GENERAL DISCUSSIONProf. M. Eigen (Gottingen) said: Dr. Grunwald mentioned our results on therate of H-bond formation and dissociation. The reaction we have studied was thedimerization of ecaprolactam in non-polar solvents.1 On changing to aqueousmedia the solvent now competes for H-bonding. It turns out, however, that therates of dissociation change only slightly; so that the figure of 109-1010 sec-1 isprobably correct. The rates of H-bond formation are smaller in H20 than in non-polar solvents because of this competition. From studies of the rates of structuralconformation changes (helix +random coil) of polypeptides in aqueous solutionthe above figure for the rate of dissociation of a single H-bond (in the unwindingof the helix) was also obtained. Specific influences which might be related to thephenomena reported by Dr.Grunwald were observed 2 in the recombination ratesof substituted ammonium compounds with OH-. We interpreted these resultsas a strengthening of the H-bond structure at the site of proton transfer induced bythe substituted (non-polar) groups. The strengthening of the H-bond structurewould increase the reaction cross-section for proton transfer and thus lead to anincrease of the rates.Prof. R. M. Noyes (University of Oregon) said: The observation that protontransfer to phosphorus is slower than that to nitrogen raises the interesting questionof the extent to which equilibria and rates of proton transfer reactions can be general-ized in simple terms. The proton is a unique chemical species in that it has no extra-nuclear electrons.Hence it will locate in a region occupied by an electron pairand will be a sensitive probe for indicating regions of high electron density. In anysystem at equilibrium, the protons will be located in regions of maximum electrondensity. That this effect is large can be illustrated by the isoelectronic bases F-,OH-, NH; and CH;. The density of excess negative charge in this series variesby about a factor of 4, but the ionization constants of the acids HF and CH4 probablydiffer by 50 powers of 10 in aqueous solution.Although equilibrium constants can be interpreted from electron densities,rates of proton transfers do not correlate simply with equilibrium acidities. Thus,proton transfers to and from carbon tend to be much slower than those to andfrom oxygen.When acetylacetone ionizes to CH3COCHCOCH; , the electrondistribution in the anion is very different from that in the neutral conjugate acid.Transfer of a proton to the anion will be accompanied by polarization effects thatgreatly shift electron positions. Such polarization will be manifested by a greateractivation energy than when the proton is transferred to an oxygen atom where theelectron cloud is already localized.In summary, the position of equilibrium in proton transfer reactions will bedetermined by electron densities in protonated and unprotonated species, whilethe rates of these reactions (if they are not diffusion controlled) will be determinedby polarizability effects.Prof.B. E. Conway (Ottawa) said : Prof. Noyes raised the question of the isotopeeffect in proton conductance and Wannier’s work.3 In the latter, H3O+ ion rotationwas considered following the work of Huckel.4 This mechanism is now not gener-1 cf. Ber. Bunsenges. physik. Chem., 1963, 67, 819.2cf. Angew. Chem. int. ed., 1964, 3, 1.3 Wannier, Ann. Physik., 1935, 24, 545, 569.4 Huckel, 2. Elektrochem., 1928,34, 546.13GENERAL DISCUSSLON 131ally accepted.1 The H/D isotope effect in the mobility is close to J2 but signifi-cantly temperature dependent.2 The ,/2 value is equal to the expected ratio ofdielectric relaxation frequencies for H20 and D20 3 and hence supports the rotationcontrolled proton conductance mechanism.Prof. M. Eigen (Gottingen) said: Prof.Strehlow concluded from his observationsthat for the catalysis by H30+ the rate-limiting step of the hydration reaction isthe transfer of the proton to the carbonyl group. Assuming a pK-value of thecarbonyl group he gives in his paper the rate of protonation should indeed be ofthat order of magnitude. On the other hand, if we consider the catalysis by acidsother than H30+ there is a linear dependence of log k on pK with a slope appreciablysmaller than one. We must conclude that in those cases the simple protonationof the carbonyl group is slower than the hydration since for protonation log kdepends linearly on pK with a slope of one. (The reverse reaction cannot be fasterthan diffusion-controlled.) Therefore, as I pointed out in my paper, we have toassume a co-operative mechanism for the hydration.It might be possible thatthe two mechanisms, i.e., the stepwise mechanism with a rate-limiting protonationof the carbonyl group (followed by fast hydration) and the co-operative mechanism(requiring the water molecule in the right position for hydration) lead to equal ratesfor H30f. This means that the two lines for log k as functions of pK intersect atthe H30+-pK. (For pK> - 1 the k-values for the stepwise mechanism are belowthose for the co-operative mechanism.) Such an accidental intersection at the H38 +-pK cannot be excluded from the present data, but neither can it be safely concluded.The results for non-aqueous media (dimethylsulphoxide and acetone) might bemisleading if a preferential hydration of H3O+ is present (the H20 concentrationswere >2M).Thus it is also possible that even for H3O+ as a catalyst only theco-operative mechanism is valid. In water as a solvent the lifetime of the proton-ated carbonyl group is so short that it seems a little difficult that all the steps followingprotonation (i.e., orientation, binding and splitting of H20 at the carbonium site)occur faster than the deprotonation of the carbonyl group. Further experimentsin non-aqueous solvents at lower 1320 concentrations might answer this question.The present data are not sufficient to prove either mechanism for H3O+, but theyclearly favour the co-operative mechanism for catalysis by acids with a pK> - 1.Mr. R. P. Bell (Oxford) said: Two pieces of evidence can be quoted to supportthe assumption that a hemihydrate of acetaldehyde is formed in concentratedaqueous solution.(i) The corresponding hemihydrate of monochloroacetaldehydecan be isolated.4 (ii) Cryoscopic and vapour pressure measurements with aqueousformaldehyde solutions show the presence of (CHzOH)20 and of higher polymers.5Prof. Maurice M. Kreevoy (Utziversity of A4innesota) said: Dr. Evans and Dr.Miller, in the Physical Chemistry Laboratory, Oxford, have recently studied theacetaldehyde hydration, dehydration rates in dilute aqueous perchloric acid solu-tion by n.m.r. methods and also by a method based on trappigg the acetaldehydewith semicarbazide as it is formed. Dr. Evans obtained 560 1. mole-1 sec-1 forkCA at 25.0" by the scavenging method.Since the equilibrium constant,1 e.g., see Conway in Modern Aspects of Electrochemistry, vol. 111, chap. 2, ed. Bockris and2 Lewis and Doody, J . Amer. Chem. Sac., 1933,55,3504. LaMer and Mason, J. Chem. Physics,3 Auty and Cole, J. Chem. Physics, 1952, 20, 1309 ; cf. Davidson, Can. J. Chem., 1957, 35, 458.4 Natterer, Monatsh., 1582, 3, 449.5 Bezzi and Iliceto, Chim. Ind., 1951, 43, 212. Bezzi, Dallaporta, Giacometti and Iliceto,Gazz. cliim. Ira!., 1951, 81, 915.Conway (Butterworths, London, 1964).1935, 3, 406132 GENERAL DISCUSSION[CH3CH(OH)2]/[CH3CHO] is 1.17 1 this leads to kAC, 655 1. mole-1 sec-1. Dr.Miller studied n.m.r. line widths as a function of HC104 concentration and obtained605 1.mole-1 sec-1 for kAC and 490 1. mole-1 sec-1 for kcA, both at 26". The generalagreement among these values, the older values, and those of Ahrens and Strehlowgives a good deal of support to the various non-traditional methods used. Muchof the residual scatter seems traceable to the variety of temperatures involved.Prof. F. A. Long (Cornell University) said: Dr. Eigen noted the possibility thattransition states for certain systems where both acids and bases serve as catalystsmay involve a simultaneous attack by an acid and a base with presumably watermolecules playing a connecting role. In his paper Dr. Strehlow does not requiresuch participation, but he does refer to it as sometimes needed; consequently,our data on some kinetic solvent isotope effects may be of interest.Dr. Hung,Dr. Robinson and I have been studying deuterium solvent isotope effects for the muta-rotation of tetramethylglucose as catalyzed by the following acids and bases : HOAc,OAc-, H3O+, H20. The reason for utilizing tetramethylglucose is to eliminatethe irrelevant exchangeable protons of ordinary glucose and to minimize the numberof isotopic fractionation factors which must be considered. For every one of thecatalysts, it is possible formally to fit the data by assuming a conventional mechanism,i.e., direct attack of base on the hydroxylic proton for base catalysis and for acidcatalysis preliminary protonation on the ether oxygen in a pre-equilibrium stepfollowed by attack by a base of the hydroxylic hydrogen, i.e., the familiar two-stepacid-base catalysis mechanism.The alternative is to assume for acid catalysis that there is simultaneous attackby the acid and a solvent molecule.In accordance with Eigen's suggestion thiscould be synchronous by the utilization of some connecting solvent water molecules.We have analyzed our deuterium solvent isotope data for the acid catalysts H3O+and HOAc utilizing a simultaneous mechanism with the intervention of two solventwater molecules. It is possible to get good agreement between the data for experi-mental results by selecting plausible fractionation factors, treated as parameters.Further, the fractionation factors which are needed both for the acid catalysis andfor the related base catalysis form a satisfactory internally consistent set of para-meters, i.e., the fractionation factors needed to explain the data for catalysis bythe base, H20, can be taken over directly and used for catalysis by the acidic speciesH3Of.One cannot utilize this result as proof of a simultaneous mechanism.Thedifficulty is that there are too many parameters to permit a unique solution.However, the simultaneous mechanism does give a satisfactory fit, which is at leastas good as can be obtained by the two-step mechanism.Dr. C. F. Wells (University of Birmingham) said: Dr. Strehlow and Mrs. Ahrenshave suggested as evidence against the rapid pre-establishment of an equilibriuminvolving the formation of protonated aldehyde the fact that their rate constantdoes not decrease proportionately as the Hammett acidity function HO increases.I submit that the use of Ho in such a manner is not valid, as HO measures the totalbasicity of these aldehyde + water mixtures to an added indicator.A knowledgeof the partition of protons between water and aldehyde, such as the concentrationquotient Kc = [CH3CHOH]/[CH3CHO] [(H20)4H+], is really required ((H20)4H+is the hydronium ion). We have shown 2 ~ 3 for various oxygen-containing organic41 Lombardi and Sago, J. Chem. Physics, 1960,32,635.2 Wells, Nature, 1962, 196, 770.3 Gila and Wells, Nature, 1964, 201, 606GENERAL DISCUSSION 133molecules that curves of HO against concentration with p-nitroaniline as the indicator-base B can be analyzed at the water-rich end in terms of the competition of B forprotonated water (H20)4H+ and protonated oxygen-containing molecules.Valuesof the concentration quotient Kc calculated for some carbonyl compounds 2, 3 in5 % solution from such an analysis are given in table 1, together with some valuesfor simple alcohols 1-3 for comparison. Acetaldehyde should have a value a littleTABLE 1.KC KC(moles-1 1.) (moles-1 1.)methanol 0.088 methyl ethyl ketone 0.99ethanol 0.27 ethyl acetate 0-96isopropanol 0.34 acetic acid 0.14acetone 0.63less than that of acetone. Protonation constants obtained in H20 + H2S04 mixtures,such as those quoted by Dr. Strehlow and Mrs. Ahrens, are considerably smallerthan our constants and are not applicable to their largely aqueous system. Ourhigh values of Kc for ketones suggest that the rate constant for protonation ofacetaldehyde will be high.Therefore, the possibility of the rapid pre-establishmentof the protonation equilibrium in the hydration of acetaldehyde is not eliminatedby equilibrium measurements when the appropriate data are used. These highvalues of Kc do not preclude the possibility of a concerted mechanism for thecatalyzed hydration.Prof. Dr. H . Strehlow (Max-PIunck-Inst. Physik. Chem., Gottingen) said : TheHammett function Ho is a measure of the thermodynamic activity of the hydrogenion and therefore determines equilibria involving these ions-only approximately-because of the non-thermodynamic assumptions involved. The observed maxima ofHammett functions in many non-aqueous solvent + water mixtures at constantconcentration of a strong acid, are explained by an increase of the basicity of waterwith its depolymerization brought about by the added organic solvent.4 Thatwater is the stronger base in alcohol containing a small amount of water has beendemonstrated from conductivity measurements (e.g., ref.(5)). A pK-value ofabout - 8 obtained for protonated carbonyl compounds by HO measurements withrather concentrated sulphuric acid is applicable to dilute aqueous solutions sincethe reference state of the Hammett function HO is the aqueous solution of a strongacid at unit molality corrected for non-ideal behaviour at that concentration.Dr. C . F. Wells (University of Birmingham) (partly communicated) : Braude’s 5explanation of the increase in HO in water-rich conditions in terms of the breakdownof the water structure is not applicable at low concentrations of organic molecules.Glycerol has an apparent Kc of zero, and the addition of up to 50-60 % glycerolhas no effect on the pK of p-nitroaniline : 6 the maximum density data 7 suggestthat structure breaking does not occur at low concentrations of alcohols.Structurebreaking will occur at higher concentrations of oxygen-containing organic molecules ;e.g., my relationship 8 ~ 9 gives curved plots 10 at alcohol concentrations > 50-60 %.1 Wells, Nature, 1962, 196, 770.3 Wells, Trans. Faraday SOC., in press.4 Braude and Stern, J. Chem. SOC., 1948, 1976.5 Strehlow, 2. physik. Chem., 1960, 24, 240.6 Braude and Stern, J.Chem. SOC., 1948, 1976.7 Wells, Trans. Faraday SOC., in press8 Mitchell and Wynne-Jones, Disc. Faraday SOC., 1953, 15, 161.10 Giles and Wells, Nature, 1964, 201, 606.2 Giles and Wells, Nature, 1964, 201, 606.McHutchison, J. Chem. SOC.,1926, 1898. 9 Wells, Nature, 1962, 196, 770134 GENERAL DISCUSSIONWynne-Jones has suggested that in H20 + H202 mixtures structure breaking probablyoccurs at very low concentrations of H202, and in agreement with his maximum den-sity data,l plots of my relationship 293 are curved at very low concentrations of H202.2Probably the difference 4 between our measurements and those in H20 + H2SQ4mixtures 5 : 6 is largely due to the different species of hydronium ion (H20)hH+involved: in our aqueous system (and that of Dr.Strehlow and Mrs. Ahrens)/i = 4,7 whereas in the 25-70 % H2S04 mixtures 9 9 6 where the mole fraction ofwater is small h<4 and will approach h = 2.7, 8 These latter H2S04 mixturescorrespond to the breakdown in water structure of Braude.9 The proton affinityin (H20)4H+ will be partially satisfied by the structural hydrogen bonding withinthe entity and will therefore be less than in (H20)2H+, i.e., (H20)2H+ will be aweaker acid than (H20)4H+ to an oxygen-containing organic molecule.Dr. C . F. Wells (University of Birmingham) (communicated) : Prof. Conwayhas suggested that the minimum in proton mobility in water+methanol mixturesmust be related to the basicities of water and methanol. However, the basicity ofwater and methanol vary with cornposition.10, 11 In passing from pure water to puremethanol, the protonated species gradually changes from (H20)4H+ to water-solvated CH30H2,2 until sufficient methanol is present to break down the waterstructure,l2, 13 when the equilibrium will tend to shift from CH30H2 to HiO, withHZO less acidic than (H20)4H+.Only at nearly 100 % methanol will this equilibriumshift to methanol-solvated CH30H2.14, 15 This suggests that (H20)4H+ in structuredwater and methanol-solvated CH30H2 in methanol involve a higher proton mobilitythan CH3OH2 and HZO in the collapsed water structure. Minima in proton mobilityhave also been observed for mixtures of water with ethanol, isobutanol, n-butanoland n-propanol.14~ 16Prof. B. E. Conway (Ottawa) said: I agree substantially with the remarks ofDr.Wells; we had, in fact, made some semi-quantitative estimates of the freeH30+ concentration in methanol+water solutions of HCl and related this to thecomposition at which the minimum conductance occurs. The minimum is ob-served only at high methanol concentrations where MeOH; -+MeOH proton transferevents can compete effectively with the alternative process of MeOHi +H2O protontransfer. The change of basicities of water and methanol with changing com-position will be an important factor in more quantitative calculations as Dr. Wellspoints out.+f++f1 Mitchell and Wyiine-Jones. Disc. Faraday SOC., 1953, 15, 161. McHutchison, J. Chenz. SOC.,3 Giles and Wells, Nature, 1964, 201, 606.5 Campbell and Edward, Can.J. Chem., 1960,38, 2109.6 Den0 and Wisotsky, J. Amer. Chem. Soc., 1963, 85, 1735.7 Bell, The Proton in Chemistry (Methuen, London, 1959), chap. VI.8 Wyatt, Disc. Faraday SOC., 1957,24, 162.10 Salomaa, Acta Chem. Scmtd., 1957, 11, 125.11 Wells, Nature, 1962, 196, 770. Giles and Wells, Nature, 1964, 201, 606.12 Braude and Stern, J. Chenz. SOC., 1948, 1976.l 3 Wells, Trans. Faraday SOC., in press ; see also this Discussion.14Goldschmidt and DahII, Z. physik. Chem., 1925, 114, 1 ; 2. physik. Chem., 1924, 108, 121.Goldschmidt and Thuesen, 2. physik. Chem., 1912, 81, 30. Thomas and Marum, Z. physik.chem., 1929, 143, 191.16 Goldschmidt, Z. physik. Chem., 1926,124,23. Goldschmidt and Mathiesen, Z. physik. Chem.,1926, 121, 153. Goldschmidt and Thomas, Z.physik. Chem., 1927, 126, 24.1926, 1898. 2 Wells, Nature, 1962, 196, 770.4 Wells, Trans. Furaday SOC., in press.9 Braude and Stern, J . Chem. SOC., 1948, 1976.15 De Ligny, Rec. trav. chim., 1960, 79, 733GENERAL DISCUSSION 135Prof. R. J. Gillespie (McMaster University) said: Dr. Brouwer has claimed thatthe collapse of the methyl peaks in the n.m.r. spectra of solutions of hexamethyl-benzene, prehnitene, and 5,8-dimethyltetralene in anhydrous hydrogen fluoride withincreasing temperature shows that an internal proton shift in the formed carboniumions occurs at temperatures lower than that at which exchange with the solventoccurs. It is also claimed that for hexamethylbenzene the reaction rate is the samein the less basic solvents HF+BF3 and HF+SbF,. However, for solution of hexa-methylbenzene in fluorosulphuric acid, Dr. Birchall and I were unable to obtainany evidence for this internal proton shift.1 The n.m.r. spectra obtained were con-sistent with the collapse of the methyl peaks being due to proton exchange betweenthe carbonium ion and the solvent. In support of this conclusion we found thatthe solvent line broadened simultaneously with the collapse of the methyl peaksand that the rate of proton exchange was increased by addition of fluorosulphatewhich increases the basicity of the solvent and decreased by addition of SbF5 whichdecreases the basicity of the solvent. It seems remarkable that an internal protonexchange should be observed in hydrogen fluoride but not in the less basic solventfluorosulphuric acid.1 Birchall and Gillespie, Can. J. Chern., 1964, 42, 502