIMENTALMATERIALSThe deuterated compounds were obtained from Merck, Sharp and Dohme of Canada,Limited. All compounds were purified by means of gas chromatography.Mass-spectro-metric analysis indicated that n-pentane-dl;? contained 10 % CSDllH ; CyClOprOpa.ne-d63P . AUSLOOS AND S . G . LIAS 37contained 7-4 % C~DSH; propylene-d6 contained 4.3 % C3DsH; ethylene44 contained3.4 % C2D3H; propane-d8 contained 3 % C3D7H; and isobutaned10 contained 8 %C4DgH. Phillips reagent-grade CH4 was purified by repeated slow distillation from - 195to -220°C. Assayed reagent-grade hydrogen, xenon, krypton, and argon were obtainedfrom the Air Reduction Company. The deuterium gas, obtained from General DynamicsCorporation, contained 0.5 % HD.IRRADIATION AND ANALYSISPyrex reaction vessels of 500 ml provided with breakseals were used in all experiments.Prior to being filled, the cells were heated under vacuum close to the m.p.of Pyrex. Thecells were irradiated at 40f5"C in the National Bureau of Standards 50,000 Curie 6OCosource. Dosimetry was based on the measurement of the saturation current in a speciallyconstructed reaction vessel.1 Taking a W value for H2 of 36.3 eV, the energy absorbed byhydrogen is 1 . 3 5 ~ 1018 eV/mole sec.After irradiation, quantitative analysis of the products was carried out by expandingan aliquot of the irradiated material into an F and M fractometer provided with a silicagel or an alumina column, a flame ionization detector, and temperature programming.Subsequently, hydrogen was removed through a spiral trap immersed in liquid hydrogenas a refrigerant. Methane was then distilled and the remainder of the sample was intro-duced on to a Perkin-Elmer vapour fractometer (silica gel column) from which the productcompounds were collected separately from the helium stream at the exit of the instrument.All fractions were introduced into a Consolidated 21-101 mass spectrometer in order todetermine their isotopic compositions.RESULTS AND DISCUSSIONOnly a summary of the experimental observations is presented in this paper.Many of the results quoted are derived from detailed product analyses not givenin their entirety here.Tables containing the isotopic compositions of the differentproducts and the derivations of the ion pair yields and relative reaction rates dis-cussed can be obtained from the authors upon request.Nitric oxide was added as a free radical scavenger in most radiolysis experiments.Although removal of free radicals by NO usually simplifies the interpretation ofthe radiolysis mechanism, NO may also, to some extent, interact with some of theions present in the system.In several cases, therefore, the concentration of NOwas varied over a certain range in order to assess the effect of NO on the ionicreaction mechanism.PROTON TRANSFER TO n-PENTANEWhen hydrogen is irradiated in the presence of a small amount of alkane, RH,the following reaction mechanism will occur2 :Hz+H2+Hz+HH3f +RH-+H,+RHzThe protonated hydrocarbon thus formed will eliminate H2 or a smaller alkane,and the corresponding alkyl ion. For example, it was inferred that protonatedn-pentane-dlz decomposes as follows ;C5D12H'-+CD,H+sec-C4D: (3)C5D, 2H+-+C2D5H + sec-C,D; (4)C5Dl,H+-+C3D7H + C,Df ( 5 38 GAS-PHASE RADIOLYSISThe fragment carbonium ions (CnDin+1) formed in reactions (3), (4) and (5) reactwith n-pentane in a hydride transfer reaction :forming the corresponding fully-deuterated hydrocarbons as products (table 1).TABLE 1.-PROTON TRANSFER TO n-PENTANE-& ; MOLECULES FORMED PER ION PAIRsystemn-C~D12+ NO(1 : 0.05)H2+n-CsD12+NO(300 : 1 : 0.5)Ar+ H2+ n-C5D12+ NO(1100:400: 1:0.6)Kr+H2+n-C5D12+NO(500 : 300: 1 : 0.5)Xe+ H2 + n-C5D12 + NO(300 : 300 : 1 : 0.5)H2+ Xe+ n-C5D12+ NO(300 : 1 : 1 : 0.5) *H2+Xe+n-C5D12+NO(300 : 10 : 1 : 0.5) *0.56 0.440.34 0-29n.d.0.29n.d. 0.0990.60 0.480.77 0.67(0.63)0.14 0.140.050 0.0580.053 0.058- 0.390.1 1 0.130.12 0.40(0.13)0-069 0.0650.029 0.0240.029 0.026- 0.120.054 0.0650.065 0.15(0.062)* Values given are molecules formed per H i initially produced.Values in parentheses are valueswhich may be attributed to proton tansfer reactions after correction for contributions from charge-transfer radiolysis. n.d., not determined; -, not formed.No protonated alkanes except CH3 and C2Hj have beed observed in the massspectrometer, indicating that the higher homologues have a relatively short lifetime.This is corroborated with the protonated pentane by the fact that, even at a pressureof 25 cm, the ion pair yields ascribed to processes (3), (4) and (5) (0.56, 0.14, and0.07, respectively) account for 0.77 of the HZ ions formed.This is reasonably closeto unity if one considers that other modes of decomposition may still be unaccountedfor. For example, the decomposition of CsD12H+ to give hydrogen and a pentylion would be difficult to detect by the analytical methods used in this study.It is of interest to determine whether protonated pentane could be producedusing other proton donors such as ArH+, KrH+ or XeH+, and whether the relativeprobabilities of the modes of decomposition would vary with the AH of the reaction.Therefore, several H2 +inert gas + C5D12 + NO mixtures were irradiated, in whichthe ratio inert gaslhydrogen was kept sufficiently large so that more than 95 %of the energy was absorbed by the inert gas. With argon and krypton,Kr++Hz+KrH++H (7)Ar++Hz-+ArH++H (8)are the major modes of reaction of Ar+ and Kr+ in such a system.4 Charge transferfrom krypton to hydrogen cannot occur because the krypton ionization potentialis lower than that of hydrogen, while mass spectrometric studies5 have shown thatcharge transfer from argon to hydrogen should be of minor importance comparedto reaction (8).If KrH+ or ArH+ formed in these mixtures does transfer its proton to n-CsD12,the major lower hydrocarbon products which should be observed would be theproducts resulting from decomposition processes (3), (4) and (3, followed by thP.AUSLOOS AND S . G . LIAS 39hydride transfer reaction (6). The fact (table 1) that yields of CD3H, C2DsH andC3DjH are comparable to those of C4Dl0, C3Ds and CzD6, respectively, showsthat processes (3), (4) and (5) occur in these systems to the exclusion of chargetransfer from the inert gas to pentane, which would lead to the formation of a largeyield of C3Ds and smaller amounts of C2D6 and C4D10, all in the absence of partiallydeuterated products.6 If the product yields are calculated relative to the energyabsorbed by the inert gas (as in table l), the sum of processes (3), (4) and (5) accountfor about 50 % of the Ar+ and Kr+ ions.On the other hand, the product yieldsaccount for more than 10 times the maximum number of H i ions which could beformed in the system by reaction (l), thus demonstrating conclusively that the protondonors were ArH+ and KrH+.In contrast to the results obtained on the krypton and argon mixtures when aXe+H2+CsD12+NO mixture is irradiated with more than 95 % of the energyabsorbed by Xe, the observed product distribution is closely similar to that ob-served in the xenon-sensitized radiolysis of pentanes in the absence of hydrogen,clearly indicating that charge transfer from xenon to pentane is the major processtaking place.On the basis of the product yields, at least 80 % of the Xe+ ions canthus be accounted for.These observations are not unexpected since there is no definite proof in theliterature for the formation of XeH by reaction (9)Xe++H2-+XeH++H. (9)Stevenson and Schissler 7 point out that this reaction has not been observed in themass spectrum of xenon + hydrogen mixtures and hence the cross-section of thisreaction must be less than 50 times that for the analogous formation of ArHf.In an attempt to form XeH+, a mixture of hydrogen and xenon was irradiatedin which the relative concentrations were adjusted so that at least 86 % of theenergy was absorbed by hydrogen.When pentane was added to such a mixturein concentration equal to that of xenon (table 1, 6th experiment), the product dis-tribution and yields were nearly identical to those observed in the radiolysis ofH2 + C5D12 +NO mixtures in the absence of xenon. Because, according to Thompsonand Schaeffer,ss 9 reaction (1 0),H l +Xe+XeH++H, (10)occurs with a rate comparable to the collision frequency, one would expect that aproton is transferred to xenon from H'j at least as rapidly as to C5D12, thus, leadingto diminution in product yields if XeH+ does not, in turn, transfer its proton topentane, but rather undergoes neutralization.10 Such a drop in yield is not ob-served even when the concentration of xenon is increased to 10 times that of pentane,thus indicating that proton transfer from XeH+ to pentane may be efficient.The relative probabilities of the different modes of decomposition of the pro-tonated pentane ion formed by proton transfer from H;, ArH+, KrH+, XeH+ and,as recently reported,ll CHS, do not change although the AH of the reaction differswith the different proton donors. Although a variation in the ion distribution withthe AH of the proton transfer process has been observed in the mass spectrometer,l*these changes are apparently not due to a change in the relative probability of theinitial fragmentations, but rather to a lower or higher degree of decomposition ofthe fragment carbonium ion.No evidence for such fragmentation is seen at themuch higher pressures at which the radiolytic experiments were carried out,indicating that these ions are apparently stabilized because of the higher collisionfrequency40 GAS-PHASE RADIOLYSISPROTON TRANSFER TO CYCLOALKANESWhen a proton is transferred to a cycloalkane, the resulting protonated entity,C,H;,+ 1, is isomeric to a carbonium ion and, therefore, in contrast to the protonatedalkanes such as C5D12H+ discussed above, can be expected to have some stability.That is, the protonated cycloalkane may subsist for a sufficiently long time causingit to react as a carbonium ion instead of decomposing, in which case the structureof the ion can be determined.Proton transfer to cyclopropane was studied indetail. However, exploratory experiments indicate that analogous processes occurwith other cycloalkanes such as methylcyclopropane and cyclobutane.That fragmentation of the protonated cyclopropane formed in reaction (1 1)is not an important process is shown in the radiolysis of H ~ + c - C ~ H ~ mixtureswhere the ion pair yields of the products up to C3 account for only about 12 %of the Hf in the system. The major products up to C3 are methane and ethylenewhich are formed in equal amounts with ion pair yields of about 0.12 at an H2/C-CfH6ratio of 630.These yields vary little with the amount of NO added to the reactionmixture or a change of the ratio H~/c-C~H~.A plausible mechanism for the formation of methane would be reaction (11)followed byThe vinyl ion, which would be formed in this decomposition, may react with cyclo-C3H; -+CH,+ C2H:. (12)propane,C2Hz + c-C~H~+C~H: + CzH4,to form ethylene. This would account for the fact that, in the scavenged experi-ments, the ethylene yield is approximately equal to that of methane. The factthat the ethylene fraction observed in the radiolysis of H2 + c - C ~ H ~ + c-CJD~ mixturesconsists entirely of C2H4 and C2D4 is consistent with the suggestion that theethylene is formed as a result of a reaction such as (13).Because only a fraction of the C3H; ions formed in reaction (11) apparentlydecompose, it is of interest to determine the modes of reaction, as well as the structureof the stabilized entity, C3H3.A derivation of this structure is of particular interestin view of the suggestion of Meyerson et aZ.139 14 that C3H3 ions, originating fromthe fragmentation of alkane ions in the mass spectrometer, have a protonated cyclo-propane ring structure. On the other hand, Stevenson 15 has postulated that theseions acquire the sec-propyl ion structure.Since propane is a minor product (ion pair yield <0.007) in the scavenged H2-cyclopropane radiolysis, it is evident that the protonated cyclopropane producedin reaction (1 1) does not undergo a hydride transfer reaction,C3H; +C-C3H6-*C3Hg+C3H: (14)with cyclopropane.If the protonated cyclopropane does acquire the sec-propylion structure, this observation would be expected since it was demonstrated in oneexperiment that sec-propyl ions, produced in the radiolysis of isobutane-dlo,l6 didnot react with additive cyclopropane to produce propane-&,but reacted entirely with the isobutane to form propane-dg,sec-C3DT +iso-C4D10+C3D8 + C4D,f. (16P. AUSLOOS AND S. G. LIAS 41However, from these results, it could be inferred that sec-propyl ions react withc-C& at a rate which is 0.64 that of reaction (16), to form a more complex ion whichis probably removed from the system by NO.Thus sec-propyl ions, if produced, cannot be detected in the hydrogen-cyclo-propane system unless some compound is added to the system with which thepropyl ion undergoes a hydride-transfer reaction.Cyclohexane was chosen forthis purpose, not only because sec-propyl ions react readily with cyclohexane 17but also because an experiment in which cyclohexane was added to hydrogen in theabsence of any other compound showed that proton transfer to cyclohexane producesessentially no product compounds below c.6.It is seen (table 2) that, when c-C6D12 is added in various concentrations toH2 + C-C3H6 + NO mixtures, propane becomes the major product. More than90 % of the propanes formed consist of CH3CHDCH3, thus demonstrating thatreaction (1 7) occurs and that protonated cyclopropane rearranges to the sec-propylion structure prior to or during reaction.If C3H3, formed in a process such as(ll), retained its ring structure, then, on statistical grounds, it would be expectedthat CH2DCH2CH3 should be a major product. The above observations are alsosupported by the fact that, in the radiolysis of D2 + ~C3D6-t- C-C~H~Z. mixtures,propane consists entirely of CD3CDHCD3.From the observation that the propane-& formed in the radiolysis of H2+C-C3D6 + C-CgD12 mixtures contained approximately 80 % CH2HCD2CD3, it may beconcluded that the protonated c - C ~ D ~ isomerizes mainly to the structure CD2HCDC3which is statistically favoured over the alternate configuration CD3CHCD;.From earlier results 17 and those cited in this discussion, it is calculated thatthe rate of propyl ion addition to cyclopropane is 0.42 times as fast as reaction(17).Knowing this, it is possible to calculate, from the ion pair yields ofsec-C,H; +c-C,D,,+CH,CHDCH, + C,DTl, (17)TABLE 2.-PROTON TRANSFER TO CYCLOPROPANE; MOLECULES FORMED PER ION PAIRCH4 GH4 CHSCHDCH, CH2DCHzCHzD calc. C3H:H2f C-C3H6+ C-CgD12f NO 0.084 0.076 0.22 <0.002 0.65H2+ C-C3H6+ C-CgD12+ NO n.d. 0.047 0.26 <0*002 0.36Ar+ H ~ + c - C ~ H ~ + c-CgD12+ NO 0-032 0.04 0-25 <0*002 0-53KrS H2-t c-C&+ c-C&2+ NO n.d. 0.07 0.17 0.005 0.37Xe+ H2+ c-C3H6+ c-C&2+ NO 0-021 0.17 0.040 0.086 0.086(176 : 1 : 0.2 : 0.1)(195 : 1 : 1 : 0.5)(169 : 69 : 1 : 0.37 : 0.15)(99 : 67 : 1 : 0-37 : 0.15)(65 : 72 : 1 : 0-36 : 0.15)CH3CHDCH3 given in table 2, the total ion pair yield which can be ascribed to propylions.In addition, if one assumes that a proton is transferred with the sameprobability to C-C3H6 as to c-CgD12, noting that process (12) occurs in this systemwith an ion pair yield of 0-08, an ion pair yield of 0.91 is obtained for H;. Thus,it may be concluded that the proton-transfer reaction is highly efficient and that themajority of the protonated cyclopropane ions arrange to the stable sec-propylion structure and react further with either cyclopropane or the additive cyclohexane.A few H2 f C-C3H6 4- C-CsD12 + inert gas mixtures were irradiated for which theratio inert gas/hydrogen was kept sufficiently large so that more than 95 % of thetotal dose was absorbed by the inert gas. Processes (7) and (8) should again b42 GAS-PHASE RADIOLYSISthe major modes of reaction of the Krf or Art ions.It can be seen (table 2) thatCH3CHDCH3 is the major product formed in these systems, indicating that processes(7) and (8) are followed by the proton-transfer reactions :ArH+ 4- C-C,H,+Ar+ C,HT (18)KrH' +c-C3H6+Kr + C3HT (19)C3H6D2 is not a product, clearly showing that charge transfer does not occur inthese systems and direct radiolysis is unimportant, as it has been demonstrated 18that the cyclopropane parent ion, if formed, would undergo the D ;-transfer reactionThe yields of the observed products in these experiments calculated relative to theenergy absorbed by the inert gases account for about 80 % of the Ar+ and 50 %of the Kr+ ions. The proton-transfer reaction is, therefore, a major process under-gone by ArH+ and KrH+ in these mixtures.This finding, as well as the resultsof the pentane experiments discussed above, contradicts the assumption 1 9 s 20 thatthe only fate of ArH+ in the presence of hydrocarbons is the neutralization processArH++e-+Ar+H. (21)Since processes other than (21) occur, the rate constants derived in those studieswill be in error.When xenon is added to the H2 + c - C ~ H ~ + c-C6D12 +NO system, CH3CHDCH3is only a minor product, and CH2DCHDCH3 is formed in appreciable yield,demonstrating that parent cyclopropane ions are formed by the charge-transfermechanism :followed by reaction of C3H: with c-CgD12 (reaction (20)). Thus, again the H+-transfer reaction (9) occurs at a much lower rate than the analogous processes (7)and (8).It may be expected that CHS, which is formed in the radiolysis of methane 11by the processwould effectively transfer a proton to c - C ~ H ~ ,Xe++C-C3H6+Xe+C3Hl (22)CH: + CH4+CHS + CH3 (23)CHf $C-C,H6+CH,+C3HT.(24)This process is exothermic by about 65 kcal if one accepts 21 a value of 234 kcalof AHf(CHf), and a sec-propyl ion structure for C3H3. The formation ofCH3CHDCH3 in the radiolysis of a CH4 + C-C& + c-CgD12 mixture (table 2)confirms the occurrence of process (24). Also, the fact that the value of 0.69, whichcan be calculated for the ion pair yield of sec-C3Hi, agrees closely with the estimatedvalue of 0.68 for the ion pair yield of CHj,ll indicating that reaction (24) occurswith a high efficiency.In the radiolysis of H2+C4H8 (600: 1) mixtures, fragmentation of the C4Hi,which may be expected to be formed by the proton-transfer reactionH: + c - C ~ H ~ + C ~ H ~ +Hz (25)is of minor importance.Ethane and ethylene are the major lower hydrocarbonproducts, but they account for only 10 % of the H3 ions. On the other hand,CH3CH2CHDCH3 is a product in the radiolysis of a H2 + C-C4Hg + c-CgD12 + NP. AUSLOOS AND S. G . LIAS 43(400 : 1 : 0.4 : 0.2) mixture, indicating that a fraction of the C4Hi ions formed inprocess (25) are stabilized and rearrange to the sec-butyl ion structure prior to,or during reaction.PROTON TRANSFER TO OLEFINSProton transfer to an olefin gives a protonated entity which, as with the proton-ated cycloalkanes, is isomeric to a carbonium ion and, in view of the results presentedabove, may be expected to react as such.Unlike the protonated cycloalkane,however, the protonated olefin can assume a carbonium ion structure withoutring opening or rearrangement.Hydrogen or deuterium was irradiated in the presence of small amounts ofethylene, propylene, 1 -butene, cis-2-butene, and trans-2- butene. Only thoseexperiments with propylene additives will be discussed in detail.When hydrogen is irradiated in the presence of a small amount of added propylene,the formation of H3 by reaction (1) shouId be foIlowed byThe only products up to C4 which are observed in such an experiment are methaneand ethylene (ion pair yield = 0.023), thus demonstrating that fragmentation ofthe protonated propylene occurs only to a minor extent.Because propane is notobserved, the C3H'; produced in reaction (25) must undergo a hydride-ion transferreaction with propylene to a limited extent, if at all. On the other hand, in anindependent experiment, propyl ions formed in the radiolysis of isobutane wereobserved to add to propylene, probably to form a heavier ion which can be removedfrom the system by NO. Therefore, in order to measure the yield of the propylions, varying amounts of cyclohexane were added as interceptor to hydrogen+propylene mixtures in a series of experiments. As with the analogous cyclopropaneexperiments discussed above, CH3CHDCH3 is the major product produced in theirradiation of a H2 + CH3CHCH2 + C-CgH12 +NO mixture, thus showing thatreaction (26) is followed by reaction (17) in this system and, therefore, that the pro-tonated propylene, C3H';, has the sec-propyl ion structure prior to, or during re-action.That some rearrangement occurs in the protonated intermediate is shownby the irradiation of H2 + CD3CDCD2 + C - C ~ D I ~ mixtures, where the productpropanes consist of CD~CDZCD~H and CD3CDHCD3 in a ratio of about 3.5 to 1.Knowing that the rate of propyl ion addition to propylene is 1.53 times fasterthan hydride transfer with cyclohexane 17 (reaction (17)), we may deduce from theyields of propane attributed to reaction (17) a total propyl ion yield of about 0.44.Further, assuming that a proton is transferred from H i to cyclohexane at the samerate as to propylene, it is calculated that the average yield of HS accounted for bythe observed products is 0.80, indicating that the proton-transfer reaction to propyleneoccurs with a high efficiency.In similar experiments with ethylene as the added olefin, CH3CH2D appearedas the major product with an ion pair yield of 0.19 when the C ~ H ~ / C - C ~ D I ~ ratiowas 2-0.From relative reaction rates determined in separate experiments (ethylion addition to ethylene occurs with a rate which is 1.2 times as fast as hydridetransfer with cyclohexane), it is calculated that about 98 % of the H+ ions have beenaccounted for, thus demonstrating that proton transfer from HS to ethylene ishighly efficient.Again, in similar experiments with 1 -butene, cis-2-butene, and trans-2-buteneas the olefins added to H2 with c-C6D12 as the butyl ion interceptor, n-butane-&having the structure CH3CHDCH2CH3 appeared as a product in all three cases,CH,CHCH2 + H l -+C,H; + H2 (2644 GAS-PHASE RADIOLYSISthus demonstrating that the protonated butene reacts as a sec-butyl ion.Isobutanewas not a product, indicating that C4HZ does not always rearrange to the thermo-dynamically more stable t-butyl ion.Proton transfer reactions from carbonium ions to larger olefinsCmH,+,+ 1 + CnH2n+CmH2m + CnH,+,+ 1 (27)are usually exothermic, but may occur with low probability because addition tothe double bond, as well as hydride-transfer reactions, may constitute alternativemodes of reaction of the C,Hl,+l ion.From the radiolysis of C3D~fC3H6-NOmixtures, it could be deduced that all three reactions occur. In these mixtures,C3Dg is the source of the C2Dg ions 22 which can either react with C3D8C2D: +C3H6-,C2D5H+C3H5f (30)C2Di + C3H6+C5D5Hz. (31)On the basis of the above mechanism, approximate values of the relative rates ofthe different processes can be obtained from the yields of C2D6, C~DSH, and C2D4and their variations with C3H6 concentration. However, because a fraction ofC2D6 and C2D4 is also formed by reactions 23 other than (28) and (29), more accuratedetermination was based on the isotopic analysis of the ethane and ethylene formedin the radiolysis of C3Dg + C3H8 (1 : 1) mixtures in the presence of various amountsof C3H36 From these results it is calculated that the relative probabilities of reac-tions (28), (29), (30) and (31) are 1.00, 1.07, 0.16 and 0.93, respectively. Radiolyticstudies 3, 17 have shown that CZHf and n-C3Hf can also transfer a proton to otherorganic compounds such as CH30H and CH3COCH3.This research was supported by the U.S. Atomic Energy Commission.1 Ausloos and Gorden, J. Chem. Physics, 1964,41, 1278.2 Ausloos and Lias, J. Chem. Physics, 1964, 40,3599.3 Sandoval and Ausloos, J. Chem. Physics, 1963,38,2454.4 Stevenson and Schissler, J. Chem. Physics, 1955,23, 1353.5 Giese and Maier, J. Chem. Physics, 1961, 35, 1913.6 Ausloos and Lias, J. Chem. Physics, 1964,41,3962.7 Stevenson and Schissler, The ChernicaZ and Biological Action of Radiations (Academic Press,8 Thompson and Schaeffer, J. Amer. Chem. Soc., 1958,80,553.9 Schaeffer and Thompson, Rad. Res., 1959,10,671.10 Maschke and Lampe, J. Amer. Chem. Soc., 1964,86,569.11 Ausloos, Lias and Gorden, J. Chem. Physics, 1963,39, 3341.12 Chupka and Lindholm, Arkiv. Fysik., 1963, 25, 349.13 Rylander and Meyerson, J. Amer. Chem. SOC., 1956,78, 5799.14 Grub and Meyerson, Mass Spectrometry of Organic Ions (Academic Press, New York, 1963),15 Stevenson, Trans. Faraday SOC., 1953,49,867.16 Borkowski and Ausloos, J . Chem. Physics, 1963, 38, 36.17 Borkowski and Ausloos, J. Chem. Physics, 1964, 40, 1128.18 Ausloos and Lias, J. Chem. Physics, in press.19 Smith, Corman and Lampe, J. Amer. Chem. Soc., 1961,83,3559.20 Futrell and Tiernan, J. Chem. Physics, 1963, 38, 150.21 Lampe and Field, J. Amer. Chem. Soc., 1959,81, 3242.22 Ausloos and Lias, J. Chem. Physics, 1962, 36, 3163.23 Ausloos, Lias and Sandoval, Disc. Faraday Soc., 1963, 36, 66.London, 1961), vol. V, pp. 249-254.p. 518

ISSN:0366-9033    

DOI:10.1039/DF9653900036

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

GENERAL DISCUSSIONProf. J. J. Weiss (University of Newcastle) (contributed) : Several referenceshave been made to a mechanism where the proton transfer occurs along a hydrogenbond, a point which was emphasized particularly by Eigen. In water, this maygo with the intervention of a hydrogen-bonded water molecule whereas in aproticsolvents one should have the direct formation of a hydrogen bond between the tworeactants prior to proton transfer. If the proton transfer goes along a hydrogenbond, the distance which the proton has to travel is relatively short, viz., 0.4-0.6A.Thus, the width of the potential barrier which the proton has to penetrate wouldbe sufficiently narrow to allow processes of tunnelling to compete favourably witha classical mechanism of proton transfer.I am referring here to a real tunnel effectand not to the type of mechanism discussed by Caldin which is essentially a classicalone with a correction for tunnelling only at the very top of the potential barrier. Inthe simple model which I have used,l proton transfer is considered to take placealong a line joining the proton donor and acceptor hence it may be described by aone-dimensional barrier of given height and of width equal to the distance whichthe proton has to travel along the hydrogen bond. This problem, using an Eckartbarrier, which from a physical point of view is the most realistic one, can beformulated in such a way that one can do an exact calculation.The model leads to an expression for the rate constants which gives a reasonableaccount of the isotope effect.Moreover, a Bronsted type of relation can also beobtained from this theory without any further assumptions. As a first approxim-ation (i.e., if one takes only the first term of a particular series expansion) one obtainsfor the Bronsted coefficient a, the following theoretical expression :a = (1 - (2PkT/Q')),where p = (a/h> (2 M)*, where 2a denotes the barrier width and M is the massof the proton (or deuteron). Q is an energy, i.e., the difference between the protonaffinities of the donor and acceptor molecules. Thus the Bronsted exponent adepends also on the value of the energy difference Q. Approximate constancyof a within a certain series therefore presupposes an approximate constant valueof Q, otherwise there should be a variation of the Bronsted coefficient with thevalue of Q. Eigen has discussed the variation of the Bronsted coefficient with thepK difference of the proton donor and acceptor.Eigen's results are fully compatiblewith the above equation as Q, i.e., the difference in the proton affinities of donorand acceptor, is directly related to the difference of their pK values. Eigen'sdiscussion moreover shows that the current derivation of the Brijnsted relation haslittle significance from a theoretical point of view, based as it is on the a prioriassumption of the proportionality between the activation energies and the freeenergies, thereby virtually presupposing what is meant to be derived.Prof. F. A. Long (Cornell University) said: It is true, as Dr.Eigen has said,that the general acid catalysis for the detritiation of azulene-1-t follows the Bronstedrelation over a wide range of acidity. This is similar to the results obtained foressentially the same reaction by Kresge with 173,5-trimethoxybenzene-2-t. Theonly reservation is that the density of points from acids is not high and as a conse-quence, much of the behaviour hinges on the results for the two acids, H20 and1 J. Chem. Physics, 1964,41,1120.446 GENERAL DISCUSSIONH30+, acids which are difficult in Bronsted plots. We also have data giving a moreprecise study of catalysis for groups of acids 1 which show that the catalysis to somedegree depends upon charge type, hence it would be useful to have a greater densityof points for acids of identical charge type to support the statement of Dr.Eigen.Prof. A. J. Kresge (IlZinois Institute of Technology) said: For the detritiation of1,3,5-trimethoxybenzene-2-t, we fowd that the data conform to the Bronstedrelation very well over a wide range of reaction rate (fig. 1). For a spread of 9Ipowers of 10 in catalytic coefficient which corresponds to 17 pK units in catalystacidity, the standard deviation in rate constant from the usual linear log-log relationis only 30 %. On this scale the deviations which Prof. Long found would be small,and this may be the reason why our data appear to obey the Bronsted relation sowell.I would like to add that the reactions with which we are dealing are all slow:none of the rate constants approach the values expected for diffusion-controlledreactions.Therefore, the sort of deviations from the Bronsted relation whichProf. Eigen discussed would not be expected to occur here.Prof. M. Eigen (Giittingen) said: I should like to add a few remarks regardingthe discussion contributions of Lord Wynne-Jones, Mr. Bell, Prof. Long andProf. Kresge. The concerted or co-operative mechanism mentioned in my paperwas introduced in order to remove the inconsistencies occurring in the mechanismof consecutive acid-base action in prototropic changes. Such a consecutive mech-anism would require rate constants for the rate-limiting step which depend linearlyon the pK of the acid or base catalyst-even in the range of diffusion control, whereall directly measured rate constants of proton transfer become independent of pK.(In some cases they even exceed the limiting value of about 1010 M-1 sec-1 appreci-ably, not showing the asymptotic behaviour a-+O).While the concerted or co-operative mechanism must be effective in these cases, it does not necessarily apply1 Thomas and Long, J. Arner. Chern. SOC., 1964,86,4770GENERAL DISCUSSION 47to other cases where a constant a has been observed over a large pK-range. HOW-ever, a necessary condition for a constant a is that the rate constant of the actualprocess is sufficiently below the limiting value mentioned above. This is indeedtrue for all the mentioned reactions, in which one step of proton transfer is probablyrate limiting, especially in the isotope exchange of hydrocarbons discussed by Longand Kresge.The hydrolysis of orthoesters mentioned by Wynne-Jones shows O ~ Yacid catalysis, and inspection of the structure of these compounds suggests thatthere is no suitable place for an attack of the base in the initial (possibly rate-limiting)step. Nevertheless, the co-operative mechanism does not require a symmetricalbehaviour with respect to acid and base catalysis, although in principle both shouldbe possible. In general, all types of mechanisms might be found (e.g., pre-equilibrated protonation or deprotonation leading to specific H+- or OH--catalysis,rate-limiting proton transfer with or without pre-equilibrated protonation or de-protonation and co-operative proton transfers-the latter cases all leading to generalacid or base catalysis).The co-operative mechanism will be preferred only if thetwo sites on the substrate can interact favourably via the solvent structure. Theoxygen groups in the above-mentioned cases fulfil this condition. CK-groups,however-as involved especially in the isotope exchange of hydrocarbons or inketo-enol tautomerism (cf. fig. 26 in my paper) might disfavour the co-operativemechanism.Wynne-Jones also mentioned the fact that a correlation of activation energieswith AH is usually not as good as that of log k with AP (or pK). Since AF andAH, as well as the corresponding kinetic quantities, include contributions fromdifferent sources (e.g., formation of transition complex, solvation, proton transferin H-bond, etc.), one cannot expect that all these contributions are correlated bythe same factor a.Thus, a good correlation for any of the quantities can be ex-pected only for a series of homologous substances. A more detailed treatmentwould require a further knowledge about the structure of the transition state towhich isotope studies might contribute greatly.Prof. B. E. Conway (Ottawa) said: Prof. Eigen remarked on the conditionsunder which a in the Bronsted relation may be expected to be constant over a reason-able range of pK values. A constant a will be expected if the crossing region ofpotential energy curves for the proton transfer in the acid-base reaction is well aboveELECTROCHEM I CAL ACID- BASEAE*=ah(RTlnK1A (RTI R K 1RE ACTION COOR DI N ATEthe zero-point levels of the reactants and/or products.Some justification for thea priori expectation of reasonably constant a follows by comparison with the analogouscase for electrochemical proton transfer at a cathode. Here the energy of the initialstate is changed relative to that of the final state by an energy zFAY for a chang48 GENERAL DISCUSSIONof metal-solution p.d. of AY and the electrocheinical rate i is modified byexp[-zF/3AV/RT]. Hence In i is proportional to AV and a symmetry factorp(+ 0.5) analogous to Bronsted’s a is involved. The activation energy is modifiedby an energy zF’AV. The analogy between the two cases is shown schematicallyin fig. 1, for varying base strength of the entity B. At some electrodes, e.g., Hg,In i is exactly proportional to AY over 9 decades of rate (i.e., for AV changing byca.1 V) so that a similar effect with regard to changing pK values in the Bronstedrelation is entirely reasonable on the basis of potential energy diagrams.1 A changeof 1 V in the electrochemical case is equivalent in RT In K units to a change of about17pK units in the acid-base case, which is a wider range than that normally en-countered in real chemical cases.Mr. R. P. Bell (Oxford) said: Mention may be made of another reaction forwhich it has been claimed that the catalytic power of a series of acids ranging fromH20 (PK 15.75) to H30+ (pK - 1.75) can be represented by a linear Bronsted rela-tion. This is the decomposition of the diazoacetate ion in aqueous solution, studiedby King and Bolinger,2 for which the special mechanism suggested by Eigen for thehydration of carbonyl groups and related reactions does not seem appropriate.However, the dependence of velocity on catalyst concentration in the decomposi-tion of the diazoacetate ion is an unusual one, and the catalytic constants given byKing and Bolinger were derived by an arbitrary extrapolation procedure.It wasshown by Bell and McTigue 3 that the kinetics of this reaction can be interpretedquantitatively by assuming two consecutive steps of comparable rate, making itpossible to derive the velocity constants for the first step N2 : CHCO; +At-+N2 . CH2CO; +Bi, where Ai-Bi is an acid-base pair. (In addition to the valuesalready given we find kA = 1-4 x 104 M-1 min-1 for catalysis by acetic acid.) Therevised catalytic constants do not give a convincingly linear Bronsted plot over thewhole range, and in particular the considerably decreased value found for hydrogen-ion catalysis suggests a flattening of the curve of the kind found by Eigen for normalproton-transfer processes.A further experimental study of the diazoacetate de-composition would be desirable.It is worthwhile emphasizing that the concerted or co-operative mechanismfor proton transfer in aqueous solution is probably confined to a certain type ofreactions, and probably always involves one or more water molecules. Thus,there is no evidence for such a mechanism for keto-enol reactions in water, andno evidence for kinetic terms such as k[AcOH][AcO-] in the hydration of carbonylcompounds or the mutarotation of glucose.4 Additional evidence for the concertedmechanism in carbonyl hydration comes from the results of Strehlow 5 on the kineticsof the reaction MeCOC02H+H20+MeC(0H)2C02H7 which shows a “ spon-taneous ” rate which is much too large to be attributed to catalysis by solvent mole-cules.It was therefore attributed to intramolecular catalysis by the carboxylgroup, and this seems much more probable if the proton transfer takes place throughone or more intervening water molecules.Prof. A. J. Kresge (Illinois Institute of Technology) said: Mr. Bell’s calculationsand the earlier ones to which he referred all predict that the value of the isotope4-1 Bell, The Proton in Chemistry (Cornell Univ.Press, Ithaca, N.Y., 1959), p. 170.2 King and Bolinger, J. Arner. Chem. SOC., 1936, 58, 1533.3 Bell and McTigue, J. Chem. Soc., 1960,2983 ; see also Bell, The Proton in Chemistry (Cornell4Bell and Clunie, Nature, 1951, 167, 363; Proc. Roy. SOC. A , 1952, 212, 33.5 Strehlow, 2. Elektrochern., 1962, 66, 392.U.P., 1959), p. 136GENERAL DISCUSSION 49effect on a given reaction will pass through a maximum as the structure of the transi-tion state is varied from one extreme to the other. I now present some data whichseem to provide the first clear indication of such a maximum. The reaction isaromatic hydrogen exchange, and the isotope effect is that on bond-breaking inthe phenonium ion which is the intermediate in the reaction :kH kDH i- +ArD cHArD+ -+ HAr + D +TABLE 1Ic,lk, exchanging relative rateposition of exchange HArC6H6 - 1 3.4 h0.2C6H5CH3 3 6 3.4 f0.2 *C6HSCH3 2 4 x 102 4-6 f0.4 aC6HSCH3 4 4 x 102 5.5 f0.3 aC6H50CH3 2 2~ 104 7.2C6H50CH3 4 6x 104 6.71 ,3,5-C6H3@CH3)3 2 1 x 1010 6.7 f0-2 Cazulene 1 3x 1011 5.6(a) Olsson, Arkiv.Kerni., 1960, 16, 482. (c) Kresge andChiang, J. Arner. Chern. Soc., 1962, 84, 3976. ( d ) Schulze and Long, J. Arner. Chem. Soc., 1964,86, 331.(b) Russell, M., private communication.These experiments cover a wide range of reactivity and presumably, therefore,a considerable variation in transition state structure. Other data, such as thosepresented by Mr. Bell, are more limited, and this may be the reason why they showonly an upward or downward trend in isotope effect without a maximum.Mr.R. P. Bell (Oxford) said: I should like to amplify two points made in mypaper. First, the conclusions embodied in table 3 are supported by calculationsrecently published by Willi and Wolfsberg,l who examined the dependence of iso-tope effect on the symmetry of the transition state for various values of the barriercurvature. For zero curvature (corresponding to Westheimer’s treatment) there isa fairly sharp maximum in kH/kD for a symmetrical transition state, but for realisticvalues of the curvature this maximum becomes very flat and kH/kD is almost con-stant over a wide range of transition states: this corresponds to the decrease inA1, and the approach of v?/v? to unity shown in my table 3.Secondly, the correlation between kH/kD and p shown in table 2 is improvedif a correction is applied for the secondary isotope effect of the non-ionizing deuteriumatoms in CD2- and CD3-groups.Streitwieser and van Sickle 2 have found that thesubstitution of two deuteriums in the methyl group of toluene diminishes the rateof exchange of a third deuterium by 24 %, i.e., there is a secondary isotope effectof kH/kD = 1.1 5 per deuterium, and in my laboratory Mr. D. M. Goodall has founda similar value for the ionization of nitroethane by hydroxide ions. It is thereforereasonable to correct the observed values of kH/kD by dividing by 1.15 for CD2-groups and by 1-152 for CD3-groups: these corrected values follow the samesequence as /? with the single exception of acetylacetone.One further value maybe added to table 2 by using the observation 3 that k;/kD = 7.7 for the brominationof acetone catalyzed by &0+. The rate-determining reaction is Me& : OH+ + H20,and the appropriate value of p is 1 - a, where a is the Bronsted exponent for acidcatalysis, having the value 4 0.62.1 Willi and Wolfsberg, Chern. and Ind., 1964, 2097.2 Streitwieser and van Sickle, J. Arner. Gem. SOC., 1962, 84, 254.3 Reitz, 2. physik. Chern. A , 1937,179,119. Reitz and Kopp, 2. physik. Chem. A , 1939,184,429.4 Bell, Acid-Base Catalysis (Oxford, 1941), p. 9150 GENERAL DISCUSSIONDr. D. B. Matthews (Univ. of Virginia) said: With regard to the interrelationshipof the Bronsted coefficient, the isotope effect, activation energy and heat of reaction,I shall employ the potential energy profile method of Horiuti and Polanyi :REACTION COORDINATE ( x ) -Change in configuration of the activated state produced by a change in heat of reactionwhere I refers to the initial state and I1 the final state.A change in heat of reactionwill cause a vertical shift of one curve with respect to the other (dotted curve).The effect is a change in activation energy, and also in configuration of the activatedstate with respect to the initial or final state configurations, which causes a changein zero-point energy of the activated state and hence a change in the isotope effect.This direct dependence of the isotope effect on the heat of reaction is often overlookedin comparison with, say, the role played by charge on the acceptor molecule and therole of proton tunnelling. My second point concerns the dependence of the degreeof proton tunnelling on the heat of reaction.Mr. Bell has shown that the isotopeeffect for proton transfer as a function of charge on the acceptor goes through amaximum. Similarly it is found that the dependence of the isotope effect, due toproton tunnelling, on heat of reaction goes through a maximum, the position ofthe maximum corresponding to AH0 = 0.Dr. N. A. J. Rogers (Birmingham University) (communicated) : I should like torefer to a point raised by Dr. Matthews, who pointed out a property of potential energyprofiles, illustrated in fig. 1.He argued that in a situation in which an initial stateFIG. 1.B can lead, via a transition state, TS or TS’, to two final states, A or A’, then thealtering of the final state from A to A’ will not only affect the activation energy ofthe reaction, but will also alter the configuration of the transition state. Thisconsequence is clearly shown in fig. 1. I should like to refer to some of our work,in which the concept appears to be of importance. We have investigated thGENERAL DISCUSSION 51protonation of a series of conjugated dienol ethers of the types I and 11, in which theattached groups are alkyl or hydrogen.We have, in the first instance, studied the preferred site of protonation ( a or y)of these compounds. Broadly, it appears from our results that the transoid-dienolethers (11) are less sensitive in their behaviour to alkyl substitution than are thecisoid-isomers (I).Thus the transoid-dienol ethers examined to date protonate\/I I1exclusively at the y-carbon atom, yielding the more stable, conjugated products.Further, a linear plot of log (k,/k,) against (qa-q,,), the charge densities as calculatedby the simple Huckel m.0. method, has been obtained for the cisoid-dienol ethers.In attempting to rationalize these results, we have found it necessary to invokethe idea of a difference in relative configuration of the transition states for the pro-tonation of these two series of compounds. There appears to be a greater degreeof bonding in the transition state for protonation of the transoid compounds, wherethe more stable products are formed, than in that for protonation of the cisoidisomers, where dependence on a property of the initial state (Aq) is observed.For the reactions below,a cisoid-(111) and a transoid-(1V)-dienol ether give rise to the same product (V)on y-protonation.From our earlier arguments, the initial state for the protonationenergy profile of (111) should lie at a higher level than that of (IV). We havemeasured the equilibrium between I11 and IV and have found that IV is the morestable by about 1 kcal/mole. Thus, on Dr. Matthew's diagram, fig. 1, I11 cor-responds to A', IV to A and V to B.Secondly, Prof. Havinga referred to his studies of the protonation (deuteronation)of the excited state of anisole (I), in which 0- and m- attack is observed, the relativeI I1reactivities at these positions correlating with the calculated charge densities.Wehave been studying the formally similar anion-radical of anisole (11), the calculatedcharge densities of which would predict protonation at the o-position preferentially.This protonation is the crucial step in the metal/ammonia reduction of anisole52 GENERAL DISCUSSIONPreferential protonation at the 0- and at the m-position have been strongly sup-ported by different groups. We now have experimental results, based on an e.s.r.study of the anisole + dimethoxyethane +potassium system, which are most readilyinterpreted in terms of preferential o-protonation of 11. This result, like those men-tioned by Prof.Havinga for I, is the one predicted on the basis of charge densitycalculations.Dr. M. Fleischmann (Newcastle upon Tyne) said: If the viscous relaxation timecan be shown to affect the rate of proton transfer in examples where the solvent isnot directly involved, it would follow that the classical transfer would have to betreated as if the system were non-conservative.1 Both the “ frequency factor ”and the “ energy of activation ” will then depend on the velocity of the particlesinvolved in the transfer, low velocities giving the highest probability of surmountingthe barrier. One would predict that the ‘‘ reaction co-ordinate ” would involvethe movement of the whole molecules or sections of the molecules in the solventcage rather than the stretching of single bonds.In one particular example of proton transfer it is now possible to state thatsolvent reorientation is unimportant.The electrochemical hydrogen ion dischargeon mercury has been measured up to heterogeneous rate constants of the orderI000 cm sec-1.2 The logarithms of these high values of the rate constants plottedagainst eIectrode potential fit on the same “ Tafel line ” as the values measured atless negative potentials. Since the velocity of the nuclei at the high rate constantsis certainly in the range or above that corresponding to the viscous relaxation timeof water, it may be concluded that the reorientation of the solvent is unimportantin this particular discharge reaction.Dr. H. W. Numberg (Kernforschungsanlage Jiilich) (communicated) : Employingthree advanced techniques of polarographic nature (pulse-polarography (PP), square-wave-polarography (SWP) and especially high level faradaic rectification (HLFR) 3),which permit reduction of the measuring time t1 after the start of polarization of theelectrode to finally 1 psec, we have been able to study at 20°C the kinetics of thehydrogen evolution at the mercury electrode up to very negative potentials and con-sequently up to high rate constants for the charge transfer step of this electrodeprocess.4Over the whole range of measurements leading from kct = 5 x 10-5 cm sec-1 at- 1.32 V (SCE) to kct = 800 cm sec-1 at -2.1 1 V (SCE) the results fit to a Tafel-line with a slope of b = 110 mV indicating a constant apparent charge transfercoefficient of aa = 0.53 (fig.1). No double-layer correction has been applied,but as the solutions contained always a constant high supporting electrolyte con-centration of 1 m LiCl the influence of the double-layer effects on b over the wholemeasured range of potential remains quite small (<lo %) and will not exceed 3 %The data between kct = 100 and 800 cm sec-1 could only be obtained withacetic acid making use of the fact that due to the participation of the prior homo-geneous chemical reaction of the dissociation of the weak acid the current-voltage-curve rises less steeply and reaches therefore the limiting current at a more negativeelectrode potential than when a strong acid (HCl) is employed as proton donor.After the homogeneous dissociation rate constant of acetic acid had beenfor aa.1 cp.Bass, Proc. Roy. SOC. A, 1964, 277, 125.2 Barker, Niirnberg and Bolzan, Report Jiil-l37-CA, 1963, Kernforschungsanlage, Julich.4 (a) Barker, Numberg, and Bolzan, Report JiiZ-l37-CA, (Kernforschungsanlage Julich, 1963).Barker and Numberg, Naturwiss., 1964, 51, 191.(b) Barker, Niirnberg, Bolzan and Gardner, Electrochim. Acta, in pressGENERAL DISCUSSION 53determined from the limiting current 1 the charge-transfer rate constant kct could beevaluated from the rising part of the polarogram.The dotted part of the Tafel-line in fig. 1 leads to the most negative potentialobtained in the limiting current region of the polarograms before the discharge ofthe Lif ions occurs.An extrapolation to this potential seems justified if one assumesthat a, remains constant.Several general conclusions relevant to a number of points made by variousauthors at this discussion are possible from our results.(i) Our measurements are concerned with a range of charge-transfer rate constantsand corresponding potentials above that region normally hitherto studied as it isnot accessible by the more conventional techniques due to their larger measuringtimes. Exceptions reaching the lower part of our Tafel range are the experiments10' -10' -loz -10' -.4I 8 l -1 0 ' -10' -1 0 ) -Id' I'amHLFR f-- 4"P*HAC ///S WP, I, = 2.22 rnsec. HClPP. I, =40 msec,HCl-1.3 -1.4 -1.5 -1.6 -1.7 -1.8 -1.9 -1.0 -1.1 -2.2 -23voltsFIG.1 .-Tafel plot for h.e.r. at Hg from 2.5 x 10-4 m HCI, 1 m LiCl and 1 x 10-3 m HAc, 1 x 10-2 mNaAc, 1 m LiCl respectively at 20°C (from ref. (24). The upper measuring limits of the employedtechniques due to their respective measuring time t l and to the respective experimental conditions(strong or weak acid) are indicated. The Tafel-line has been smoothed through a great number ofexperimental points, which had to be omitted because of the small dimensions of the figure.by Bockris and Azzam.2 However, the data obtained by conventional methodsat lower potentials may be fitted fairly well to a Tafel plot of the same slope.3 Con-sequently, for the h.e.r. at mercury from acidic aqueous solutions a constant slopeb and therefore a constant charge-transfer coefficient a over a range of more than1.9 V corresponding to 17 powers of ten in the charge transfer rate constant kot isexperimentally established.This behaviour is consistent with theoretical consider-ations of Christov 4 predicting for an Eckart barrier with a height of Er = 1.5 x 10-12erg at q = 0 a constancy in a over more than 2 V of overvoltage q (see p. 128 in ref.1 Nurnberg, in Proc. 3rd Znt. Congr. Polarography, Southampton, 1964), ed. Hills (Madllan3 cf. Vetter, Elecirochemische Kinetik, (Springer, Berlin, 1961), p. 432.4 Christov, Ber. Bunsen. physik. Chem., 1963, 67, 117.Ltd., London, 1965). 2 Bockris and Azzam, Trans. Fmahy SOC., 1952,48, 1454 GENERAL DISCUSSION(6)). Thus the h.e.r. on mercury may well be regarded as the electrode process onwhich the most extended knowledge with respect to charge transfer rate constantsand charge transfer coefficient is available at present.(ii) With respect to the comparison by Prof.Conwayl between the constancyof the Bronsted coefficient over a certain pK-range and the constant behaviour ofthe charge-transfer coefficient for the h.e.r. at mercury over a given range of elec-trode potentials, the figures have to be enlarged appreciably including our results?In the 2RT In K scale the range of 2 V for which a constant U-value has been ob-served is equivalent to 30.5 pK-units which is a much wider region than generallyobserved in homogeneous proton transfer cases.(iii) One deduces from fig. 1 further that there is no indication for any fast priorchemical reaction in the discharge of hydrogen ions from a strong acid as HCl.On the other hand, we have observed for a number of metal ions reductions studiedat the mercury electrode with the new technique of high level faradaic rectificationa normally fast prior chemical step which is to be attributed to partial dehydrationor decomplexation of the metal ion before electron transfer occurs.2as 2 The absenceof this step in the discharge of hydrogen ions is to be expected by analogy with themechanism pointed out by Eigen 3 for the homogeneous recombination of the H3O+ion and the anion of an acid (see also my discussion remark on the paper of Salomonand Conway and that of Bockris, Srinivasan and Matthews).(iv) Several authors4 have raised the question of the relations between therelaxation time for the reorientation of water molecules and proton transfer.Ourstudy on the h.e.r. at mercury has been cited as furnishing evidence for a protontransfer process where solvent reorientation is unimportant (cf. discussion remarkof Dr. Fleischmann) as well as showing that the dielectric relaxation time of waterseems to be smaller than the time for proton transfer via the interphase electrode/solution (cf. discussion remark of Prof. Christov). A detailed inspection of theproblem reveals that both statements are not strictly correct.Taking our highest (extrapolated) kct value of 2 x 104 cm sec-1 in fig. 1 and assum-ing a " reaction layer " pct for the charge transfer step of only 0.5 A 5 an equivalenttime zct = 2.5 x 10-13 sec is obtained.The dielectric relaxation time z * for H20dipoles in 1 m aqueous alkali halide solutions has been determined 6 as 76 = 9 x 10-12sec at 20°C and is thus a factor 40 larger than our smallest zct. However, the neces-sary rotation of a water molecule into a position favourable for proton transfer isaccelerated by the field of the H3O+ approaching the water molecule.7 For therelevant rotation time in proton transfer one has 'trot = 2 x 10-32,. Regardingfurther that on the average only 1 of 9 rotations leads to a position of the H20molecules adjacent to an H30+ favourable for proton transfer, the effective rotationtime becomes zfrot = 9 x 2 x 10-3 z, as the proton transfer step itself along thehydrogen bridge formed to the favourably orientated H20 molecule occurs in atime negligible with respect to zfrot (see, e.g., ref.(13)) and the measurements on theproton mobility in ice by Eigen and collaborators 8). Inserting these figures one* The viscous relaxation time (cf. Discussion remark Dr. Fleischmann) may be identified with7 0 for practical purposes.1 cf. Discussion remark of Conway.2 Barker, Nurnberg and Gardner, 13 CITCE-Meeting, Rome, 1962 ; Electrochim. Ada, in press.3 Eigen, 2. physik. Chem., 1954, 1, 154.5 (a) Salomon and Conway, this Discussion. (b) Nurnberg, this Discussion.6 Lane and Saxton, Proc. Roy. SOC. A , 1952,214, 531.7 Bockris, Conway and Linton, J. Chem. Physics, 1956,24, 834.8 Eigen and De Maeyer, 2.Elektrochem., 1956, 60, 1037. Eigen, De Maeyer and Spatz, Ber.4 cf. Bell, this Discussion.Bunsen. physik. Chem., 1964, 68, 19GENERAL DISCUSSION 55obtains drat = 1.62 x 10-13 sec. Thus, the Tct value equivalent (with the assumptionfor pet) to our largest (extrapoIated) value of the charge transfer rate constant kctis approaching the order of the effective orientation time Trot' for the H20 moleculesin proton transfer. Though the reorientation of the H20 molecules during protontransfer is therefore not affecting our present results the design of such experimentsseems feasible in principle.(v) In terms of the paper of Salomon and Conway 1la and an earlier published cri-terion by Conway 1 the slope of our Tafel-line in fig. 1 indicates no significant contri-bution of proton tunnelling to the charge transfer step while according to Bockris,Srinivasan and Matthews 2 a moderate tunnel effect cannot be excluded.Dr.Roger Parsons (University of Bristol) said: Some light may perhaps bethrown on the problem of the reorientation of adjacent solvent molecules duringproton transfer by asking the question, " Why are proton transfers fast comparedwith electron transfers? " It is known from the theories of Marcus 3 and Hush 4that homogeneous electron exchange between hydrated transition metal ions andthe electron transfer is retarded by the necessity to reorganize the surroundingsolvent into a configuration intermediate between initial and final states. Thiscan lead to energies of activation in the region of 9 kcal mole-1 for an ion of thesame size as H30+.The absence of such activation energies in proton transferreactions (except when they are endothermic) seems to suggest that solvent reorgan-ization does not play an important role probably because the distance over whichthe proton is transferred is small compared with the electron jump distance and thefield on the surrounding solvent is, relatively, much less altered as a result of thetransfer.Prof. M. C. R. Symons (University of Leicester) said: Although, as stressed byDr. Parsons, the barrier to proton transfer caused by the need for solvent reorgan-ization is likely to be small in comparison with that for many electron-transferprocesses, since the distance through which the proton needs to move is small, thiswill only be true of overall reactions in which the rate-determining step is the transferof a proton.There must be many proton-transfer reactions in solution for whichthis is not the rate-determining step, in which case the solvation " barrier " may beextremely important. Various possibilities may be illustrated by the followingmodel, which also serves to underline in a simple manner several other points raisedduring this discussion.Consider the symmetrical transferThis may be conventionally represented by an energy diagram which showsmovement of the proton as it moves across from one A- to the other (fig. I). How-ever, such curves imply a fixed A - - - A distance, so one needs to draw up a familyof curves for different A - - - A distances.This can be done by imagining a seriesof curves, of the type shown in fig. 1, through each point on the curves given infig. 2, where energy is plotted against the A - - - A distance.Initially, as A- approaches AH, the proton will remain bonded to its initialpartner, and the A - - - H bond length will hardly alter. There will, however, bea solvation barrier to overcome, which could be pictured as a replacement of one ofthe solvent molecules associated with A- by the acid HA. If the system [A - - - H - - - A*Isil,. has some stability, as is often the case, the curves in fig. 2 will now fall.1 Conway, Can. J. Chem., 1959,37, 178.2 Bockris, Srinivasan and Matthews 1 his Discussion.3 Marcus, J. Chem. Physics, 1956, 24, 966 ; for general review see Marcus, Ann.Rev. Physic.A,,,. + HA+(A- - -H- - -A),,,.+AH +A~'lvv. (1)Chem., 1964,15, 155. Hush, J. Chem. Physics, 1948, 28,96256 GENERAL DISCUSSIONOne extreme is that from this minimum there is still a large barrier to protontransfer, so that solvent reorganization will be only a minor effect (cf. Dr. Parsons’explanation). This would be the case, e.g., if the preferred path were a combinationof curve (3), fig. 2, and curve (3) of fig. 1. An important alternative is that, withinthe solvated complex [A - - - H - - - A]& proton transfer is rapid and hence the( A-H -- A )-FIG. 1.( A - - - - A ) h n(A---A) distanceA-SOLV- -- HA ( A-- - H ---A )-SOWFIG. 2kinetic barrier is the initial solvation barrier discussed above.This would be thecase for curves (1) of the figures. One example of this latter situation is the exchange(1) in which A- is F-. Here the HF; ion is stable, and its dissociation will resultin an equal distribution of protons between the two fluoride ions. The same modelmay be used to illustrate the situation envisaged during the discussion, in which aproton is “ pushed across the barrier ” by the group to which it is attached. Thismeans that the A - - - A distance most suitable for proton transfer is less than theequilibrium distance for normal hydrogen bonding. Jn general, the solvent barrieGENERAL DISCUSSION 57considered here will control the rate if xg>yg+zgxd, etc., being the appropriateenergies shown in the figures.The symmetrical exchange used as a model has the advantage that the principleof microscopic reversibility is naturally accommodated.For unsymmetrical ex-changes barrier heights will vary in the manner outlined in the discussion by variouscontributors, but the factors stressed here, viz., (i) the influence of a stable hydrogen-bonded complex; (ii) the need for solvent reorganization during its formation andloss ; (iii) the operation of the principle of microscopic reversibility ; (iv) the needto consider the A - - - A (or A - - - B) distance as well as the A - - - H distance,remain worthy of consideration.Mr. R. P. Bell (Oxford) said: The width of the parabolic barrier at its base isperhaps a rather artificial quantity, since the parabolic approximation is only validnear the top of the barrier.Since the reactions so far investigated involve only arelatively small tunnelling correction (EAIEqmt = 0.79-0.95) the quantity which isdirectly derivable from the experimental data is really the curvature of the top ofthe barrier, and this might be preferable to the width as a basis of comparison.(In terms of the parameters used by Caldin and Kasparian the curvature is 2Eqmt/a2,and it can also be written as 4n%~v$, where m is the reduced mass and iv* theimaginary frequency). At a given temperature the curvature is directly related to&/Eqmt, and the distinction between aqueous and alcoholic solutions, mentionedby Caldin and Kasparian, appears more clearly in terms of curvatures than in termsof barrier widths.Prof.J. J. Weiss (Newcastle) (communicated): I should like to comment onthe question of the barrier width in proton transfer reactions. The distance whichthe proton has to travel along the hydrogen bond between the two reactants wouldbe the mean distance between its initial and final state and this can be deducedfrom known bond distances. The distance of the centres in hydrogen bonds is2.6-2-8 A; as the OH or NH bond lengths are about 1.0-1-1 A, the actual distancewhich the proton has to travel along the hydrogen bond will be 0.4-0-6A. This,however, must not be confused with the closest distance of approach of the reactantsin the initial state (collision diameter) which, on the basis of these figures would benot less than 1.5-1.7 A.These distances cannot, however, be defined very pre-cisely since the molecular surfaces are not infinitely sharp.Prof. M. M. Kreevoy (University of Minnesota) said: Dr. Paul Steinwandand I have recently investigated the reaction of allyl-mercuric iodide with aqueousacid in the presence of traces of iodide ion. The reaction proceeds as shown:H+I -CH,=CH--CH2HgI+CHSCH=CH2 + HgI2.The rate is independent of iodide concentration and the rate-determining step isproton transfer from H+ to the y-carbon. We have studied the primary isotopeeffect on this reaction by a competition method as a function of temperature. Fromthe isotopic difference in the pre-exponential factors we have evaluated the widthof a hypothetical parabolic barrier as 1.3 A.Regardless of its exact physical sig-nificance, this result certainly supports the Caldin and Kasparian generalizationthat the " width " estimated in this way is primarily a function of the solvent. Thereaction is of quite a different type from those reported in table 3, but the width isvery similar to the others reported for aqueous solution.Also, it seems that there is a close relation between proton tunnelling and un-equilibrated transition states. It is the essence of tunnelling that it is very fast.Although the proton apparently moves only about 1 A, the centre of positive charg58 GENERAL DISCUSSIONmoves several times that distance. It seems unlikely that much solvent re-orientationcan take place during the tunnelling. It then follows that the solvent shell will beunequilibrated during much, and perhaps all, of the transfer.I would also like to ask if anyone has thought about tunnelling in the contextof a double jump mechanism :HIA-@ -0-0 +B.Dr.J. R. Jones (Battersea College of Technology, London) (communicated) : Thereare a number of points in Dr. Caldin’s paper upon which I would like to comment.Table 3 contains seven reactions which have been considered in terms of non-classicalbehaviour either by using isotopes or by observing curvature of the Arrheniusplot, and in some cases which involve the use of different catalysts. We believethat curvature alone cannot be taken as conclusive evidence for the existence oftunnelling and would suggest that it should only be used as confirmation when atleast one isotope effect is known together with the activation energy differencesand frequency ratios.This point is important for di-isopropyl ketone where, ifHulett’s account is accepted, a large isotope effect should be observed togetherwith an abnormal ratio for A D / A ~ . Our data for the rates of detritiation showthat ET-EH = 500 cal/mole, AT/& = 1-16 and K H / K ~ = 2-0 at 20°C.Secondly, the calculations of the energy barriers and widths all involve theassumption of a one-dimensional barrier of a symmetric parabolic shape. As manyof the substrates have widely different structures and the kinetic isotope effectsoften cover a wide range of values (cf. the value of 2.0 for di-isopropyl ketone andan equivalent value of 6-0 for acetone), this method is limited and discussion of smalldifferences in barrier widths must be of doubtful value.In the same way as isotopicsubstitution can be used to eliminate a number of uncertainties, then a study of anumber of structurally similar compounds using the same catalyst should be ofconsiderable benefit.Dr. J. R. Hulett (University of Leeds) said: I am interested in the table of barrierwidths given by Dr. Caldin and Dr. Kasparian in their paper. As we gain in-formation upon barrier dimensions, I believe we can use it to decide betweenpossible reaction mechanisms. Since I wrote my review 1 I have examined moreclosely the paper by Stewart and van der Linden.2 Although the Arrhenius para-meters for the permanganate oxidation of PhCH(OWCF3 and PhCD(OH)CF3are not given, they may be calculated from the rate constants shown at threetemperatures for the reaction at pH 13.0.From these results, I find E D - E ~ = 2290 cal mole-1 and AD/& 3-03.These figures are consistent with the author’s suggestion that the large isotope effect,k ~ / k ~ = 16, may be due to tunnelling.I have used Bell’s equations 3 for a para-bolic barrier and find = 11.42 kcal mole-1, ED(^, = 12-73 kcal mole-1, a =0.553Three features call for comment: (i) the degree of barrier penetration issimilar to that found for protons in the bromination of various ketones; (ii) thedifference in barrier heights, 1.3 kcal mole-1 is much the same as the difference inzero-point energies for the stretching modes of C-H and C-D bonds.Thisgiving EH/EH(*, = 0.85 and ED/ED(~, = 0.95.1 Hulett, Quart. Rev., 1964, 18, 227.2 Stewart and van der Linden, Disc. Furuduy Suc., 1960, 29, 211.3 Bell, Trans. Furaduy Soc., 1959, 55, 1GENERAL DISCUSSION 59implies almost complete rupture of the C-H bond in the transition state, unlessthere are considerable contributions from the bending modes. This result may becontrasted with those found for ketone bromination 1 ; (iii) the barrier width issimilar to those observed for proton transfer between an uncharged body and ananion in aqueous solution. Stewart and van den Linden suggest several possiblemechanisms for the reaction. Two involve termolecular collisions of solute species,of which one, the unionized alcohol, is present in only very small proportions atthis pH.- AS* for this reaction is probably too small to support such an unfavour-able process. The two other possible mechanisms are (a) hydride ion transferbetween the alcoholate anion and a permanganate ion or (b), simultaneous protontransfer from the alcoholate anion to water and electron transfer to a permanganateion. The authors favour the first of these mechanisms, although the very smalleffect of substituents in the aromatic nucleus is more easily accommodated by thesecond.The barrier width may help to decide between these alternatives. Mechanism(a) requires the approach of two negative ions and the transfer of a negativelycharged species between them. The electrostatic repulsions should lead to anenergy barrier considerably wider than those obtained for proton transfer betweenan anion and a neutral molecule.This is not the case, and thus the hydride iontransfer mechanism seems unlikely. The calculated barrier width, however, isconsistent with mechanism (b) if the major effect of deuterium substitution is onthe proton transfer, and secondary effects on the electron transfer are small. Atthis point we must leave aside any question of electron tunnelling. Althoughconsideration of the barrier width does not definitely establish mechanism (b),it seems most probable that the reaction involves a proton transfer between ananion and a neutral species-certainly mechanism (a) is unlikely.Prof. G. J . Hills (Southampton University) said: Of the arguments for protontunnelling, the non-linearity of the Arrhenius relation is perhaps the weakest.Thenormal enthalpy of activation defined by the isobaric relationship,(1)consists of two related components, viz., the internal energy of activation AU*and the product of the volume of activation and the so-called internal pressure ofthe system, i.e.,AH' = RT2(a In k/aT),,aTAH: = AU,f+-AV*, (2) Bwhere a and /3 are the coefficients of cubical expansion and compressibility, and AU*is defined by the isochoric relation,AU* = RT2(a In k/aT),,and AV* by the isothermal relation(3)AV* = RT(d In k/dP),. (4)The quantum-mechanical aspects of the reaction will determine the AU* termand thus may contribute to the temperature-dependence of AH*.However, a,/3 and A V are all temperature dependent and, more important, they and possiblyalso AU * are volume or density dependent. Since density is temperature-dependent,the isobaric variation of AH* with temperature is to be expected as a thermodynamicconsequence of eqn. (2). The question then arises as the magnitude of the termccTAV*/P. In aqueous solutions at room temperatures, it is small because a is small1 Hulett, Proc. Roy. Soc. A, 1959, 251, 27460 GENERAL DISCUSSION(zero at 4°C). At other temperatures and in other solvents it is large (1-2 kcal mole-1)even for modest values of AY*. The isobaric temperature dependence is then acomplex quantity even in classical terms and as a criterion of detailed aspects of areaction mechanism, should be used cautiously.Dr.J. R. Hulett (University of Leeds) (contributed): I have tried to apply theequationAH: = AUz+(aTAV:/P>to my results for the bromination of di-isopropyl ketone,l assuming that AU: istemperature independent. Preliminary calculations suggest that the experimentalresults cannot be reproduced by this equation unless AVT varies quite considerablyin the temperature range used (0-SOOC).Prof. B. E. Conway (Ottawa) said: I would like to comment on the questionof barrier widths 2a in proton transfer reactions, e.g., as deduced in the paper ofCaldin and Kasparian and referred to elsewhere in several papers in this Discussion.Most values of 2a seem too large compared with the real barrier width which wecould define as the mean distance between the proton in its initial and final state,and which might be deduced from molecular radii and covalent bond distances.A good example is the autoprotolysis in ice; here the 0-0 distance is relativelyfixed and is close to 2.8 A and the OH internuclear distance in the initial state is0.98 A.After proton transfer, an ion pair OH-H30+ is temporarily created andthe distance the proton is transferred can hardly be greater than about 2.8 - 2 x 0.98 A,i.e., 2a = 0.84A. Similarly, in proton conductance in acids, where the initial andfinal states are identical, 2a f O-SA. The deduced barrier widths, which are largerthan these figures by a factor of 2-3, probably reflect a degree of inapplicability ofthe tunnelling permeability equations as expected; for the Eckart barrier case, agreater “ barrier width ” will be required to reproduce a given curvature at thetop of the barrier, as remarked by Mr.Bell, than would be the case for the parabolictype of barrier model. The real barrier width will probably depend on the strengthof the hydrogen bond between the acid-base pair between which proton transferoccurs.In the model of Bockris, Srinivasan and Matthews, however, the barrier widthof ca. 4 tf seems too large on any basis, e-g., as indicated by Christov’s calculations.2Prof. J. O’M. Bockris (University of Pennsylvania) (communicated) : The Eckartwidth barrier of our calculation is not an assumed one, as is implied by Prof. Conway.It is that which is indicated by the only interpretation which we can make quanti-tatively consistent with the dependence of separation factor on potential (and thevalue that we have used comes out very near to that of Caldin and Kasparian).3Prof.S. G. Christov ( h t . Physic. Chem., Bulgaridn Academy of Sciences) said:The role of the tunnel effect in the proton-transfer processes in solutions and onelectrodes has been investigated by Bell 4 and the writer.5~ 6 Essentially the samemethod for the determination of the dimensions of the potential barriers was appliedindependently in both cases. This method consists in inserting the experimentaldata for reaction rate 0, activation energy E’ and frequency factor K’ in the rate1 Hulett, J. Chem. Sac., 1965, 430.2 Proc.1st Australian Con$ Electrochemistry, ed. Friend and Gutmann (Pergamon Press, 1965),4 (a) Bell, Trans. Fmahy Suc., 1959, 55, 1. (b) Bell, Fendley and Hulett, Proc. Roy. Soc. A,5 Christov, (a) Dokl. Akad. Nuuk. S.S.S.R., 1959,125, 141 ; (b) 2. physik. Chem., 1960,214,40;6 Christov, 2. Elektrochem., 1958, 62, 567.p. 723.1956,235,453.see also (c) Electrochim. Acta, 1961, 4, 306 ; 1963, 9, 575.3 Caldin and Kasparian, this DiscussionGENERAL DISCUSSION 61equation o = K’ exp (-E’/kT) for two isotopes (H and D or T). The quantitiesEk, EI;, KAY KI; (or KH/KA), VH and UD (or OH/UD) are in principle functions of thebarrier width zo and barrier height Ec and also of the corresponding zero-point energies(80 and 60) and entropies (SO and SO) in the initial and the transition states.Thesolution of a system of many transcendental equations by fitting is difficult, but areduction of the unknowns is possible 6 ; in absence of tunnelling we obtainthe relation 0.5 < K b / K h <2 for the ratio of the classical frequency factors, includingthe corresponding activation entropies (So - Sd) ; (ii) by introducing the “ true ”classical activation energies EH and ED (E = E,-(Q-&)) as unknowns. Ac-cording to (i) a low value for the tunnel correction is obtained by assuming Kb /K$ < 1.Accordingly, the decrease of the barrier height (E<E, because he0 = EO-EO>O)leads to an underestimation of the tunnel correction, too. It seems possible in thismanner to determine a lower limit for the net tunnel correction, which is thereforenot cancelled by the contribution of Aeo.The simplest and most direct way for estimating the role of the tunnel effectconsists in the evaluation of the characteristic temperature TK, defined by the con-dition, that the probabilities of transferring through the barrier and over it are equal2For smooth barriers this temperature is given by the relation 1TK = h JG/2nkJ12m, L, = -(d2V/dx2)x=xm,in which rn is the mass of the particle, h the Planck constant, k the Boltzmann con-stant and Lm the barrier curvature at the maximum.Expressions for TK for variousbarrier shapes were derived previously.4 This characteristic temperature permitsthe determination of the temperature ranges 5 : T> 2 T X (negligible tunnelling),TK < T< 2Tx (weak tunnelling), TK/Z < T< TK (moderate tunnelling) and T< T K / ~(large tunnelling).For each of these regions the tunnel correction varies betweenwell-determined limits and corresponding approximations are applicable.2A lower limit for TK can be obtained by insertion for L, the lowest value, corres-ponding to a parabolic barrier with the greatest possible width (e.g., double layerthickness) and the smallest possible height E = E‘ (E’ = experimental activationenergy). Calculations show 3 that proton-transfers normally occur in the regionof moderate or weak tunnelling.Under these circumstances it is not important whether we assume a symmetricalor an unsymmetrical barrier. Calculations of the barrier dimensions on the basisof experimental data give only a barrier equivalent to the real one.29 6 This meansthat both barriers have almost the same permeability in the given temperaturerange, which is possible under the condition that their upper permeable parts nearlycoincide.In this way we obtain different barrier widths for different equivalentbarrier models.2, 6 The effect of non-zero reaction heat (unsymmetrical barrier 39 4)may be important for large tunnelling which normally does not occur in proton-transfer processes. (This seems also to be the opinion of Mr. Bell.)The question, is there justification for the application of a one-dimensionalbarrier in calculating the tunnel corrections, is an important one. It appears,however, that the treatment of Johnston and Rapp 5 does not give a final answer,because they applied the Eckhart potential in evaluating the tunnel corrections fordifferent profiles of the potential energy surface ; this may lead to an underestimationof the tunnel effect.The true barrier profiles may have an intermediate course1 Christov, (a) Dokl. Akad. Nauk., S.S.S.R., 1960, 136, 663 ; (b) Ann. Physik, 1963, 12, 20.ZChristov, Ann. Physik, 1965, 15, 87.3 Christov, Proc. 1st Austral. Con$ Electrochemistry, 1963, p. 723.4Christov, 2. Elektrochem., 1960, 64, 840.5 Jonston and Rapp, J. Amer. Chem. Soc., 1961, 83, 162 GENERAL DISCUSSIONbetween the Eckart and parabolic barrier (with the same curvature at the top),which can be expressed better by means of the generalized Eckhart potential.4bs 5It is therefore possible that the difference between the mean permeabilities ofthe one-dimensional and two-dimensional barriers are not very significant, at leastfor moderate tunnelling.Physically one can expect some cancellation between theeffects of the side-wall repulsion and the more favourable situation, for tunnellingalong the profiles, which are longer than at the saddle point; the tunnel correctionfactor r = V/V' will be larger for these profiles than at the saddle-point, althoughthe barrier height is bigger (and therefore the mean permeability smaller).Finally, the writer agrees with Mr. Bell that relaxation in the water moleculeorientation may contribute to the activation energy of some proton-transfer reactionsin solution.Although it seems that new measurements of hydrogen overpotential 1do not confirm the conclusion that the dielectric relaxation time for water is greaterthan the time of a proton-transfer in the electrode double-layer.Dr. E. F. Caldin (University of Leeds) (partly communicated): I agree withProf. Hills and Dr. Jones that curvature of the Arrhenius plot is not by itself con-clusive evidence for tunnelling ; the evidence is strengthened, however, if (as in ourwork) alternative reasons for the curvature can be eliminated. Isotope effectsshould be investigated wherever possible ; this cannot be done, without changingthe solvent, for the reaction to which our experimental results refer, but work isin progress on the reaction (10) in table 3 of our paper.The possibility of a lag in the reorientation of solvent molecules, mentionedby Mr. Bell, should be investigated further ; it would affect both the isotope effectand the Arrhenius plot.It does not, however, provide a complete explanation ofthe phenomena attributed to tunnelling. In a given solvent the effect of such a lagshould presumably increase (and so the apparent barrier-width should decrease)with increase of - AS*, which reflects the change of solvation on forming the transi-tion state; this does not agree with the data on reactions (7) and (8) in our table 3,which have values of - AS* of 16 and 9 cal deg.-1 mole-1, respectively, and apparentbarrier-widths of 1 -66 and 1 -46 A, respectively.Prof. Hills assumed that AUZ is a simpler quantity than AH*, but for chemicalreactions it does not appear to be so, either in terms of thermodynamic expressions(cf.Dr. Kohnstam's contribution to the discussion of Prof. Hills' paper) or in termsof the physical model. When the transition complex is formed, there is a changein the orientation of solvent molecules, and consequently a change in volume, inthe neighbourhood of the transition complex. Under constant-pressure conditions,the structure and volume of the bulk of the solvent is unaltered, so there is an overallchange in the volume of the solution. Under constant-volume conditions, thelocal volume change around the transition complex must be compensated by anequal and opposite volume change in the bulk solvent, effected by a change in theexternal pressure. Thus, the structure of the bulk solvent is altered, as well as thatnear the transition complex, and the two effects may be difficult to disentangle;AU: will not be determined by the reaction alone.The interpretation would prob-ably be simplest if AV* were kept constant, and this is nearly achieved in constant-pressure experiments at 1 atm.In reply to Prof. Conway, the barrier-widths quoted in our paper are certainlylarger than the distance that a proton would move if it were initially hydrogen-bonded, but (as we point out) if the distance is calculated from the van der Waalsradii there is reasonable agreement. Hydrogen bonding would be expected to beweak in most of the systems reported. If we were to assume that a hydrogen-bonded1 Nurnberg, Barker and Bolzan, Report Jul-1 37-CA, 1963, Kernforschungsanlage Julich)GENERAL DISCUSSION 63complex is formed as a first step in the reaction, followed by proton-transfer withinthe hydrogen bond, we should have to take into account the contribution of the heatof formation of the hydrogen bond to the energy of activation.Dr.M. J. Henchman (Leeds University) said: While tbe concern of this meetingis with proton transfer processes in solution, these processes have been extensivelystudied in the gas phase. Dr. Ausloos has shown how radiation chemistry can bemade to reveal much concerning these processes in the gas phase. The otherinstrumental method uses the mass spectrometer 1 : results by Mr. Ogle and myselfusing this technique indicate some of the factors governing proton transfer reactionsin the gas phase.The reaction chamber is a metal box, traversed by an electron beam: the gasunder investigation flows in through a hole and positive ions are formed by electronimpact in the electron beam : a positively-charged repeller plate pushes these ionsout of the chamber and once they reach the exit slit, they are rapidly accelerated,sorted according to their mass-to-charge ratios by a magnetic field and counted.At low gas pressures, the mass spectrum is an aliquot of the ions formed in theelectron beam and these are the conditions for analytical mass spectrometry. Athigher gas pressures, there is a chance that the positive ion will collide with a moleculealong its track in the reaction chamber, in which case the ion recorded in the massspectrum will be the product of the chemical reaction occurring at that collision.Mass spectra are recorded as a function of increasing pressure when the relativeintensities of the reactant ions decrease and those of the product ions increase.In this way, ion-molecule reactions can be identified and their rate constants measured.We have, e.g., characterized two proton transfer reactions( 2 )CH2NHl + CH3NH2+CH3NH: + CH2=NHoccurring in methylamine and we obtain the following rate constants kl = 9.6 x 10111.mole-1 sec-1, and k?. = 3.6 x 1011 1. mole-1 sec-1. These very high rate constantsare characteristic of ionic reactions in the gas phase and are due to both the ion-permanent dipole and ion-induced dipole interactions. A calculated value for therate constant based on this model 2 gives k = 11.4 x 1011 1.mole-1 sec-1, agreeingwell with the experimental value for kl, and emphasizes here that the rate constantis primarily determined by simple electrostatic considerations. k2 is lower thanthe calculated prediction since here the activated complex may break down by analternative way, according to a hydride ion mechanism,CH2NH,f + CH3NH2+CH3NH2 + CH,NHl.What determines the relative contributions of these two breakdown modes remainsan intriguing question since neither reaction possesses an activation energy.Quite detailed information can be obtained about the mechanism of reaction(1) and the effect of translational velocity upon this.From electrostatic con-siderations, the collision consists of a positive ion striking a dipole and one couldsuggest that, at low incident ion velocities, the ion-dipole interaction will alignthe collision complex in the configuration of minimum potential energy, as shownin fig. 1. Chemically this must be of the form shown and the nitrogen hydrogens1 see Lampe, Franklin and FieId, Progr. Reaction Kinetics, ed. Porter (Pergamon Press Ltd.,2 Moran and Hamill, J. Chem. Physics, 1963,39, 1413.London, 1961), 1, 67, for method and tabulation of reactions and rate constants64 GENERAL DISCUSSIONwill be sterically better placed for transfer. One could predict that the probabilityfor transfer of the nitrogen hydrogen would be greater than that of the carbonhydrogen, i.e., that PN(H)/Pc(H)> 1. On the other hand, at high ion velocities,the ion-dipole interaction will be unimportant, the collision complex will have arandom configuration and both types of hydrogen will be transferred with equalprobability, i.e., PN(H)/Pc(H) +l .CONFIGURATION OF COLLISION COMPLEXCH3NHl+ CH3NH2 +CH3NH: + CHzNH2 { CH3NHLOW ION VELOCITIESIon-dipole interaction determines configurationCH3I n H-N-HHIGH ION VELOCITIESElectrostatic forces unimportant c3- random (+--I configurationFIG.FIG.1.-The configuration of the collision complex at low and high incident ion velocities.4v !3un %,x 2 2 a,00 60 120 I80N molecules/d x 10-13with N, the concentration of methylamine in the reaction chamber.2.-The variation of the functions P@)/Pc(H) and PN(H)/Pc@), as defined in the textGENERAL DISCUSSION 65Differentiation between the two kinds of hydrogen can be achieved using deuteriumlabelling and the proton transfer reaction has been studied for CD3NH2 and CH3NDz.The rate constant for overall proton transfer is in both cases similar to kl, the rateconstant for the reaction for CH3NH2, but in both cases the relative probabilityfor transferring the nitrogen hydrogen increases as the gas concentration in thereaction chamber increases (fig.2). At the lowest gas concentrations, there is littlereaction and so little depletion of the primary ion beam as it passes down the trackout of the chamber. The ions are being constantly accelerated by the repeller fieldand reactive events are occurring over a range of velocities.On the other hand,at the highest gas concentrations, most of the ions react before they leave the chamberand so many ions react at low velocities that few remain to react at high velocities.Increasing the gas concentration emphasizes the low-velocity events and thus theresults in fig. 2 provide powerful support for the model outlined in fig. 1. At lowvelocities, the ion-dipole interaction is aligning the collision complex to give theconfiguration of lowest potential energy but as the velocity increases, this alignmentbecomes less important. In the language of chemical kinetics, the reaction, whichpossesses no activation energy, can proceed through two possible transition states.At high temperatures there is no discrimination between these. At low temperatures,even though there is no activation energy to determine the path, the reaction is stillmade to proceed via the transition state of lower potential energy by the electro-static interaction.These results suggest that both the rate constants and the mechanisms of protontransfer processes in the gas phase are powerfully influenced by electrostatic forcesand this emphasizes the difference between these reactions occurring in the gas andthe liquid phase.Dr.R. A. Ross (University College, London) said: Dr. Ausloos discusses Meyerson’ssuggestion that C3H; ions formed in the mass spectrometer have a protonatedcyclopropane structure. Baldwin, Maccoll and Miller 1 also find that their measure-ments of appearance potentials of C3H; ions from propyl halides are consistent witha common ion being formed. Further, the appearance potentials of the C4Hgions formed from n-butyl, sec-butyl, isobutyl and t-butyl halides again suggest thata common species is formed. While the common ion in the propyl case could be thesecondary ion this is rather improbable in the butyl case and they suggest that asimilar species to that proposed by Meyerson for the propyl ion, with the protonreplaced by CH; , is consistent with the measured appearance potentials.Dr. P. Ausloos (National Bureau of Standards) (communicated): In reply toROSS, the results of recent radiolytic studies have demonstrated conclusively thatthe butyl ions formed in the decompositions of the parent ions of n-pentane andneopentane do not have identical structures. The butyl ion produced in the radio-lysis of n-pentane 1 readily takes part in a hydride-ion transfer reaction with n-pentane-& to form CH3CH2CHDCH3, thus demonstrating that the precursor ionacquires the sec-butyl structure prior to or during reaction. On the other hand,the butyl ion produced in the radiolysis of neopentanez does not react with n-pentane-d12 to form either n-butane or isobutane, demonstrating that this butylion has a different structure from that formed in the n-pentane radiolysis. Thefailure of this ion to undergo a hydride transfer reaction with n-pentane indicatesthat it probably has a t-butyl structure, since reaction between a t-butyl ion andn-pentane is endothermic 3 and thus would be expected to compete efficiently with1 Baldwin, Maccoll and Miller, paper to ASTM Mass Spectrometry Conference (Paris, 1964).2 Ausloos and Lias, J. Chem. Physics, 1964,41, 3962.3 Ausloos and Lias, J. Chern. Physics, in press.66 GENERAL DISCUSSIONneutralization. If the butyl ion in the two systems discussed above were initiallyformed with an identical structure such as the one referred to by Mr. Ross, onewould be faced with the unlikely situation that they rearrange upon reaction todifferent structures whose form is dependent on the nature of the precursor molecule.Finally, if the protonated cyclobutane referred to in our paper, or the butyl ion formedin the radiolysis of n-heptane rearranges to a structure such as the one mentionedby Mr. Ross, one would expect that, contrary to the results, the hydride ion transferreaction with deuterated higher hydrocarbon molecules would result in the formationof CH2DCH2CH2CH3 rather than CH3CHDCH2CH3

ISSN:0366-9033    

DOI:10.1039/DF9653900045

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Kinetics of Proton Transfer to Weak Aromatic Bases*BY B. C. CHALLIS AND F. A. LONGDept. of Chemistry, Cornell University, Ithaca, N.Y., U.S.A.Received 1 1 th January, 1965The rates of proton transfer to and from the weakly basic aromatic species azulene are sufficientlyslow that the rate of approach to a displaced equilibrium can be measured in a fast-flow apparatus.Independent measurements of the equilibrium protonation then permits calculations of the separaterate coefficients ky and kr for protonation and deprotonation respectively. For the ionization ofthe azulenium ion as an acid, AGO = -2-4 kcal ; ASo = - 11 cal/mole deg., and AH' = -5.8 kcal.At 7.3", over the acidity range 1.5-4-0 M HC104, kf = 1-52 hk26 and kr = llOh;P.68 where ho isthe Hammett acidity function.The Arrhenius parameters for the protonation reaction are similarto those for the acid catalyzed-tritium exchange of azulene-1-t in aqueous media and the aciditydependence of protonation is similar to the acidity dependence of the tritium exchange for sub-stituted azulenes, suggesting related mechanisms. Quantitative evidence that the exchange reactiondoes go via a conjugate acid intermediate, is provided by the fact that, after some necessary medium,isotope effect and statistical corrections are made, the specific rate for protonation of azulene isfound to agree within limits of error with a protonation rate derived from the tritium exchangedata assuming the 2-step A - sE2 mechanism for the latter.The exchange of hydrogens attached to an aromatic nucleus with the hydrogensof an aqueous solvent has been studied in recent years by a number of groups.1-*The reaction is catalyzed by acids but not by bases and exhibits general acid catalysis.This and a number of the other details, including the kinetic isotope effects and themagnitude of the Arrhenius parameters, strongly suggests that the reaction involvesan electrophilic attack of the solvated proton on a carbon atom of the aromaticsystem. This offers a particularly simple example of electrophilic acid substitution.There has been considerable discussion of the mechanism for the reaction. Thebulk of the evidence clearly supports a two-step A - S E ~ mechanism.H L H LIt is clear that the kinetic data themselves do not require that the reaction gothrough the comparatively stable conjugate acid as an intermediate. On the otherhand, such a mechanism, leading to a " two hump " diagram for the free energyof activation as function of reaction co-ordinate (fig.3), is very plausible and isentirely consistent with the similar mechanism proposal made earlier by Melanderfor more obvious electrophilic substitutions on aromatic ring systems.6One way to obtain evidence on the validity of this mechanism, and the subjectof this paper is to study directly the rate of proton transfer from acidic media to andfrom an aromatic ring system. That this direct proton transfer reaction shouldoccur, and at a measurably slow rate, is indicated by the free energy of activationdiagram of fig.3 which is approximately to scale, if one assumes that the reactiondoes go via the conjugate acid.* work supported by a grant from the Atomic Energy Commission.668 PROTON TRANSFER TO AROMATIC BASESThe rate of tritium exchange of azulene occurs at an easily measurable rate forcatalyst concentrations of around 0.001 M hydrogen ion; kex the specific rate co-efficient for the reaction between azulene-1-1 and H3O+, is 0.183 1. mole-1 sec-1at 25O.3 However, to measure the proton transfer directly one must go to sufficientlyacidic solutions so that the equilibrium is measurably displaced. The base strengthof azulene is such that it is half-protonated at about 2-2 M perchloric acid. Hencea displacement of equilibrium of a magnitude which permits easy spectroscopicanalysis involves working at acidities of around 2 M strong acid, at which pointthe reaction rates bewme large.The solution to this has been to use a fast flowapparatus of the Partridge-Roughton type, with a mixer and observation systemthat permits analysis within less than 10 msec after mixing.Symbolizing the neutral azulene molecule by AzH, the reaction concerned isfkrAZH + H + +AZH,+,where kf and k, are the first-order rate coefficients for protonation and deprotonationrespectively. The observed reaction is the first-order approach of a displacedsystem to a new equilibrium. The first-order rate coefficient kut for this processis linked to those above bywhere the subscript e denotes equilibrium concentrations. Hence a measurementof kist and of the indicator ratio I = [AzHg],/[AzH], permits determination of kfand kr.For later use, one can further define a second-order rate coefficient forprotonation, kbi = kf/[H+].EXPERIMENTALThe fast-flow apparatus was specifically fabricated to permit use of concentrated aqueousmedia. A 4-jet circular tangential mixer was employed; this permitted mixing to within1 msec. In a typical experiment one syringe contained nearly saturated azulene (about10-5 M) in 1 M perchloric acid. Another syringe of equal cross-section contained 5 Mperchloric acid. These solutions were mixed in equal volumes by a plunger moving at apredetermined rate. After mixing the solution flowed along a quartz observation tube ofknown diameter. Extent of reaction at a given point along the tube was determined bymeasuring the change in light absorption at 350A, a wavelength where azulenium ionabsorbs strongly and azulene almost negligibly2 Final acid concentration was determinedby titration of the mixed solution. Temperature was measured for the flowing solutionwith a thermocouple; temperature control was always to within f0.4".A few experi-ments were performed in which equilibrium was approached from the more acid side;the results agreed with those from studies of the above type.A typical signal record for a kinetic experiment is shown in fig. 1. The lines A and Brefer to the intensities of the initial perchloric acid and unreacted azulene solutions respec-tively ; line C refers to the combined solutions under flow and line D to the hal equilibrium.The calculation of the rate coefficient involves2.303 [AzH,C], - [AzH'] 2.303 D -(A + B)/2log10 D-C 9o = -loglo [AzH2f],- [AzHl], t klst = - twhere D, A, B and C are heights of the above lines from the base line and where t wascalculated from the flow rate and geometry.The normal procedure was to plot valuesof log10 D-(A+B)'2 D-c against i for a series of measurements at constant acidity but varyinB . C . CHALLIS AND F. A . LONG 69reaction times. The slope of the line gives 2.303/klSt. Reactions were accurately first-order up to at least 85 % reaction.The equilibrium ratio [AzHz],/[AzH], was determined in a Cary model 14 spectrometerwhich permitted temperature control to 0.1".Spectra were recorded at several wavelengthsbetween 2200 and 3600A and the indicator ratio was calculated from the results in theusual way. At each temperature data were recorded at a number of acidities. Therrno-dynamic dissociation constants KA& were determined by plotting values of log ([AzHz]/[AzH][Hf]) against perchloric acid concentration and extrapolating to zero acid con-centration .9Itime-,FIG. 1.-Diagram of the recorded signal for a single kinetic point.RESULTSTable 1 records the acidity constants KAH; as a function of temperature. Thesedata lead. for the ionization reaction, to AGO = -2.4 kcal; ASo = -11 cal/moledeg. ; AHo = - 5.8 kcal. As the small value of AGO indicates, the azulenium ionis of comparable stability to azulene itself.TABLE 1 .-THERMODYNAMIC DISSOCIATION CONSTANTS KA~HZAS FUNCTION OF TEMPERATUREtemp.O C KA& mole 1-15.7 106 f 615.5 69 f 625.0 56 f 638.6 36 f 6The results for an extended list of kinetic studies at 7.3" are given in table 2where the data are analyzed into values of kf and kr. Fig. 2 gives plots of log kfand log kr against the acidity function Ho. For the protonation reaction the datalead to log kf = 0.18 + 1.26 (-Ho). The deprotonation reaction, whose stoichio-metry does not require acid, also varies strongly with acidity, i.e., kr~Ch,0-68, in-dicating a very pronounced medium effect on the reaction70 PROTON TRANSFER TO AROMATIC BASBSTABLE 2.-vARIATION OF kist, kf AND k, WITH ACIDITY AT 7.3f0.4"CIHCI041M1.461 -481 -491.161-711.781 -922-052-092-502-572-582-692.852.993-003.283-313-653.91- HO0.5100.5270-5320.5920.6440.6800.7500,8050.8251.011 -401.w~1 -091-161 *221 a231.341.351.541.67[AzHiI *1-10.1390.1480.1500.1950.2450,2840.3850.5040-5431-281 *461 -491-862.603-513.576.2 16.4313.221.6kist sm-151.351.448-348-348.852-052.447.852.250.053.951.159.868.974.673.783-998.3139174kf sec-16.36.66.38.79-611.514.616.018.628.132-030-638-949.858.157.672.385.6129166* from equilibrium studies at 57°C.TABLE 3.-vARIATION OF kist, kf AND k, AS FUNCTION OFTEMPERATURE FOR HC104 = 1.93 Mkist sec-1 kf sec-1 temp.O C7.3 f0.4 0.398 52-4 14.912.7 f0.4 0.427 79.0 23.616.7 f0-2 0.450 108 33.519.9 f 0.2 0.470 138 44.1[AzH;l1-1kr sec-144.944.642-039-639.240.534.832.034221.921.920-518.919.115.516.111.612.49.87.7k, sec-137.555.474.593.9Table 3 gives the temperature dependence of the rate coefficients for reactionin the single medium 1-93 M perchloric acid. These data lead to the followingprotonation (sec-1) deprotonation (sec-1)AG', kcal 16.0 15.6AH*, kcal 15.1 12.4AS*, cal/mole deg. - 3.2 -11.0Arrhenius parameters. These data imply that the thermodynamic parameters forequilibrium in this solvent differ appreciably from the values for an infinitely diluteaqueous medium.Specifically for equiIibrium in 1-93 M perchloric acid, AG =-0.4 kcal; AH = -2-8 kcal; A S = -8.1 cal/mole deg.DISCUSSIONIf the acid-catalyzed tritium exchange of azulene-1-t goes by a two-step A - S E ~mechanism (see fig. 3), via the conjugate acid as an intermediate, then the first stepof the reaction should proceed at essentially the same rate as the direct protonationof azulene. Qualitatively, there is marked indication of similar behaviour. Forthe exchange reaction, with rate coefficients, in 1. mole-1 sec-1, the Arrhenius para-meters for reaction in dilute aqueous solution 3 are : AGTx = 19 kcal ; AS; B . C . CHALLIS AND F . A . LONG 71- 10.1 cal/mole deg. ; AH: = 16 kcal. These values are similar to those forprotonation in 1.93 M perchloric acid. A further point is that for azulenes withexchange rates slow enough to be measurable in concentrated acid solutions, first-order rate coefficients for exchange vary with acidity much like the protonationreaction does.Thus, for I-CN-azulene-3-t,3 kexcchk2. It is therefore of interestto see if there is quantitative agreement between the rates for the protonation re-action and for the presumed first step of the exchange reaction.If mechanism (I) is correct for the exchange reaction, then with tritium used attracer level so that k-2 can be ignored,kzx = kT/(l+ k'! JkT),where the superscript T means only that the rate coefficient is for mechanism I.The comparison which is desired is between kbi for protonation and k? in the samemedium for the hypothetical exchange of hydrogen atoms.Comparison for thesame medium is important because of the strong medium effects indicated by fig. 2.5 0-5 1.0 1.5(-H0)?5'FIG. 2.-Plot of loglo k, and loglo kf against ( -Ho), 7.3".For the exchange reaction we shall first assume that kT = ky, i.e., that thesecondary isotope effects on the rate of protonation are negligible.11, 12 The problemthen is to calculate kl from available information on the exchange rates. For-tunately this can be done if data on both kinetic and deuterium solvent isotope effectsare available. This is true for the azulene exchange reaction.2, 3 Using the availabledata and the procedure of Kresge 1 we obtain 72 PROTON TRANSFER TO AROMATIC BASESThe latter ratio then leads to a predicted kinetic tritium isotope effect on k2 of kF/kT= 12 which in turn means a predicted A(AG*) between the two humps of fig.3of 1500 cal mole-1 at 25". Since by fig. 3 it is evident that kT, /kT is, except forsecondary isotope effects, the same as ky/kT, we may write,kY kYkT, = 1 +(ky/kT) = 1 + exp (+ 1500IRT)'The Schulze and Long tritium exchange data 3 interpolated to 7.3" give k', =0.032 1. mole-1 sec-1 for reaction in an aqueous solution containing 0.1 M electro-lyte. Application of the above equation then gives k? = 0-5 1. mole-1 sec-1 asthe rate coefficient for protonation at (a single) 1-position of azulene at thistemperature.reaction co-ordinateFIG. 3.-Free energy profile for two-stage A-S& mechanism, 25".Two corrections need be made on the data for kbi before comparing it with k?.The simpler is a statistical correction of 2 to take account of the fact that the pro-tonation reaction has two equivalent base sites to attack (the 1- and 3-positions),lowhereas k y is for attack at only one of these.The more difficult correction is formedium effects. The most reasonable procedure is to extrapolate the data forkbj and k,. to give the limiting rate coefficients k& and k," for infinitely dilute aqueoussolutions. The appropriate type of extrapolation is indicated by the following.The coefficient k& is defined by the equations-d[AzH]/dt = kf[AzH] = kbj[AzH][H+] = k,",[AzH][H+]f,,~f,+/f*,where f k ~ f , f ~ f and f* are activity coefficients of azulene, hydrogen ion andtransition state respectively.If, as usual, the activity coefficients are referred to a value of unity at infinite dilution,Rearrangement and taking logs leads tolog k& = log kf-log [H+]-log (j&JH+/f*).log k; = lim [log kf-log [H']].[H+]+OThe form of the activity coefficient ratio involved here is similar to that enteringin the similar extrapolation for ionization equilibrium data.Since a linear extra-polation has been found to be valid in the equilibrium case,9 it should be valid foB. C . CHALLIS AND F. A . LONG 73the kinetic case also. A similar development for the deprotonation reaction leadsto the relationshiplog k," = lim [log k,.][H+]-OThe appropriate plots for the kinetic data are shown in fig. 4 ; both are linearwithin limits of experimental accuracy.Least-squares treatment leads to kii =1-17 1. mole-1 sec-1 and k," = 135 sec-1 at 7.3" and infinite dilution. A good testof the extrapolation is to compare the calculated equilibrium constant k,"/k,"i = 115with the previously determined value of K ~ H ; = 100 at the same temperature.The agreement is as good as can be expected.FIG. 4.-Linear extrapolation for protonation and deprotonation reactions at 7.3".The plot of fig. 4 can also be used to give a value of kbi at the electrolyte con-centration of 0.1 M which is the one of interest for comparison with kf: from theexchange data. From fig. 4, the value of kbi at this concentration is 1-28 1. mole-1sec-1. Hence the desired comparison between protonation data and ky fromexchange data is between kbi/2 = 0.64 1. mole-1 sec-1 and k y = 0.5 1. mole-1 sec-1.These agree to well within the experimental error of the assumptions and extrapola-tions involved. We conclude that results on the protonation reaction provide semi-quantitative evidence for the validity of the proposed 2-step A - S E ~ mechanismfor the exchange.1 Kresge and Chiang, (a) J. Amer. Chem. SOC., 1959, 81, 5509 ; (b) Proc. Chem. SOC., 1961,81;2 Colapietro and Long, Chem. and Ind., 1960, 1056.3 Schulze and Long, J. Amer. Chem. SOC., 1964, 86,322,327,331.4 Gold and Satchell, J. Chem. Soc., 1956,2743 ; 1960,2461 ; see also ref. (5).(c) J. Amer. Chem. SOC., 1961, 83, 2877; (d) J. Amer. Chem. SOC., 1962, 84, 397674 PROTON TRANSFER TO AROMATIC BASES5 Gold, Friedel Crafts and Related Reactions (ed. Olah) (Interscience, New York, 1964), vol. 11,p. 1253.6 (a) Melander, Isotope Eflects on Reaction Rafes (Ronald Press, New York, 1959) ; (6) Melanderand Olsson, Acta Chem. Scand., 1956, 10, 879.7 Eaborn and Taylor, J. Chem. Soc., 1960,3301.8 Thomas and Long, J. Amer. Chem. Soc., (a) 1964, 66, 4770 ; (b) J . Physic. Chem., 1964, 29,9 Paul and Long, Chem. Rev., 1957,57, I .10 Heilbronner and Simonetta, Helv. chim. Actu, 1952, 35, 1049.11 Streitweiser, Jagow, Fahey and Suzuki, J. dmer. Chem. SOC., 1955, 80, 2326.12 Olsson, Arkiv. Kerni, 1960, 16,487.341 1

ISSN:0366-9033    

DOI:10.1039/DF9653900067

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

General Acid Catalysis in Moderately Concentrated AqueousSulphuric AcidBY A. J. KRESGE, L. E. HAKKA, S. MYLONAKIS AND Y. SATODept. of Chemistry, Illinois Institute of Technology, Chicago, Illinois 6061 6Received 12th January, 1965Rates of aromatic hydrogen exchange in 1,3-dimethoxybenzene are found to be 23-3 timesgreater in moderately concentrated sulphuric acid than in moderately concentrated perchloric acidwhen comparison is made on the basis of ho values. Somewhat smaller differences in the oppositedirection are found for the equilibrium protonation of azulene in these same solutions. Com-bination of these kinetic and equilibrium data provides a measure of the rate of aromatic hydrogenexchange in sulphuric acid through acidic species other than H3O+, and a comparison of this excessrate with concentrations of various solution species indicates that the additional reaction occursby proton transfer from HSOZ and H2S04.This establishes general acid catalysis for aromatichydrogen exchange in moderately concentrated sulphuric acid and shows that there is no fundamentaldifference between proton transfer from strong acids and proton transfer from weak acids.A considerable body of evidence indicates that acid-catalyzed aromatic hydrogenexchange occurs through simple protonation and deprotonation of the aromaticsubstrate.1-* The reaction has a single intermediate, the cationic species HArH+,in which the exchanging hydrogens occupy equivalent positions :H’Ar + HA%H’ArH+ + A-+HAr + H’AA reaction of this type should show general acid catalys-h, and general acid catalysisis observed for aromatic hydrogen exchange in dilute aqueous solutions of weakacids.19 2 Under these conditions, the rate of exchange is the sum of contributionsfrom all acidic species present in the reaction mixture, and these contributions arecorrelated well by the Bronsted relation.However, in concentrated solutions of strong mineral acids, rates of aromatichydrogen exchange seem to bear no simple relationship to the concentrations of theacidic species, and general acid catalysis seems to be absent3 10 The rate is governedinstead by an acidity function which is a measure of the thermodynamic acidity of thesolution. This difference between the behaviour in dilute solution and the behaviourin concentrated solution is unexpected,los 11 and it has led to the suggestion that afundamental difference exists between proton transfer from weak acids and protontransfer from strong acids.12 In order to determine whether such a distinction isnecessary, we have undertaken the examination of aromatic hydrogen exchange inmoderately concentrated solutions of strong acids.It is difficult to say just how general acid catalysis in concentrated solutionsshould be recognized.Proton transfer is an ionic reaction and is certain to shownon-ideal behaviour in solutions containing large amounts of dissociated electrolytes.In concentrated solutions of strong acids, therefore, quite large deviations from pro-portionality between rate and concentration can be expected.Any estimate of thesedeviations will necessarily be uncertain, for added to the normal problems encounteredin treating forces between stable ions at high concentrations is the considerable diffi-culty of handling transition states. It would seem, therefore, that an analysis along75(176 CATALYSIS I N SULPHURIC ACIDthe lines usually carried out for dilute solutions will not be diagnostic. The conclusionwhich has been reached from one analysis of this type, i.e., that general acid catalysisis absent in aromatic hydrogen exchange,g should perhaps be taken with somereservation.A method of detecting general acid catalysis in concentrated solution which seemsmore reliable because it corrects for non-ideal behaviour in an empirical way is thecomparison of reaction rates in solutions of different kinds of acids.In a moderatelyconcentrated aqueous solution of a strong monobasic acid such as perchloric or nitric,the chief acidic species is the hydronium ion, H 3 0 f . In a moderately concentratedsolution of a polybasic acid such as sulphuric or phosphoric, there are appreciableamounts of other acidic species as well as H30+. If these other acids contributeto the rate of reaction, i.e., if the process is catalyzed by general acids, then reactionwill be faster in the polybasic acids than in the monobasic acids. This kind ofcomparison of rates has already been made for a number of reactions which showgeneral acid catalysis in dilute solution, and faster rates in polybasic than in monobasicacids have usually been found.Thus, sulphuric acid is a more effective catalyst thanperchloric acid in the iodination of acetophenone.13 More recently, this test hasbeen applied to several reactions which involve rate-determining proton transferfrom catalyst to substrate and which, therefore, are more closely analagous to aro-matic hydrogen exchange. In the hydrolysis of aryl boronic acids, the effectivenessof strong acid catalysts increases in the order HC104, H2SO4, H3P04;14 similarily,H3P04 and H2S04 are better catalysts than HCl, HN03, and HC104 for the hydrationof mesityl oxide and crotona1dehyde;ls and H2S04 is more effective than HC104in the dehydration of Ph.CH(OH).CH2.COMe and p-N02.C6H4.CH(OH).CH2.COMe,l6 in the formation of 1 ,Zcyclohexadione from its en0l,l7 and in the hydrolysisof vinyl mercuric iodide.18This test has also been applied to aromatic hydrogen exchange, but the resultsare equivocal. With p-chlorophenol, exchange was more rapid in H3PQ4 than inH2SO4, but with p-cresol, exchange was slower in H2SQ4 than in HCl.19 Thisinconsistency, however, may be more apparent than real, for these comparisons weremade on the basis of old values of the acidity function ho * before it was generallyrecognized that acidity function values are strongly dependent on the structure of theindicator bases used to measure them.Recently, ho has been redetermined insulphuric 20 and perchloric 21 acids by spectrophotometric methods with the sameset of primary amine indicators.These new values of ho are significantly differentfrom older values 22 over considerable regions of acidity. They make available,for the first time, a self-consistent and precise acidity function for two acids, and,in so doing, provide a most suitable basis on which to make fine comparisons of acid-base phenomena in two different concentrated acids. As part of our attempt todetermine whether aromatic hydrogen exchange is subject to general acid catalysisin concentrated aqueous acids, we therefore have measured rates of exchange of arepresentative aromatic substrate in perchloric and sulphuric acids.There is still another way in which general acid catalysis in concentrated solutionsmight be detected for aromatic hydrogen exchange.This method is based on thefact that the intermediate in exchange, HArH+, can be detected easily and its con-centration measured accurately when it is present in sufficiently large amounts in* It would seem that an acidity function is the most suitable basis on which to make a com-parison of rates which is designed to detect any reaction through an acidic species other than H3O+.An acidity function contains within it the effect of non-ideal behaviour on equilibrium proton transferfrom H3O+ ; it comes closer, therefore, to making the proper correction for kinetic proton transferfrom H3O+ than other measures of acidity such as CH~O+ or wt. 74 acid which are purely stoichio-metric quantitiesA . J . KRESGE, L. E . HAKKA, S . MYLONAKIS AND Y .SAT0 77acidic solutions.39 49 23 In any acidic solution, the concentration of HArH+ willbe governed by the thermodynamic acidity of that solution toward an aromatic base.If the solution is aqueous, the thermodynamic acidity is determined by the advitiesof the H30+-H20 pair, and the concentration of HArHf is controlled by its rate offormation from HAr and H30+ and its rate of destruction by reaction with H20.The presence of other acids and their conjugate bases may raise the rate of inter-conversion of HAr and HArHf, but it should not alter the concentration of HArHffrom the value determined by H3O+ and H20. This phenomenon can be used todetect aromatic hydrogen exchange by any acid other than H30f, and it can alsoprovide a measure of the additional rate.To do this, rates of aromatic hydrogenexchange and concentrations of MArHf must be measured in two different acids,one monobasic and one polybasic. If the rate of exchange in the polybasic acid isgreater than the rate of exchange in the monobasic acid at the same concentrationof HArH+, then reaction must be occurring through some species other than H@+in the polybasic acid. The extent to which the rate in the polybasic acid exceeds thatin the monobasic acid is a measure of the rate of reaction by other acids.This four-fold comparison of rates and equilibria of aromatic protonation shouldmost properly be performed on a single aromatic substrate. But in the region ofacidity where measureable amounts of HArH+ are formed from a given substrate,rates of aromatic hydrogen exchange for that substrate are very fast. We thereforeused two substrates of somewhat different reactivity and basicity to make this com-parison.For the kinetic experiments, we chose 1,3-dimethoxybenzene labelled withtritium at the 4-position and measured aromatic hydrogen exchange by the rate ofloss of tritium from this substance. This substrate is the least reactive aromaticfor which general acid catalysis has been detected in dilute solutions of weak acidsand for which the parameter a in the Bronsted relation is known ;24 its rate of exchangecan be measured by conventional methods in sulphuric or perchloric acid up to30-40 wt. % acid. None of the simple benzene derivatives is sufficiently basic to giveaccurately measurable amounts of HArHf at these acidities, but azulene is a stronglybasic aromatic whose kinetic and equilibrium protonations are closely similar to thoseof the methoxybenzenes.294 We therefore used azulene for the equilibriumexperiments.EXPERIMENTALMATERTALS.-~ ,3-Dimethoxybenzene-4-t was prepared by treating a solution of 1,3-dimethoxyphenyl magnesium bromide in tetrahydrofuran with tritiated water.Theproduct was isolated in the usual way and was purified by fractional distillation at atmos-pheric pressure. All other materials used were the best available commercial grades;azulene was purified by vacuum sublimation and recrystallization.KINETICSRate measurements were made at 24-69f0.02"C by a method essentially the same asthat already described.1 The wholly aqueous reaction mixtures were prepared by mixingone part of a solution approximately 10-3 M in aromatic with 75 or 100 parts of an acidsolution.Usually, ten 5-ml kinetic samples were taken at time intervals which providedcounting rates ranging from 105 c.p.m. down to a few times background (30 c.g.m.) ; thisfurnished kinetic data over at least ten half-lives. Acidimetric determinations were carriedout directly on samples of reaction mixture.The kinetic data are presented in table 1. They obey the following relationships : inHC104, loglo k = -(3.273 &-0-017)- (1.139 Ilt0.009)Ho ; in H2S04, loglo k = - (2-858 f0.013)- (1.175 f0.023)Ho. (Errors are standard deviations.78 CATALYSIS IN SULPHURlC ACIDTABLE 1.-bTE OF LOSS OF TRITIUM FROM 1,3-DIMETHOXYBENZENE AT 25°Cacid wt.% - H$ l@kl (min-1)HClO4 9-94 0.35 0.134Y Y 18.82 0.91 0-582Y Y 26-82 1-39 2.0 1Y Y 33-95 1-86 7.08¶ Y 37-10 2.1 1 13.540.21 2.37 26-7Hzi04 5-21 - 0.08 0.0983Y Y 10.42 0.34 0.33611 15.62 0.70 0.960¶ Y 23-15 1 *23 3.99Y Y 28.35 1.61 10.5¶ S 20.83 1 -07 2.49values for HC104 from ref. (21) and for &SO4 froin ref. (20).TAELE 2.-EQUILIBRIUM PROTONATION OF AZULENE AT 25°Cacid~ ~ 1 0 ~Y Y¶ YY Y9Ywt. % "8-7511-4213.7816-1418.762 1 *2425-9927.5629.8830.827.008.7911.3113-6715.8 118-3320-2622.8825-1927.5429-273 1.50-If$0.270.440.590-740.901 -051.341-441-581 -640-088.220.400.580-720.901 -831.221.381.541-661 -82a solutions contain 0.5 7; methanol by volume.0 values for HCI04 from ref.(21) and for H2S04 from ref. (20).EQUILIBRIAloglo CHA~H+ KHA~ - 1-29- 1.01- 0.75- 0.49- 0.20 + 0.080.690 821.091.21- 1.58 - 1.41- 1.12- 0.89- 0.67- 0.40 - 0.20 + 0.080.360.620.831.13Absorbance measurements were made at 250+ 0-5"C and 2790 A using a Beck.mannDK-2 spectrophotometer. Replicate determinations and independent measurement of thetemperature coefficient of the extinction coefficient at this wavelength showed that absorb-ances were accurate to at least 0.01 unit. Substrate concentrations (2 x 10-5 M) and opticalpaths (1 cm) were chosen so as to make full use of the 0-1 absorbance scale. All solutionscontained the same amount of substrate and were in an aqueous solvent which contained0.5 % methanol by volume.Acidimetric determinations were performed directly on thesA . J . KRESGE, L . E . HAKKA, S . MYLONAKIS AND Y . S A T 0 79solutions. It was found that azulene is slightly unstable under the conditions of theseexperiments. The greatest amount of decomposition occurred in the region of half-protonation, but even here absorbance decreased at a rate less than 1 % per hr., and thisinstability had no effect on the accuracy of these measurements.The equilibrium data are presented in table 2. In both acids, the relationship betweenloglo (CHArH+/CHAr) and HO is not quite linear. In HC104, the slope of a 1 0 g ~ O ( c ~ H f / c ~ )against HO plot increases from 1.69 to 1.87; in H2SO4, from 1.25 to 1.53.Agreementbetween these data for HC104 and published values 4 is good when account is taken of theslightly different NO scales used.DISCUSSIONAromatic hydrogen exchange, as measured by the rate of loss of tritium from1,3-dimethoxybenzene-4-t, is significantly more rapid in sulphuric acid than in per-chloric acid. Fig. 1 shows that the ratio of the rates in the two acids increasessomewhat with acidity; at HO = -0.3, ~ H , s o , / ~ H C ~ O , = 2.5 and at Ho = - 1.7,kH,sO,/kHc1o4 = 3.0. These ratios are similar to those observed in comparisonsmade on other general acid catalyzed reactions, and the interpretation made in theother cases would seem to be allowed here : in sulphuric acid, there is another reactionin addition to proton transfer from H30f.FIG.1 .-Rates of loss of tritium from 1,3-diniethoxybenzene-4-t at 25°C.Behaviour in perchloric acid different from that in sulphuric acid is also foundfor the equilibrium protonation of azulene. Here, however, the difference is in adirection opposite to that found for the rates of aromatic hydrogen exchange of1,3-dimethoxybenzene, and it is somewhat smaller in magnitude than the ratedifference. Fig. 2 shows that the equilibrium difference also increases with acidity ;at HO = -0.2, JH,SO,/~HC~O, ( I = CHA~H+/CHA~) is almost unity and at Ho = - 1-780 CATALYSIS IN SULPHURIC ACIDIH~soJIHc~o, = 0.34. Because the rate and equilibrium differences are in oppositedirections, they will reinforce one another when rates are compared at the same-L I I I0 I 2- HoFIG.2.-Equilibrium protonation of azulene at 25°C.FIG. 3.-Rates of loss of tritium from 1,3-dimethoxybenzene-4-t compared toequilibrium protonation of azulene.concentration of protonated aromatic, HArH+, in the two acids. Fig. 3 shows thatrates of interconversion of HAr and HArH+ are as much as 5.5 times greater in sulphuriA. J . KRESGE, L. E. HAKKA, S . MYLONAKIS AND Y . S A T 0 81acid than in perchloric acid. The two lines of fig. 3, moreover, are still divergingat the highest acidities employed in the investigation, and the rate difference is likelyto be still greater at higher acidities.Rate constants for the additional reaction in sulphuric acid can be obtained bysubtracting from the observed rates in this acid interpolated values of observed ratesin perchloric acid at the same value of C H ~ H + / C H ~ .These “ excess” rates, presentedin table 3, are seen to increase rapidly with the stoichiometric concentration of sul-phuric acid. The composition of sulphuric acid solutions in this concentrationrange is known from Raman measurements,2~ and values of the concentrations ofH@+ and HSO, and the concentration equilibrium constant, (Kc)~soa, interpolatedfrom the Raman values are also given in table 3. Both the concentration of HSO,and ( K c ) ~ s 0 , also increase rapidly with stoichiometric sulphuric acid concentrationin the region of the kinetic measurements, and both quantities have not yet reachedtheir maximum values at the highest acidity employed for the kinetics.In this sense,the excess rate and the two quantities which would be expected to govern the magnitudeof the exchange reaction through HSO,, CHSO~ and ( K c ) ~ s o i , have parallel behaviour.TABLE 3.-ANALYSIS OF RATE DATA102kHC104b 102kexcess ‘HSO+ ‘HSO; (Kc)HSOa cHSO-[(K~)HSOi] %0.54 0-0521 OW6 1 0.70 0-37 0.28 0.191.13 0.118 0.21 8 1.5 0.77 0.59 0.591.76 0.260 0.700 2.3 1.2 1.00 1.2243 0.582 1.91 3.2 1.7 1.5 2.02.74 0.840 3.15 3.6 1.9 1.7 2.53.48 2.05 8.4 4.6 2.4 2 3 3.6= stoichiometric molar concentration of sulphuric acid.b interpolated values of observed rates in HC104 at equivalent values of CA~H+/CW.d molar concentrations and concentration equilibrium constants interpolated from values pro-(kH~S04) -(kHC104)-vided in ref. (25).If, however, all of the excess rate were the result of exchange through HSO,,then, since the value of Bronsted’s a for aromatic hydrogen exchange in 1,3-dimethoxybenzene is 0.5,24 there should be a close proportionality between the excess rate andthe product cHSO~[(&)HSO,-]~.A comparison of columns 3 and 7 of table 3 shows thatthis is not the case : the excess rate increases by a factor of nearly 200 over the rangeof acidity covered by the measurements while the increase in cHSO&&)HSO$ isten times smaller. This difference is most probably not the result of a rate-acceleratingsalt effect on the bisulphate-catalyzed reaction, for any salt effect on the rate shouldalso affect (K&SO, and will be included in the product CHSO~[(K~)HSO~ 3.It seemsmore reasonable to ascribe this difference to incursion of another reaction, exchangethrough molecular sulphuric acid. In the Raman investigations, none of the speciesH2SO4 could be detected below a stoichiometric sulphuric acid concentration of 14 M.The Raman line assigned to H2S04, however, is closely flanked by HSO, and SO,:!lines ; these are especially strong at low stoichiometric sulphuric acid concentrationsand can completely overshadow a weak H2S0.4 line. The Raman work, therefore,cannot be said to exclude the possibility that small concentrations of H2S04 arepresent in the sulphuric acid solutions of these rate measurements. Sinceis a stronger acid than either H3O+ or HSO,, its catalytic coefficient will be greaterthan that of the other acids present in these solutions, and quite small amounts ofH2S04 will make significant contributions to the excess rate82 CATALYSIS IN SULPHURIC ACIDSince the excess rate can be the sum of contributions from reactions throughHSO, and H2S04, it is not possible to isolate the HSO- reaction and obtain a numericalvalue for the catalytic coefficient of bisulphate ion.It seems likely, however, thatthe &SO4 contribution to the rate will be small at the lowest acidity employed(C = 0.54 M) and that the excess rate here is largely the rate constant for reactionby HSO,. On this assumption, the catalytic coefficients of HSO, and H30+ are ofsimilar magnitude, for, at C = 0.54 M, kHC10,xkexcess and CH~O+% CHSO,-.Thismight seem to be inconsistent with the acid strengths of these two species : at C = 0.54,(KC)~soT = 0.28 whereas the acidity constant of H30+ is usually taken to be 55.5.But the latter is not a true acidity constant ; the ionization constant of H30+ cannotbe measured in aqueous solution, and 55.5 is an assigned value based purely on aformalism. The fact that the catalytic effect of the hydronium ion is usually foundto be lower than the value calculated on the basis of kH30+ = 55.5 by at least an orderof magnitude 26 indicates that this is very probably not the proper value on whichto base rate comparisons.In moderately concentrated aqueous sulphuric acid, then, proton transfer fromHSO, and H2S04 occurs at a rate which is comparable to the rate of proton transferfrom H3O+.It seems, moreover, that as the stoichiometric concentration of sulphuricacid increases, the rates of the reactions through HSO, and H2SO4 increase morerapidly than the rate of reaction through H3O+. But this trend cannot continue,for as the system becomes depleted in water, proton transfer from H3Of must beginto catch up with the other reactions. Complete proton transfer from H30+ liberatesone covalently bound water molecule and several solvating water molecules as well.Proton transfer from MSO, and H2S04, on the other hand, consumes water: theconjugate bases of these acids are ions of greater charge than the acids themselvesand so will require more water for solvation.The same differences will be present,though to a smaller extent, in incomplete proton transfer, i.e., in kinetic protontransfer, from these acids, and a shortage of water, therefore, will favour reactionthrough H3O+ over reaction through HSO, and H2SO4.The former apparent absence of general acid catalysis for aromatic hydrogenexchange in concentrated solutions of strong acids led to the suggestion that there is afundamental difference between proton transfer from weak acids and proton transferfrom strong acids.12 This proposal was based on the observation that the twokinds of acids, the weak acids with which general acid catalysis was observed and thestrong acids for which it was presumably absent, differ markedly in the transferabilityof their protons.The rate of proton transfer between weak acids and water is generallyseveral orders of magnitude smaller than the rate of proton transfer between strongacids and water27 In this investigation, however, we have shown that aromatichydrogen exchange is subject to general acid catalysis by the strong acids HSO, andH2S04. Since the rate of proton transfer between HSO, and water is very rapid,28this observation vitiates the premise on which the above proposal was based andshows that there is no fundamental difference between proton transfer from weakacids and proton transfer from strong acids.This work was supported by grants from the United States Atomic EnergyCommission and the Petroleum Research Fund of the American Chemical Society.1 Kresge and Chiang, J.Amer. Chem. Soc., 1959, 81, 5509 ; 1961,83,2877.2 Colapietro and Long, Chem. and Ind., 1960, 1056. Challis and Long, J. Amer. Chem. SOC..1963, 85, 2524. Schulze and Long, J . -4mer. Chem. Soc., 1964, 86, 331. Thomas and Long,J . Amer. Chem. SOC., 1964, 86, 4770A . J . KRESGE, L. E . HAKKA, S. MYLONAKIS AND Y. SAT0 833 Kresge and Chiang, Proc. Chem. Soc., 1961, 81. Kresge, Barry, Charles and Chiang, J.4 Long and Schulze, J. Amer. Clzem. SOC., 1961, 83, 3340 ; 1964, 86, 322, 327.5Kresge and Chiang, J. Amer. Chem. SOC., 1962, 84, 3976. Kresge, Pure Appl. Chem., 1964,8,243.6 Gold, Lambert and Satchell, Chem. and Ind., 1959, 1312; J. Chem. Soc., 1960, 2461. Battsand Gold, J. Clzem' SOC., 1964, 4284.7 Eaborn and Taylor, J. Chem. SOC., 1960, 3301.8 Melander, Arkiu. Kemi, 1961,17,291 ; 1961, 18, 195.9 Gold and Satchell, J . Chenz. Soc., 1955, 3609.Amer. Cheni. SOC., 1962, 84, 4343.10 Gold in Olah, ed., Friedel Crafts and Related Reactions (Interscience Publishers, New York,11 Melander and Myhre, Arkiu. Kemi, 1959, 13, 507.12 Gold, Proc. Chem. SOC., 1961, 453.13 Zucker and Hammett, J. Amer. Chem. SOC., 1939, 61, 2791.14 Kuivila and Nahabedian, J. Amer. Chem. SOC., 1961, 83, 2159.15 Bell, Preston and Whitney, J. Chem. SOC., 1962, 1167.16 Noyce and Reed, J. Amer. Chem. SOC., 1958, 80, 5539.17 Long and Bakule, J. Amer. Chem. SOC., 1963, 85, 2313.18 Kreevoy and Kretchner, J, Amer. Chem. SOC., 1964, 86, 2435.19 Gold and Satchell, J. Chem. SOC., 1955, 2622.20 Jorgenson and Hartter, J. Amer. Chem. SOC., 1963, 85, 878.21 Yates and Wai, J . Ainer. Chem. SOC., 1964, 86, 5408.22 Long and Paul, Chem. Rev., 1957, 57, 1.23 Schubert and Quacchia, J. Amer. Ckrem. SOC., 1962,84, 3778 ; 1963,85, 1278.24 Kresge and Sato, unpublished work.25 Young and Maranville in Hamer, ed., The Structrire of Electrolytic Solutions (John Wiley26 Bell, Acid-Base Catalysis (Oxford Univ. Press, London, 1941), p. 93.27 Caldin, Fast Reactions in Solution (Blackwell Sci. Publ., Oxford, 1964), p. 263.28 Eigen, Kurtze and Tamm, 2. Elektrochem., 1953, 57, 103.1964), vol. 11, chap. XXIX.and Sons, Inc., New York, 1959), chap. 4

ISSN:0366-9033    

DOI:10.1039/DF9653900075

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

Proton Transfer to OlefinsBY V. GOLD AND M. A. KESSICKDept. of Chemistry, King’s College, Strand, London, W.C.2.Received 3rd March, 1965The experimental evidence relating to deuterium fractionation between hydrogen ions and wateris outlined. The conclusions which can be drawn from these measurements about the structureof the hydrogen ion are indicated. In the light of this information the rate and product isotopeeffects on the hydrogen-ion-catalyzed hydration of isobutene, the rate-limiting first step of whichinvolves a proton transfer to olefin, are discussed. The results are adequately described by thetheory of isotope effects on slow proton transfer from H30f to substrate. It is shown that morecomplicated mechanisms, in particular indirect pro ton transfer from H3O+ via a water molecule(or direct transfer from the outer protons of the H904f cluster), are likewise compatible with the results,but offer no advantages in the context of isotope effects.Recent discussions of proton transfer processes from the hydrogen ion in solutionhave been in terms of the formula H90,+ for that ion, whereas discussions of deuteriumsolvent isotope effects on protolytic reactions and equilibria have most frequentlybeen based on the formula H3O+.We here examine one aspect of the questionwhether the two approaches are mutually exclusive. The discussion is based inthe main on recent experimental work on olefin hydration, a reaction which lendsitself particularly well to the detailed study of kinetic solvent-isotope effects. Theinterpretation of reaction velocities and equilibrium constants of hydrion (i.e.,proton, deuteron or triton) transfer reactions in H20 + D20 mixtures involves theconsideration of isotope fractionation effects, especially between hydrogen ions andwater.This topic is reviewed as an essential preliminary to the main theme.1. ISOTOPE FRACTIONATION BETWEEN AQUEOUS HYDROGEN IONS ANDWATERA number of different methods of measurement point to the conclusion that,in a solution of a strong mineral acid in an H2O+D20 mixture, the deuteriumisotope is not randomly distributed between hydrogen ions and water moleculesbut shows a preference for the latter. The interpretation of the results of someof these measurements requires an assumption concerning the formulation of thehydrogen ion which makes the measurements particularly relevant to our maintopic.All these results for hydrogen ions are, to varying extents, complicated bythe unavoidable presence of the complementary anions.The main methods that have been used are :A. GALVANIC CELLS WITHOUT TRANSFERENCE.-The comparison of the e.m.f.of pairs of cells employing isotopically different media, such as 1Pt(D2) 1 DCl in D20 I AgCl, AgPt(H2) I HCl in H20 I AgCl, Ag8V. GOLD AND M. A . KESSICK 85in conjunction with the equilibrium constant for isotope exchange between hydrogenions and water, leads to a constant L', defined by(H + (H, 0),) ( DT)2X+ ' (C1 -)2(D +(D20)x)2(H20)2"''(C1-) 2'L' =where a bar over a formula denotes solution in heavy water and absence of a bara solution in light water.The numerical value of L' is independent of x, i.e., ofthe chemical formula assigned to the hydrogen ion in solution. However, L'measures not only the exchange constant but contains a contribution arising fromthe free energy of transfer of the ions between the two waters. The experimentalvalue of L' (at 25') is ca. 18. Swain and Bader 2 consider that the free energy oftransfer of hydrogen ions can be neglected and suggested a method for calculatingthat of the chloride ion. On this basis it is possible to correct L' for the transfereffect and to calculate a value of L = 8.25,3 where[ H + ( H2 O),l [ D 01 2x+ '[ D + (D, O)x] [ €3, O] 2x+L =(On the assumption that the transfer term for hydrogen ions is negligible it is nolonger necessary to distinguish between the different phases.)B.GALVANIC CELLS WITH TRANSFERENCE.-The e.m.f. Of the Cells 4Pt(D2) I DCl in D20 i KCl (sat.) in H2O I Hg2C12, HgPt(H2) I HCl in H20 i KC1 (sat.) in H20 I Hg2C12, Hg2D2 + Hg2C12 (solid)+2Df (in D20) +2C1- (in H20) f2Hg2H2 + Hg2C12 (solid) +2H+ (in H20) + 2C1- (in H20) + 2Hgare related to the electrode reactionsandbut, in addition, some changes will occur at the liquid junctions. On the assumptionthat the associated diffusion potentials cancel when the difference between the twoe.m.f. is taken, Purlee 5 has calculated L, and considers the value L = 11.0 to carryan uncertainty of only k0.2 arising from the neglect of the diffusion potential.As has been suspected before, 29 3 there appears to be a logical flaw in this procedure,and it can be shown that the true value of L must be less than 11.(A discussionof this problem and estimates of the size of the correction will be presented elsewhere.)C. PROTON MAGNETIC msoNANcE.-In a system in which there is rapid exchangeof protons between all non-equivalent positions, the chemical shift is the con-centration-weighted mean of individual shifts 6 attributable to these positions.We assume that the introduction of a strong monobasic mineral acid HX into watercreates j sets of different positions, each set containing vj members. These " posi-tions " include any which are created by the structure-forming or structure-breakingeffects of the ions.If we express the acid concentration a as the stoichiometric atom fraction ofhydrogen nuclei added in the form HX, i.e., a = [HX]stoich./([HX]stoich.+2[HzO]),irrespective of the isotopic nature of HX and H20, and measure the position of theresonance signal A relative to the resonance frequency in pure water, we can write -A = azvjSj = 6a. (1.3)iThe implied proportionality between a and A will hold provided that the nature andnumber of positions available to the protons inj-sites is not altered by the additio86 PROTON TRANSFER TO OLEFINSof acid. This will be true only for low values of a, i.e., as long as there is no inter-action between the “ spheres of influence ” of solute particles. Eqn. (1.3) is there-fore a limiting expression (aj-0).If the hydrogen nuclei in the system are not all protons but contain an atomfraction n of deuterium the development of the general equation requires someadditional assumptions.We expect the main effect to be due to the non-randomdistribution of isotopes between water and the j-positions and that the most im-portant fractionation will involve the positions most different from water, i.e.,the hydrogen nuclei that are actually part of the structural core of the hydrogen ion.The treatment assumes that the actual number of positions (vi andj) available to thehydrogen nuclei does not depend on isotopic composition, on the ground that struc-tural differences are more likely to arise beyond the more strongly held first one ortwo layers of water molecules around an ion, and at such distances the fractionationeffect relative to water is probably unimportant.It is also assumed that the ajvalues are independent of isotopic composition, which implies, for example, thatthere are no secondary isotope effects by neighbouring deuterons on proton shifts.This is not likely to be seriously wrong since the position of the proton resonancein water differs from that in deuterium oxide containing a small fraction of protium(and where nearly all protium would be present as HOD molecules) by only ca.0.01 p.p.m.6If A’ is the corresponding chemical shift in an isotopically mixed solvent charac-terized by n and at the same acid concentration a, it can be shown thatl-n+mjjwhere $j is the fractionation factor for thejth group of hydrogen nuclei relative towater, i.e.,The chemical shift of the water resonance caused by the addition of acids is muchgreater than that due to any other electrolyte.7 It is therefore natural to associatethis large effect with the proton or protons contained in the hydrogen ion and toascribe the smaller shifts observed with salts, as well as the differences betweenvarious strong mineral acids, to less definite “ solvation ” effects.In first ap-proximation the sums contained in eqn. (1.3) and (1.4) are therefore replaced by asingle term, so that 8A-= l-n+n$j.K *This expression is independent of v], the number of hydrogen nuclei-all consideredto be equivalent-in a hydrogen ion. With these simplifications, 4j represents aunique distribution coefficient of deuterium between hydrogen ions and water,which is in general given the symbol 1.The seriousness of these approximationscannot as yet be assessed exactly, but the following are relevant considerations.The value of t$j (or I ) determined by eqn. (1.6) for perchloric acid is 0-69+0-02and the corresponding value found for hydrochloric acid * is 0.70+0.02. Thevalues of for the two acids differ but in any reasonable division of 8 among* The previously quoted 8 value of 0-68 f0.02 involved a minor calculation errorV . GOLD AND M . A . KESSICK 87anion and cation effects,7 the molar shift due to the hydrogen ion is much greaterthan that due to the chloride or perchlorate ions. From a consideration of thesize of these effects we estimate that the true fractionation effect due to hydrogenions alone is within 0.03 of the experimental value of 1.This experimental result is again independent of any assumptions about thenumber of hydrogen nuclei in the hydrogen ion that are concerned in the fraction-ation equilibrium.However, the value L, determined by e.m .f. measurements, canbe reconciled with the n.m.r. result for 4j only if a specific structural assumptionabout the hydrogen ion is made.8 For reasons of simplicity we explicitly consideronly formulae with different numbers of equivalent nuclei, but we do not therebyimply that other formulae are necessarily to be excluded. The formula of such ahydrogen ion is written as Hzs+10;, although the number of oxygen atoms perhydrogen ion does not strictly matter.In the limiting case as n-1, the only isotopicforms of this ion to be considered will be D2z+10j; and HD2zOf,, with [D2z+~0Z]9[HD2,0:] and correspondingly [D20] 9 [HOD], so that9 (1.8) +L-- 1/2(2x+ 1)the rule of the geometric mean 9 being assumed to apply to the isotopic forms of thehydrogen ion. Values of 4j predicted in this manner from L on the assumption ofdifferent numbers x are given below. Since L is, according to the e.m.f. measure-ments, likely to be in the range 8-10, the calculations are reproduced for these twovalues of L. In comparing the calculated values with the experimental value of# a small correction for the different temperatures of the e.m.f. measurements(25") and the n.m.r.measurements (31") can be made : its effect would be unimportantin the present context. The assumption that x = 1 (v = 3) gives close agreementwith the experimental value of 4j, and gross experimental errors would have to beinvoked to make the data fit in with any other integral value of x.X 4j (calc.10 0.35 0.321 0-71 0.682 0.8 1 0.793 0.86 0.85L = 8 L - 10A somewhat different treatment of similar n.m.r. measurements 10 has led tothe same conclusion. In this work, chemical shifts were studied as a function ofboth perchloric acid concentration and n. Mutual consistency of results at differentvalues of n was obtained only if the hydrogen ions are all isotopic forms of H3O+,with I = 0-68+0-01.D. MEASUREMENTS OF COMPOSITION OF WATER VAPOUR, OVER AQUEOUS ACIDS.-111this method the isotopic composition of a sample of water vapour in equilibrium withan aqueous solution (containing more than one isotope) is analyzed.There issome fractionation in the absence of added solutes. The effect changes when saltsare added, most of the change being associated with the anion,ll but a larger changeoccurs with perchloric acid. These phenomena arise from the fact that the com-position of the vapour in the presence of solutes reflects that of the free liquid water,i.e., water not held in solvation shells. The isotopic composition of the free wateris naturally affected by the fractionation equilibrium between hydrogen ions an88 PROTON TRANSFER TO OLEFINSwater. The most informative study by this method would be to examine the frac-tionation effect for different values of It, both close to zero and close to unity.Theonly measurements available to date are confined to low concentrations of deuterium.In the detailed evaluation of the results some problems arise from the anion effectand also a minor one from the presence of some undissociated (perchloric) acidat high concentrations and the resulting isotope fractionation between undissociatedperchloric acid and water. The equilibrium constant evaluated from these measure-ments, for the limiting case n-+O at 13.5"C isThis constant is related to L, on the basis of the rule of the geometric meanaccording to(1.9) L- 1/2(2x+ 1) KL = -2x+1so that it is again possible to calculate KL from e.m.f.values for L or, by combinationof eqn. (1.8) and (1.9), from the n.m.r. result for $j. Heinzinger and Weston 3 per-form the latter calculation and show that the data are mutually consistent on the as-sumption that x = l(vi = 3). The calculation for KL from eqn. (1.9) is givenbelow for different values of x. In a more accurate comparison the temperaturedifference between the e.m.f. (25") and vapour measurements (135") can be takeninto account.XKL (calc.)L - a L- 100 5.66 6.321 0.94 0.982 0.49 0-503 0.33 0.34E. OTHER MEAsUREMENTS.-The statistical calculation of L has been describedby Swain and Bader.2 The value obtained (8.2 at 25") falls into the range of valuesgiven by the other methods, but the reliability of the spectroscopic data on whichit is based has been questioned.3The combination of fractionation measurements on the exchange betweenaqueous hydroxide ions and water 12 with values of the ionic products for ordinarywater and deuterium oxide should likewise lead to information about L, but therequired data are not as yet sufficiently well known to give values of L of comparableaccuracy to those obtained by other methods.The main experimental approaches to the problem all lead to values of L in therange 8-10.Results from the e.m.f. methods are independent of any assumptionabout the structure of the hydrogen ion in solution, but results from methods Cand D are compatible with each other 3 and with the e.m.f. result 8 only on thesupposition that each hydrogen ion contains three hydrogen atoms in equivalentpositions.These three hydrogen atoms are isotopically fractionated relative towater. The measurements do not exclude the presence of further hydrogen nucleiwithin the structure of the ion but they do imply that such further positions (e.g., inwater molecules hydrogen-bonded to the H30+ group) have essentially the sameisotopic composition as the bulk water.As will be evident from the comments on each method, the individual deter-minations all involve some assumptions which cannot be justified in detail. How-ever, these assumptions are different in each case and, taken together, the variouV . GOLD AND M. A . KESSICK 89methods provide strong evidence for the foregoing conclusions.These con-clusions were reached without any consideration of protolytic equilibria or reactionvelocities in H20 + D20 mixtures.5 These provide an independent approach tothis problem. The significance of one such study in relation tQ the investigationssummarized above will now be discussed.2. ISOTOPE EFFECTS IN THE HYDRATION OF ISOBUTENEA variety of detailed mechanisms have been advanced for the hydration ofolefins.13 They share the common feature, required by experimental evidence,that the first, rate-limiting step of the reaction involves proton transfer from thehydrogen ion to carbon. Our reaction models are specific versions of this commonmechanism only in the respect of specifying the nature of the hydrogen ion fromwhich proton transfer takes place.The proton in transit is considered not to havereached its terminal location in the transition state, but to have acquired a uniqueposition, whereas in the initial state it occupied one of a set of equivalent positionsfor hydrogen (e.g., three in H3O+).The hydration of isobutene is an acid-catalyzed reactionleading to a single product. The equilibrium constant is very large and the rateof reversal is not significant in the context of this study. Attempts to detect generalacid catalysis have been unsuccessful.14nFIG. 1 .-Rate isotope effect; curve calculated for kH/kD = 1.40 ; r = 4.The two isotope effects to which the present discussion can be restricted arethe rate isotope effect, defined as the effect of isotopic composition of the mediu90 PROTON TRANSFER TO OLEFINSon the rate of disappearance of isobutene from solution, and the deuterium productisotope effect, which is determined by the relative copcentration of isotopes in thesolvent and in the one newly-formed C-H bond of the product [eqn.(2.2)].*5The results discussed in this paper all relate to ca. 0.44 M aqueous perchloric acidat 25". The measurements performed 15 also include product isotope effects ontritium in competition with protiurn or deuterium or with mixtures of the two inthe solvent. The tritium effects quantitatively confirm the measurements withdeuterium, i.e., they are in accord with the relation,l6 ( k ~ / k ~ ) = (kH/kD)1*442.The rate of hydration, measured by spectrophotometric observation of thedisappearance of olefin, varies with the deuterium abundance n of the medium ina non-linear fashion (fig.1). The (slightly extrapolated) rate ratio for H20 andD20 iskH/kD = 1.45&0*1. (2.1)The product of the reaction contains deuterium in only one of its nine other-wise equivalent carbon-hydrogen bonds. In this one newly-formed bond theabundance of deuterium (atom fraction m) is less than that in the solvent (atomfraction n). This product isotope effect can be expressed in terms of the ratio r,n(1- m)m(1- n)r = = 3.9 0.2which is, within the limits of experimental error, independent ofposition of the medium (table 1).(2.2)the isotopic com-TABLE 1 .-DEUTERIUM PRODUCT ISOTOPE EFFECTS?I0-1590-2040.3190.4060.4980.6000-6900.8100.930m0-0500.0520.1 120-1450.2 120.2720-3780-5250-760r3.6 f0-63.9 k0.63.7 30.34.0 f0.33.7 *0.24-0 f0-23.7 f0-23.9 h0.24.2 &0-3(The stated limits of error are estimated maximum errors due to the limits of precision ofthe isotope analysis.)We now consider these results in terms of reaction models.A.DIRECT PROTON TRANSFER FROM H3O+ TO ISOBUTENEExpressions for the velocity of such a process in a medium containing bothH2O and D20 have previously been derived 17 and can be adapted to the treatmentof product composition. According to this theory the rates of proton and deuterontransfer from all isotopic H30 groups (H20, HzDO, HD20, D30), taken together,to a substrate, in fixed concentration, can be written asuH = ( kHclQ)( 1 - n)( 1 - n + n 1' - ") 2, (2.3)VD = ( ~ ~ c / Q ) T I Z ' ~ ~ " ( ~ - ~ + I I ~ ' - " ) ~ .(2.4)In these expressions k~ and k~ are the rate constants of hydrion transfer from the(light) H3O group and the D30 group respectively; these are also the rate co-efficients for reaction in ordinary water and deuterium oxide respectively, on thV . GOLD AND M. A . RESSICK 91assumption that (transfer) medium effects for the substrate, hydrogen ions andtransition states cancel or can otherwise be neglected. The acid concentration cmeasures the sum of the concentrations of all isotopic hydrogen ions; accordingto the rule of the geometric mean, Q is given by the expression (1 --n +nZ)3 ; a isthe exponent of the Bronsted catalysis law, and I is the fractionation parameterdefined in 5 1 .The total rate should then be given by the sum of eqn. (2-3) and (2.4),the relative rate coefficient in an isotopically mixed medium being given byThe isotopic composition of the product will be given by the ratioand the product isotope effect r [eqn. (231 byr = kH/kD11+2a.Combining the results of eqn. (2.1) and (2.2) with eqn. (2.7) and taking I = 0.69,we obtain a = 0.85+0.1. In this calculation r, I, and k ~ / k ~ are subject to someexperimental uncertainty, the slightly extrapolated rate ratio being the weakestlink. The value of a indicates that proton transfer is far advanced in the transitionstate but not complete. (In the related hydration of styrenes, a is small enoughfor general acid catalysis to be detectable from experiments with aqueous buffers.18)The general correctness of this analysis can be tested by substituting a possiblevalue of a in eqn.(2.5) and hence calculating reaction velocities in H20+D,Omixtures. A calculated curve obtained in this way is shown with the experimentalpoints in fig. 1.Eqn. (2.5) and (2.7) can also be cast entirely in terms of fractionation para-meters : 199 20k,,/k, = ( 1 - n + n l ) - 3 ( l - n + n ~ , ) ( l - n + n ~ 2 ) 2 , (2Wwhere 41(= l/r) expresses the isotope fractionation [as defined in eqn. (1.5)] ofthe hydrion in transit and 4 2 that of the newly forming H20 group. It also followsthatEqn. (2.5) and (2.8) are completely equivalent ; their comparison affords insightinto the physical significance of the parameter a.The agreement between prediction and results provides a more stringent testof the adequacy of this model of proton transfer reactions than rate measurementsby themselves do.The analysis also indicates an approach to the determinationof the Bronsted exponent a for reactions in which the value of this parameter rendersthe determination of catalytic coefficients for general acid catalysis difficult. Thisprocedure therefore permitted the evaluation of all parameters required in the theory,without use of any adjustable coefficients other than the minor adjustments withinthe limits of experimental error.kD/kH = 41&-3. (2.9)B. INDIRECT PROTON TRANSFER FROM H3Of TO ISOBUTENEThe question then arises whether the results would also be reconcilable witha mechanism in which the hydrogen nucleus transferred to the isobutene molecul92 PROTON TRANSFER TO OLEFINSis not the same as the hydrion lost by the H30 group, but in which a water molecule(or more than one such molecule) acts as intermediary.Such a process would beintelligible if proton transfer from a hydrogen ion H90: 21 involved the protonmovements 223H HIH-0I0-@+ s1 3I04H1/\H HIn an isotopically mixed solvent the isotopic composition of the outer water mole-cules is, according to the equilibrium measurements, close to that of the bulk water(5 1). Such a model of proton transfer would therefore be inadequate if one wereto assume that the transfer from an outer water molecule is uninfluenced by the iso-topic nature of the inner hydrogen nucleus to which that water molecule is joinedby hydrogen-bonding.However, the position is different if one assumes that allhydrogen nuclei numbered 1-4 in the above formula are concerned in the transfer.In an H20 + D20 mixture, that reaction velocity will then depend on the respectivefractionation parameters 41-44 and Z, but the isotopic composition of the newC-H bond will depend on 4 3 alone. We can thus writekJk, = (1 - n + nl)-3(l - n + n4,)(l- n + r~4~)(1- n + 124~)(l- n + n+4)2, (2.10)with 453 = l/r, and k ~ , l k ~ = $1&-434: Z-3. Eqn. (2.10) contains similar factors aseqn. (2.8), but also some additional ones. The relationship between the presenttreatment and that based solely on the H30 model becomes even closer if we assumethat the rate of the reaction is influenced by the acidity of the hydrogen nucleusnumbered 1 by a factor of the form of the Bronsted catalysis lawt = GK"(with appropriate statistical corrections). In this way the rate becomes dependenton the isotope fractionation and isotopic composition of the H3O group.When41-~$2+-1, the transition state will be described by similar parameters to thoseof the simpler model, and k,/kH curves similar to that in fig. 1 can be constructed.It follows that the model of indirect proton transfer from H30+ (or direct protontransfer from HgOt) can likewise account for the results of this work in a rationalway and undoubtedly so can other, even more complex models.In the contextof isotope effects at present available, the assumption of indirect proton transferoffers no advantages over the simpler H3O+ model, nor does it permit additionalverifiable predictions to be made. The description of isotope effects in terms ofdirect transfer of H+ from H30+ to isobutene is adequate and simpler and, solelyfor that reason, is accordingly preferredV. GOLD AND M. A . KESSICK 931 Noonan and LaMer, J. Physic. Cilem., 1939, 43, 247.2 Swain and Bader, Tetrahedron, 1960, 10, 182.3 Heinzinger and Weston, J. Physic. Chem., 1963, 68, 774.4 Schwarzenbach, Z. Elektrochem., 1938, 44, 302. Schwarzenbach, Epprecht and Erlenmeyer,5 Purl=, J. Amer. Chem. SOC., 1959, 81, 263.6 Bergqvist and Eriksson, Acta chem. Scand., 1962, 16, 2308.7 Hindman, J. Chem. Physics, 1962, 36, 1OOO.8 Gold, Pruc. Chem. Suc., 1963, 141.9 Bigeleisen, J. Chem. Physics, 1955, 23, 2264.10 Kresge and Allred, J. Amer. Chem. SOC., 1963, 85, 1541.11 Googin and Smith, J. Physic. Chem., 1957, 61, 346.12 Heinzinger and Weston, J . Physic. Chem., 1964, 68, 2179.13 for references see Baliga and Whalley, Can. J. Chem., 1964, 42, 1019.14 Ciapetta and Kilpatrick, J. Amer. Chem. SOC., 1948, 70, 639.15 Gold and Kessick, Proc. Chem. Soc., 1964, 295 ; Pure Appl. Chern., 1964, 8, 421 ; J. Chem.16 Swain, Stivers, Reuwer and Schaad, J. Amer. Chenz. SOC., 1958, 80, 5885.17 Gold, Trans. Faraday Soc., 1960, 56, 255.18 Schubert, Lamm and Keefe, J. Amer. Chem. SOC., 1964, 86,4727.19 Saloniaa, Schaleger and Long, J. Amer. Chem. SOC., 1964, 86, 1.20 Kresge, Pure Appl. Chem., 1964, 8, 243.21 for references, see Eigen, Angew. Chem., 1963, 75, 489 ; Angew. Chem. (Int. Ed.), 1964, 3, 1.22 Kreevoy, Steinwand and Kayser, J. Amer. Chem. Soc., 1964, 86, 5013.Helv. chim. Acta, 1936, 19, 1292.SOC., 1965, 6718

ISSN:0366-9033    

DOI:10.1039/DF9653900084

出版商:RSC    

年代:1965

数据来源: RSC

摘要:

GENERAL DISCUSSIONH///’Ar+HR . C02Hs Ar I ’”\ H . . . 02C . R,’ 4 H . . . O///’ ‘\\ ,/+Ar c . R (1)\H . . . OH//’ArH + CH3 . C02H+Ar’ \\H . . . OzC.CH3 Haccording to the following line of reasoning.If the intermediate is supposed to be the same cation ArHi(1) for catalysis byhydrogen ions and by acetic acid then we can write the following four equilibria1 Streitwieser and van Sickle, J. Atner. Chem. SOC., 1962, 84, 254.2 Batts and Gold, J. Chem. Soc., 1964, 4284.9GENERAL DISCUSSION 95between reactants and intermediate, depending on which catalyst and which medium(H20 or D20) is used :1-1H,O++S+I+H,O K , (3)(4)HA + S $1 +A-' K2 ( 5 )DA + S +-Ii +K- K2 (6)_-D30++S+T'+ D20 Rlthe bar over a symbol designating solution in heavy water.On the basis of mechanism (1) and with some assumptions, it is possible to coni-bine rate measurements on various H-D-T exchange reactions to lead to thevalue 3.5 for the ratio K ~ K ~ I K I R ~ .In terms of equilibria (3)-(6)' this ratio ofconstants is the ratio of dissociation constants of acetic acid in ordinary and heavywater :[H30'][S] [%][S] [DA][D20]The directly determined ratio of these dissociation constants is 3.3. The satisfactoryagreement of this ratio with the value 3.5 supports the basis of the calculation andthat the same intermediate is formed in reactions (3) and (5).A similar treatment can be applied to the forward rate constants of the equilibria(3)-(6), or the quasi-equilibria between reactants and transition states :H30++S+TS, k ,D30+ +S+TS; ElDA+S+TS; k2___- - __HA+S+TS, k2 - - -The isotope effects on these reactionsandcan be combined to giveThe experimental value of this ratio is 1-85, very different from the value 3.3 of thefirst factor and indicative of the difference between transition states TS1 and TS2,the quotient [TS2][%;]/[TS;][TS,] having the value 0.55.The results of this investigation also allow us to give an answer to Mr.Bell'squestion about the possible consequences of neglecting secondary isotope effect96 GENERAL DISCUSSIONon the reversal step (- 1) of reaction (3) on the calculated isotope effects for the stepsof the exchange process. If a reasonable estimate of the secondary isotope effect,i.e., of the statistically corrected ratio of rate constants forandArHl +ArH + (H')ArHD'+ArD + (H')is taken, the primary isotope effect, expressed by the statistically corrected ratioof rate constants forArHz -+ArH + (H')andArHD' +ArH + (D'),has the value 8-05, whereas with neglect of the secondary isotope effect, the primaryisotope effect is 6.84.Prof.A. J. Kresge (Illinois Institute of Technology) said : In addition to supplyingvery useful data for aromatic hydrogen exchange in concentrated acids, Prof. Longhas convincingly argued for the existence of the conjugate acid of the aromatic asan intermediate in this reaction. However, the presence of this intermediate canbe deduced from two other pieces of kinetic data, viz., the reaction is subject togeneral acid catalysis and is not catalyzed by bases.For exchange to occur, the system must pass through a symmetrical configura-tion such as that in the species shown in eqn.(1) of Prof. Long's paper :L H 6 <*:This symmetrical species can be either an intermediate or a transition state; letus assume that it is the latter. Then, since the reaction shows general acid catalysis,this transition state must contain the conjugate base of the catalyzing acid attachedby a partial bond to one of the exchanging hydrogens. But if the species is to besymmetrical, both of the exchanging hydrogens have to be alike, and the otherhydrogen must be bonded to the same kind of base as well. This transition state,therefore, must contain HA and A- in addition to the aromatic, and, on this hypo-thesis, the reaction will be subject to base catalysis as well as to general acid catalysis.Since base catalysis is known to be absent, the symmetrical species cannot be a transi-tion state and must be an intermediate.The requirement that both of the exchanging hydrogens be attached to a basein a symmetrical transition state can be accommodated without base catalysis byinvoking a cyclic structure for the transition state such as ,that considered (and dis-missed) by Batts and Gold 1 :AIn exchange catalyzed by carboxylic acids, this structure would be strain-free becausedifferent oxygen atoms would be used for bonding to the different hydrogens.But1 Batts and Gold, J. Chem. SOC., 1964, 4248GENERAL DISCUSSION 97for catalysts such as the ammonium ion or the water molecule, strain-free con-figurations are not possible. This difference would be expected to result in ratesof reaction which are not correlated well by a single Bronsted relation.Since thereaction does obey a single Bronsted relation with considerable precision, such acyclic transition state can be eliminated from consideration.Prof. V. Gold (King's College, London) (communicated) : Prof. Kresge's commentthat a cyclic transition state for aromatic hydrogen exchange can be excludedsimply on the evidence of a single Bronsted relation presupposes that proton transfersdo not involve intermediate water molecules. If one were to write the reactionintermediate asH. .A. .H,' '.i/H-0 0-€3H LIstrain effects would no longer distinguish the different acids HA.This commenttherefore adds weight to the conclusion, suggested by our experiments on olefinhydration, that intermediate water molecules are not concerned in proton. transferto carbon.A carbon atom as a basic site is probably not significantly hydrogen-bonded tothe water-acid system, in distinction from lone-pair atoms (N, 0) commonly involvedin rapid proton transfer. A further example bearing on this question is the protontransfer between carbon and nitrogen which, in one direction, constitutes the measur-able rate-limiting step in the pyridine base-catalyzed halogenation of ketones and,in the reverse direction, represents a proton transfer from a pyridinium ion toolefinic (enolate) carbon.This reaction is characterized by considerable stericstrain in the transition state.1 One need not accept the unconventional and perhapsimprobable detailed suggestions that have been made about this transition state.1What is beyond dispute is the conclusion that this structure must be compact so that,in the reaction between 2,6-lutidine and pinacol, e.g., there is steric interferencebetween the methyl groups (Me) of the two moieties of the transition stateHMe-O /\ Me- -Me3C. CO . C . . . H . . . N / \HThis requirement of compactness would seem to preclude the interposition of a watermolecule between the participants.Prof. A. J. Kresge (Illinois Institute of Technology) said: Since our paper waswritten, we have extended our measurements to solutions of other strong mineralacids.I would like to present some preliminary data for nitric and phosphoricacids (fig. 4). Both monobasic acids give much the same rates of exchange whencomparison is made at the same concentration of HArH+, and these rates are1 Feather and Gold, J. Chem. Soc., 1965, 1752.98 GENERAL DISCUSSION- - . - - . - I - L . - - - - - - - I i- I 0 + Ib i o CWH+ICHA~FIG. 4.-Rates of loss of tritium from 1,3-dimethoxy-benzene-4-t compared to equilibrium pro-tonation of azulene.significantly less than those for exchange in the two dibasic acids, sulphuric andphosphoric.Mr. R. P. Bell (Oxford) said: The most certain way of establishing catalysisby bisulphate ions would be to study solutions of alkali bisulphates, with or withoutthe addition of sulphates, since in these solutions the hydrogen ion concentrationwould be low.I would like 10 enquire whether such experiments have been carriedout, or are feasible.Prof. A. J. Kresge (Illinois Institute of Technology) said: It is possible to measureexchange rates in bisulphate salt solutions, and we have done this, without addingsulphates, for another substrate. The difficulty here comes in sorting out thecontribution to the total rate made by the hydrogen ion. I am not sure that enoughRaman data for alkali bisulphate solutions exist to enable one to estimate the con-centrations of the various acidic species present. One could use the technique wedescribe in our paper, i.e., compare rates at the same Concentration of HArH+,and we have in mind to do this.The addition of sulphate salts might help, but here again the interpretation wouldnot be straightforward: bisulphate ion is a moderately strong acid and does notmake good buffers.In our work on 1,3,5-trirnethoxybenzene in fluoroacetic acidsolutions we had the trouble of changing pH with changing buffer concentration,and fluoroacetic acid is a slightly weaker acid than bisulphate ion.Dr. M. Spiro (Inzperial College) said: I would like to comment briefly on theresults shown on Prof. Kresge's fig. 4, which indicates that H3P04 and HzS04have much the same effect on the rate of the reaction. Since H2PO; has a muchsmaller dissociation constant than HSO,, the HzPOi ion would not be expectedto be as good a proton donor.However, we have recently deduced 1 from con-ductance and transference measurements that solutions of phosphoric acid containtriple ions of formula HzPO, . Hf . H2PO; which, as proton donors, would beintermediate in power between HzPO; and H3P04. It could well be that in thesolutions used by Prof. Kresge there was an appreciable concentration of such triple1 Selvaratnani and Spiro, Trans. Faraday Soc., 1965, 61, 360GENERAL DISCUSSION 99ions and that these contributed to the catalytic rate. The solutions may alsohave contained some dimeric molecules (H3PO& but it is only at high concentra-tions that these would be present in sufficient degree to affect the rate. I wouldlike to ask Prof. Kresge what concentrations of phosphoric acid were used in hiswork.Prof.A. J. Kresge (Illinois Institute of Technology) said: The experiments inphosphoric acid were done at stoichiometric acid concentrations of 1-5-53 M.We, also, find it surprising that the effects of phosphoric and sulphuric acid are sosimilar. Perhaps the reason for this is that phosphoric acid is weaker than sulphuricacid in its first dissociation as well as its second dissociation. In phosphoric acid,then, more of the excess rate could be due to catalysis by the completely undis-sociated acid, and this might make up for the smaller contribution from thedihydrogenphosphate ion.DP. M. Spiro (Imperial College) (partly communicated) : Some calculations havenow been carried out to estimate the relative contributions of the various species.The concentration chosen as an example was 4 M, which is inside the range usedby Prof.Kresge. The necessary stability constants were taken from ref. (1) andthe activity coefficients from Davies’ equation.1 This provides for the influence ofionic strength but makes no allowance for the undoubtedly big medium effects.The results are listed in the following table, in which &cid is the concentrationacid dissociation constant of the species named and in which the relative contri-butions to the rate u have been assumed equal to (cOncn.)(K~cid)+, as in the paperby Kresge et al. Statistical factors have not been taken into account.species HtPOZ H3m4 HzPOZ.H+ .H2POS (H3po4)2concn., M 0.06 2-6 0.5 0.14Kacid z x 10-7 0.013 0.004 ca.1V 3~ 10-5 0.3 0.03 ca. 0.1Thus undissociated H3P04 clearly has the greatest effect on the rate but tripleions do contribute significantly. Dimeric (H3P04)2, evidence for whose existenceis rather indirect, would appear to have an even larger influence at 4 M. Thesecalculations show how important it is to know which species are present beforzcatalytic efficiencies are assigned. Unfortunately our knowledge of the compositionof concentrated solutions of this sort is rather sketchy.Prof. F. A. Long (Cornell University) said: In our studies of azulene, we haveobserved that a number of the azulenes are unstable in strongly acidic solutions.Furthermore, the decomposition is greatest when there is an approximately equiv-alent amount of protonated and unprotonated species present.Azulene itselfdecomposes moderately and decomposition of some of the substituted azulenes issignificantly more. It is my impression that the rates of decomposition whichwe had observed were somewhat higher than those suggested by Kresge. Further-more, it is my recollection that the rate of decomposition was substantially greaterin sulphuric acid than in perchloric. This also, Dr. Kresge seems not to havefound. I wonder if he would comment on these statements.Prof. A. J. Kresge (Illinois Institute of Technology) said: We also observed thatazulene is unstable in strongly acidic solutions. Since the rate of decompositionis greatest when roughly comparable amounts of protonated and unprotonatedaromatic are present in the solution together, we believe that the reaction which occurs1 Davies, Ion Association (Butterworths, London, 1962), chap.3100 GENERAL DISCUSSIONis an electrophilic substitution of the cationic conjugate acid on the neutral molecule.There is some evidence to support this, e.g., in strongly acid solution, phloroglucinolis converted to “ phloroglucid ” which has a dimeric structure.1 Since this reactionis bimolecular, it depends on the square of the total azulene concentration, and,by working at high dilutions, we were able to reduce it to such an extent that it didnot interfere with our measurements.Prof. A. J. Kresge (Illinois Institute of Technology) (communicated) : The decom-position of azulene in acidic solutions has recently been found to be second orderin the aromatic, first order each in protonated and unprotonated forms, and therate constant reported is consistent with the rates of disappearance which we ob-served?Prof.F. A. Long (Cornell University) said: I would first comment that Prof.Gold is somewhat too modest about the significance of his results. For someyears, recognizing that the original Gross “ cubic ” formulation for deuteriumsolvent isotope effects was not the only way in which the data could be fitted, ithas been important to attempt to obtain data that explicitly showed behaviourwhich required a cubic or at least some high order of expression to give a fit. Asfar as I know, Gold’s fig. 1 which does show the behaviour expected from a cubicequation, is the first instance where this direct consequence of the Gross formulationhas been shown.This alone is a most significant result.I would make two different points about Gold’s results. A first point relatesto the formula for the solvated proton, specifically H3O+, or H90;t. In analyzingdata for the deuterium solvent isotope effects, we have noted also that one can obtainreasonable fractionation factors utilizing the I 3 9 0 model, values which are aboutas satisfactory as for H30+. Whether fractionation factors for related moleculeswill be as consistent under the H90,+ treatment as they now are under the H3O+treatment is unknown. In any case, I agree that other things being equal, the simplerH30+ treatment is to be preferred.The other comment concerns the possible role of medium effects.Prof. Salomaaof the University of Turku, Finland, when he was working at Cornell and sincethen, has been concerned with an interesting way to get values for the constant L.It consists of obtaining the ratio for the K1K2 product of a dibasic acid whichis added to the solution in anhydride form. Two examples with which Salomaahas worked are COZ and S02. Considering the first of these, the equations whichare relevant areC02+2H20 = H30++HC0,, K,HC0;+H20 = H30++COg- K2For equilibrium in this system, assuming that one works at constant activity,i.e., constant pressure of the anhydride C02, the degree of hydration of this speciesis immaterial and one obtains directly the equation :In other words, this gives a possible way of obtaining the value of L directly.Usingliterature data for the CO2 system, Salomaa reports that the resulting value of Lis about 14. His own experimental studies with sulphur dioxide have led to anL value of about 19. Both of these are large compared to the best experimentalvalue of L of 7-1 1. In a personal communication, Salomaa has suggested that this1 Beilsteiti, VI, 1099. 2 Myhre and Anderson, Tetrahedron Letters, 1965, 1497GENERAL DISCUSSION 101high L value results from the fact that there is a medium change in going from H20to D20, an effect which is for most situations absorbed into the fractionation factor,treated as a parameter. He further speculates that, because he has made his studiesof Kl and K2 under different medium conditions, the medium effect is here evident.This suggests that the assumption of ideality of the solvent, which most of theseanalyses of the H20+D20 system utilize, may be doubtful.Perhaps Prof. Goldwould comment.Prof. V. Gold (King’s College, London) said: The question whether protium-deuterium isotope effects depend on the isotopic composition of the medium ispartially answered by the results in table 1 of the paper by Gold and Kessick. Theisotope effect Y, which is an experimental quantity independent of the assumedreaction mechanism, shows no significant trend with the isotopic composition ofthe medium expressed by n. The same conclusion follows from the correspondingtritium product isotope effects in H20+D2O mixtures containing a trace amountof tritium.1Prof.Maurice M. Kreevoy (University of Minnesota) said : I would like to supportthe Gold and Kessick view on the invariance of Y under changes in the isotopiccomposition of the solvent. We have determined isotope effects both from pro-duct ratios and rate ratios for cleavage of vinylmercuric iodide 2 and allylmercuriciodide.3 I have described the mechanism of the latter reaction in my comment onthe paper of Caldin and Kasparian. The former is similar. Our results arequalitatively similar to those of Gold and Kessick. The table illustrates the in-variance of r under a change of a factor of 15 in the hydrogen to deuterium ratioin the solvent.(D/H) soh. 0.677 1-09 1.60 1.90 2.75 3-15 4.50 10-47r 7.16 7.34 7.37 7.18 7.24 7.44 6.83a 7-36 av. a 7-30f0.09(a) The average and the average deviation from the mean omit the value 6.83 which was theresult of the first measurement and varies from the mean by 5 average deviations.I also want to comment on the possibility of distinguishing between the directtransfer mechanism,A-@-+B,and the double transfer mechanism,HA--@-+O--@-+B,Ion the basis of isotope effects.I agree that this does not seem possible when HAis the hydronium ion. However, this can be done for monobasic acids. For thedirect transfer, with some reasonable assumptions, it can be predicted that RH/RD,the product ratio is a mixed solvent, should be given by [HA]~HA/[DA]~DA. Forthe double transfer mechanism RH/RD should be given by (H/D)so~v.(kHA/kDA) R,where R is the ratio of the following rates :1 Gold and Kessick, J . Chem. SOC., 1965, in press ; Pure and Appl. Chem., 1964, 8, 421.2 Kreevoy and Kretchmer, J . Amer. Chem. Soc., 1964, 86,2435.3 Kreevoy, Steinwand and Kayser, J. Amer. Chem. SOC., 1964, 86, 5013102 GENERAL DISCUSSIONMIA-@ + 0-0 -+ SMIA--@+O-@+S(M can be either hydrogen or deuterium but must be the same in both reactions.)There is some reason to believe that R should be not far from unity, and, in any event,there is no reason why (H/D)so~v.R should be consistently equal to [HA]/[DA]..We are now trying to evaluate the required quantities for allylmercuric iodidecleavage.Mr. Brian Case and Dr. Roger Parsons (University of Bristol) (communicated) :The equilibrium constant L (eqn.(1.2)) of the paper by Gold et al. can be evaluatedfrom L’ (eqn. (1.1)) without approximation from experimental data. The ratiois related to the difference in the standard real free energies a of solvation of thechloride ion in H20 and in D20 byThe difference in standard real free energies of solvation is a directly accessibleexperimental quantity since the standard e.m.f. of a cell such as 1L/L’ = [cl-]2/[cl-]2 (1)RT In (L/L’) = -2(a:;o-a:;o).Ag I AgCl I NaCl in HZO I air I NaCl in D,O I AgCl I Ag(2)(3)E*= - (a??- a$?)/F. (4)sat. sat.is given byThe potential of a cell with an air gap in which the field is zero can be measureddirectly using the Kenrick method, or indirectly using Kelvin’s method, or theradioactive method.*The calculations of Swain and Bader2 effectively replace the difference of realsolvation energy in (2) by the difference of chemical solvation energy as calculatedfrom the difference in the librational behaviour of the solvent in contact with theion. This is equivalent to the assumption that the surface potential x of H20 isequal to that of D20. In view of the importance of the libration of the solventmolecules in the isotope effect on solvation parameters it is possible that the sub-stitution of D20 for H20 in the surface may change x owing to a change in theaverage dipole orientation. While this effect is probably small because the valueof x itself is small ( N 100 mV for HzO) the effect on L may not be negligible since itis evident from (2) that L/L’ would be changed by a factor of 2 if ~D20- xH*0 were9 mV.We have recently made preliminary measurements of the e.m.f.of cell (3) byKenrick’s method using an apparatus similar to that described by Randles.4 Wefind a value of E* = -3.5f2 mV. This corresponds to a ratio L/L’ of 0-762&0.12and L = 13-9&2-0 which is somewhat higher than the value proposed by Gold andKessick. These results suggest that there is in fact a difference in x between waterand D20. Using Swain and Bader’s estimate of the difference in chemical solvationenergy for Cl-, this amounts to 6.6+2 mV.1 Parsons in Modern Aspects of Electrochemistry, I, ed. Bockris (Butterworths, London, 1954),4 Randles, Trans.Faradzy Soc., 1956, 52, 1573.pp. 115-127. 2 Swain and Bader, Tetrahedron, 1960, 10, 182GENERAL DISCUSSION 103Prof. M. Eigen (Gottirigen) said: Although there is no doubt that H3O+ in wateris further hydrated it is completely adequate-as Prof. Gold pointed out-to ascribehis results to the H30+ species. There may be some relation to the migrationmechanism of the proton in which the rate-limiting step is the incorporation of furtherH2O molecules in the outer hydration shell of H30+. Similarly, in proton transferreactions to H-bonded acceptors the partner must somehow become incorporatedin the hydration structure of H3O+. The proton will be transferred only if thenew acceptor finally becomes the centre of hydration; otherwise the proton willjump back into its original central position with a rate comparable to its valencevibration frequency (cf. Franck-Condon principle).Thus it is really the H3Of(i.e., the centre of the hydration complex) which has to reach the new acceptor,and the higher multiplicity of protons in the outer hydration sphere does not comeinto play. The situation might have been different for a mechanism in which aproton jump over the whole diameter of the hydration complex (i.e., from the originalcentre to a newly formed one) were the rate-limiting step.Prof. H. Zollinger (Swiss Federal Inst. Technology, Zurich) (communicated) : Theproblem of whether indirect proton transfers via a water molecule do or do not occurin reaction mechanisms has been investigated in connection with general base catalysisof diazo coupling reactions of 1-naphthol and 1 -naphthylamine derivatives in the2- and 4-positions.1 The effect of water molecules relative to the respective effectH H0..'* .. /f0'' 1, HCsoyIof acetate ions as proton acceptors for the intermediates in diazo coupling is 2.2times larger in the 2- than in the 4-position of 1-naphthol-3-sulphonic acid.Similarly, the effect of water relative to the respective effect of secondary phosphateions is 5.7 times larger for the diazo coupling intermediate in the 2-position of1 -naphthol-3-sulphonic acid than in the 2-position of 1 -naphthylamine-3-sulphonic.Together with the observation 2 that the activation entropy of diazo couplingsin o-positions are smaller (more negative) than that of the respective reactions inp-positions, the most feasible explanation of these results consists of a protontransfer (I) from the 2-position to the naphtholate oxygen.The water moleculeis fixed by a hydrogen bond to the oxygen and, therefore, is in a sterically favourableposition to accept a proton from the 2-position. In the 4-position such a transferis impossible. It is less probable with coupling at the 2-position of 1-naphthylaminebecause of the lower degree of solvation of the NH2 group in comparison of theO0 group of the naphtholate.Prof. V. Gold (King's College, London) (communicated) : The decision whethera particular item of experimental evidence is convincing is bound to be a subjectiveone. Prof. Long's generous remarks about our work should, therefore, not be al-1 Zollinger, Chem. and Ind., 1965, 885.2 Stamm and Zollinger, Helv. chirn. Acfa, 1955, 40, 0oO104 GENERAL DISCUSSIONlowed to pass without some mention of the fact that Purlee and Taft 1 studied thedependence of rate of olefin hydration on the isotopic composition of H20+DzOmixtures several years before us.Dr. A. Gandini and Dr. P. H. Plesch (University of Keele) said: Under strictlyanhydrous conditions the protonation of styrene and other phenylolefins by per-chloric acid in methylene dichloride is a relatively slow reaction which only goesto completion if the [HC104] is at least 102 [olefin]. We studied the reactionbetween styrene and perchloric acid with a high-vacuum spectroscopic device byfollowing the absorption at 427 mp due to the 1-phenylethyl ion. At roomtemperature (20-22") the rate law isd[CH,CHPh]/dt = k[HC104]a[C8H8],with k = (2-35&0.1) x 102 l.* mole-* sec-1 for [CgHg] in the range 10-4-10-3 Mand [HC104] at least a hundred times greater. The fractional power in acid con-centration indicates that the acid is associated in methylene dichloride. Only themost careful technique gave closely reproducible results since the carbonium ionsare very sensitive to traces of water.21 Purlee and Taft, J. Arner. Chem. SOC., 1956,78, 5807.2 Gandini and Plesch, J. Chem. SOC., in press

ISSN:0366-9033    

DOI:10.1039/DF9653900094

出版商:RSC    

年代:1965

数据来源: RSC