Temperature dependence of enthalpy, entropy, and heat capacity of activation in the neutral ester hydrolysis in 2-butoxyethanolndash;water solutions

摘要

J. CHEM. SOC. PERKIN TRANS. n 1983 Temperature Dependence of Enthalpy, Entropy, and Heat Capacity of Activation in the Neutral Ester Hydrolysis in 2-Butoxyethanol-Water Solutions Liisa T. Kanerva Department of Chemistry and Biochemistry, University of Turku, SF-20500 Turku, Finland The kinetic behaviour of the neutral hydrolysis of methyl trifluoroacetate in 2-butoxyethanol-water solutions when x, is close to 0.98 was studied and the calculated activation parameters have been com- pared with the kinetic data of the hydrolysis of chloromethyl dichloroacetate in the same conditions. In these hydrolyses not only AH* and AS* but also ACp* are highly temperature dependent, dAC,*/dT for x, = 0.980 being +38 f 2 and +67 f 5 J mol-' KW2in the hydrolyses of methyl trifluoroacetate and chloromethyl dichloroacetate, respectively.Thus the temperature dependence of ACp*is less pronouncedin the hydrolysis of methyl trifluoroacetate than in that of chloromethyl dichloroacetate. The temperature dependence of ACp*and the S-shaped Arrhenius plots in the hydrolyses studied in butoxyethanol-water mixtures can be explained on the basis of the formation of 'moving units' of clathrate hydrate- like structures when the temperature is raised. Thus with increasing temperature the solvent effects on ACp*, generally found in typically aqueous solutions, are surpassed by the effects of the transition to a microphase structure. It is well known that many solvolytic reactions in water or water-organic co-solvent mixtures do not follow the Arrhenius equation exactly but usually have a negative heat capacity of activation, AC,Z = dAH$/dT.1-3 For instance, in the neutral hydrolyses of methyl trifluoroacetate and chloromethyl dichloroacetate in water the values of ACp$ have been found to be -245 f5 and -182 f8 J rno1-I K-I, re~pectively.~ Various explanations have been given for this observation.It is usually assumed that the curvature in the Arrhenius plots can be attributed to the dependence of the related enthalpy of activation on temperat~re.l-~ Recently, it was suggested by Blandamer et al.4-8 that the two stage reaction scheme, proposed by Albery and Robin~on,~ would explain the temper- ature dependence of rate constants in various solvolyses better than the old treatment of Robertson.' Recent results obtained in this laboratory for the neutral hydrolysis of chloromethyl dichloroacetate in 2-butoxy-ethanol-water (2-BE-water) mixtures, at the mole fraction of water (x,) close to 0.98, have shown that the Arrhenius plots are S-shaped.'O This seems to be the first case when such behaviour cannot be explained by the change in mechanism with increasing temperature.In that case ACpZ is highly temperature dependent, dACp$/dT being f78 f4 and f67 f5 J mo1-I K-2 for a, 0.981 and 0.980, respectively. The positive values of ACpt found at higher temperatures are in agreement with the results of the neutral hydrolysis of p-methoxyphenyl dichloroacetate (ACpZ +937 J mol-I K-I) and 2,2-dichloropropionate (ACpS $-1 431 J mo1-I K-') in a 2-BE- water system at temperatures from 293 to 321 K when x, 0.98." According to the thermodynamic excess functions the binary system 2-BE-water belongs to a class of typically aqueous (TA) soI~tions.'~*'~The phase diagram for 2-BE-water mixtures is closed and the lower critical solution temperatures vary from 341.55 to 322.98 K when xw goes from 0.983 37 to 0.96708.'' Beyond the transition region 2-BE exists as microphases or aggregates." Further, it has been found by Ito et aZ.16 that at the 'magic' mole fraction (x, 0.98) of 2-BE-water solutions the formation of aggregates, which they call 'moving units ', is beginning.The formation of these aggregates depends considerably on temperature l5 and they have clathrate hydrate4 ke structures, 3 (H20)502-BE 16.The aim of this work was to study the behaviour of the I k2 0II R-: + CH,OH + :B I9: neutral hydrolysis of methyl trifluoroacetate in 2-BE-water mixtures when xw varies from 0.994 to 0.967 and to compare the calculated activation parameters with the data from the neutral hydrolysis of chloromethyl dichloroacetate in the different 2-BE-water mixtures studied earlier.'O The reaction is assumed to take place as a general base catalysed ester hydrolysis, BAc3, with a second water molecule acting as a general base :B.17 In the case of methyl esters of halogeno- acetic acids equation (1) represents the reaction mechanism. Experimental Methyl trifluoroacetate, a commercial product (E.Merck AG, zur Synthese), was redistilled before use. The solvent mixtures were prepared by diluting a known weight of distilled water with 2-BE (Fluka AG, purum), purified by ion exchange and redistillation, to a known volume in a volumetric flask. The initial ester concentrations were ca. 10-4~.The temperature was stable to ca. 0.01 K. The reaction was followed conductometrically as described earlier.18 Actually the concentrations of trifluoroacetic acid were measured. It has been previously shown that, at high dilutions (s 10-4~)a linear relationship between the conduct- ance and the concentration of an acid can be expected.18*19 468 J. CHEM. SOC. PERKIN TRANS. I1 1983 Table 1. Temperature range (TIK), number of data points (N), first-order rate constants (k/s-'), activation enthalpies (AH$/J mol-I),activation entropies (AS$/J mol-' K-l), and the heat capacities of activation (AC,S/J mol-' K-'), calculated from equation (2) in its five parametric form, for the neutral hydrolysis of methyl trifluoroacetate in 2-BE-water solutions with the mole fraction xwof water at 298.15 K XW T N lo3k AH$ -ASS AC,$ 0.994 0.984 273-31 3 278-3 18 11 10 7.759 6.302 38 900 f200 33 870 f120 154.9 f0.7 173.4 f0.4 -229 -623 5 29+ 34 0.980 0.982 278-321 278-321 13 12 5.106 5.512 27 940 f140 29 490 f180 195.1 f0.5 189.3 f0.6 -309 -504 f33 f43 0.973 278-321 11 3.195 26 390 f120 204.2 f0.4 +33 f29 0.967 277-3 18 11 2.324 28 240 f70 200.6 f0.2 -47 17 Table 2.Temperature range (T/K), number of data points (N), first-order rate constants (k/s-'), activation enthalpies (AHZIJ mol-I), activation entropies (AS$/J mol-' K-'), and the heat capacities of activation (AC,$/J mol-' V),calculated from equation (2) in its five parametric form, for the neutral hydrolysis of chloromethyl dichloroacetate in 2-BE-water solutions with the mole fraction x, of water at 298.15 K xw T N lo3k AH$ -ASS AC,$ 0.988 273-31 1 16 10.659 35080 5 160 165.0 f0.5 -365 f19 0.984 276-3 18 21 8.698 28 470 f250 188.9 f-0.8 -640 f54 0.981 276-31 8 20 5.804 20 650 f160 218.5 f0.5 +181 5 35 0.980 276-3 18 21 5.089 20 670 f190 219.5 f0.6 +373 f44 0.975 276-3 18 18 3.255 25 240 5 180 207.9 0.6 +614 5 40 The rate constants were calculated by Guggenheim's method.20 The standard deviations of the rate constants were in general 0.05 but were sometimes ca.0.1. The thermo- dynamic activation parameters were calculated by an extended Arrhenius equation (2), after orthogonalization, by the method Ink = A + BIT+ Cln T+ DT+ ETZ+ . .. (2) of Clarke and Glew.21 Calculations up to seven parameters show that the five-parametic equation (2) is flexible enough to represent the data correctly. Results and Discussion Kinetic data for the neutral hydrolyses of methyl trifluoro- acetate and chloromethyl dichloroacetate in the 2-BE-water solutions studied at 298.15 K are given in Tables 1 and 2, respectively. As seen from the Tables, 2-BE lowers the rates of the hydrolyses considerably in a quite narrow concen-tration range.Further, at 298.15 K AH$ and ASS are of the magnitude typical for BA,3 hydrolysis reactions,'.'' and their values decrease with decreasing water content to a minimum when the mole fraction of water is ca. 0.973 and 0.980 for the hydrolyses of methyl trifluoroacetate and chloromethyl dichloroacetate, respectively. In accord with other solvolytic reactions in aqueous organic co-solvent mixtures the changes in AH$ and AS$ counteract each ~ther.~*~~~~*~~ The logarithms of the experimental first-order rate con- stants for the neutral hydrolysis of methyl trifluoroacetate bsol;in 2-BE-water mixtures are plotted versus l/T in Figure 1. The Arrhenius plots are S-shaped when xw is 0.984, 0.982, I I 3.2 3.4 3.60.980, and 0.973.Further, in the above mole fractions, the lo3T -l/ K -llower the water content of the solution the wider is the temperature range that gives a positive value for ACp$;at Figure 1. The Arrhenius plots of -log k versus (l/T) for the neutral lower temperatures ACp$ is negative. These results are in hydrolysis of methyl trifluoroacetate in 2-butoxyethanol-water. X,agreement with the results of the neutral hydrolysis of chloro- (F) = 0.994 (A); 0.984 (B); 0.982 (C); 0.980 (D); 0.973 (E); 0.967 methyl dichloroacetate in 2-BE-water solutions when xw is from 0.984 to 0.975.1deg; Nevertheless, the S-shaped character and so the temperature dependence of ACp$(Figure 2) in the dACp$ldT for the former ester being f35 k 4 and f38 amp; 2 solvent systems studied are less pronounced in the hydrolysis J mol-' K-2 for xw0.982 and 0.980, respectively, and for the of methyl trifluoroacetate than chloromethyl dichloroacetate, latter ester +78 f4 and +67 amp; 5 J rno1-l K-2 for xw0.981 J.CHEM. SOC. PERKIN TRANS. II 1983 -2000I 1 0.99 0.98 0.97 1 0.99 0-98 X W XW Figure 2. Plots of AC,$ oersux x, for the neutral hydrolyses of methyl trifluoroacetate (a) and chloromethyl dichloroacetate (h) in 2-butoxyethanol-water mixtures. T = 278 K 0; 298 K 0; 318 K A and 0.980, respectively. The temperature dependence of AC,$ found in these hydrolyses have their origin in the fact that at the mole fractions of water close to 0.98, AH$ and AS$ go through a minimum when the temperature is raised (Figures 3 and 4).Positive values of ACp$ have usually been caused by the change in mechanism when the temperature is raised as in the neutral hydrolysis of bromomethyl chloroacetate,22 the change being from BAc3 to an SNdisplacement of bromine. 1n the neutral hydrolysis of chloromethyl dichloroacetate the same change in mechanism could be possible but is very improbable under the reaction conditions used.Io When the alkyl component is a methyl group, as in the case of methyl trifluoroacetate, the change in mechanism from BAc3 to an SNsolvolysis is impossible. Blandamer et ~tl.~-*have recently applied a hypothetical two-stage ion-pair mechanism of Albery and Robinson to explain the temperature dependence of rate constants in both SN1and SN2 solvolyses.However, Bentley and Carter 23 have seriously criticized this interpretation of AC,$ because, e.g., for the hydrolysis of ethyl bromide an ion-pair inter- mediate is firmly excluded. It is well known that a two-stage reaction scheme is evident in the neutral ester hydrolysis, but in the place of an ion-pair intermediate there is a tetrahedral intermediate equation (l). In the neutral hydrolysis of ethyl trichloroacetate in water, however, Kurz and Ehrhardt 24 have shown that the partition of the intermediate explains only a minor part (-13 J mol-' K-') of the observed AC,$ (-230 amp; 4 J mol-' K-I). Further, the two-stage mechanism cannot explain the positive values of AC,$ fouqd at higher temperatures in the hydrolyses studied because it always gives a negative ACp$.The S-shaped Arrhenius plots in the hydrolyses of methyl trifluoroacetate and chloromethyl dichloroacetate in 2-BE-water mixtures can be explained on the basis of inter- actions between ' moving units ' and both the hydrophobic character and the size of the reactant. In general, the more 60 *' c I dE 7Y bsol; 50 *m a h I 40 40 -I-0 E 2 30 3 a 20 I 10 30 50 t/OC Figure 3. Plots of AH$and TASS uerxus Tfor the neutral hydrolysis of methyl trifluoroacetate in 2-butoxyethanol-water mixtures. xW= 0.984 ; 0.982 0;0.980 A hydrophobic an organic molecule is the greater is its attraction in a hydrophobic micelle.It can be expected that this is also true in respect to microphases or aggregates. Further, it seems that when the size of an apolar ester increases the Arrhenius plots become more S-shaped and so ACp$ more dependent on temperature when xwis close to 0.98. This can be seen from the hydrolyses of the present work (Figures 1 and 2; ref. 10) and from that of p-methoxyphenyl dichloro- acetate and 2,2-dichloropropionate l1 in 2-BE-water mixtures rich in water. Consequently, in the circumstances when the formation of these aggregates has not yet begun, the values of AH$ and AS$ decrease when temperature is raised (Figures 3 and 4) because the solvation shell around the reactant in the initial state becomes looser at higher temperatures.When xwis close to 0.98 the formation of aggregates is begin- ning, Because the transition concentration for the formation of these units decreases with increasing temperature it can be expected that for the initial state the reacting ester at higher teniperatures, accommodated in ' moving units ', is in a more ordered environment than at lower temperatures. Thus AS of the hydrolyses studied first decreases and then increases when the temperature is raised. Owing to this entropy effect the rate constants at higher temperatures increase more than can be expected on the basis of the P -L I I I 10 30 50 t/OC Figure 4. Plots of AH* and TASt uersus Tfor the neutral hydrolysis of chloromethyl dichloroacetate in 2-butoxyethanol-water mixtures..Y, = 0.984 a;0.981 C'; 0.980 A Arrhenius equation. Also AH$ increases with increasing temperature. The heat capacity of activation is by far the most sensitive indicator of solvent effects. It is usually assumed that, for solvolyses in TA mixtures, ACpt has a minimum at the con- centration of water when the collapse of the solvent structure OCCU~S.~~~~~*~The position of the extrema depends on the co-solvent used and the structure of the reactant. At lower temperatures this seems to be true also in the hydrolyses of methyl trifluoroacetate and chloromethyl dichloroacetate in the 2-BE-water solutions studied (Figure 2). When the temperature is raised the minimum vanishes and at the same time moves to more water-rich regions until at higher temper- atures there is a maximum in AC,,t.This is in agreement with the formation of aggregates: with increasing temperature the normal solvent effects on AC,*, generally found in TA solutions, are surpassed by the solvent effects of the transition to a microphase structure. The present results demonstrate an extraordinary solvent J. CHEM. SOC. PERKIN TRANS. II 1983 effect on ACp*for the neutral hydrolyses of methyl trifluoro- acetate and chloromethyl dichloroacetate in 2-BE-water mixtures. However, more work is needed to better understand the solvent effects on ACptcaused by the formation of micro- phases or aggregates in solvolytic reactions. This is important because, in general, conclusions based on ACpt values cal- culated from results obtained when studying a limited temper- ature range or using different temperature intervals have been made without taking into account the possible temperature dependence of ACpt.Unfortunately, the existence of the lower critical solution temperatures in 2-BE-water solutions makes it impossible to extend kinetic runs to higher temperatures than those in this experiment. Acknowledgement 1 thank Professor E. K. Euranto for valuable discussions. References 1 R. E. Robertson, Frog. Fhys. Org. Chem., 1967,4, 213. 2 N. J. Cleve, Doctoral Thesis, University of Turku, 1973. 3 L. T.Kanerva, Licentiate Thesis, University of Turku, 1979. 4 M. J. Blandamer, R. E. Robertson, J. M. W. Scott, and A. Vrielink, J.Am. Chem. SOC.,1980,102, 2585. 5 M. J. Blandamer, J. Burgess, P. P. Duce, R. E. Robertson, and J. W. M. Scott, J. Chem. SOC.,Perkin Trans. 2, 1981, 1157. 6 M.J. Blandamer, J. Burgess, P. P. 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Euranto, in 'The Chemistry of Carboxylic Acids and Esters,' ed. S.Patai, Interscience, London, 1969, ch. 11, p. 505; E. K. Euranto and N. J. Cleve, Suomen Kem. B., 1970,43, 147; E. K. Euranto, Ann. Acud. Sci. Fennicae, Ser. A 11, 1970, No. 152. 18 L. T. Kanerva, E. K. Euranto, and N. J. Cleve, Acta Chem. Scand. B., in the press. 19 J. G. Winter and J. M. W. Scott, Can. J. Chem., 1968, 46, 2887. 20 E. A. Guggenheim, Phil. Mug., 1926, 2, 538. 21 E. C. W, Clarke and D. N. Glew, Trans. Faraduy SOC.,1966, 62, 539. 22 N. J. Cleve, Actu Chem. Scand., 1972, 26, 1326. 23 T. W. Bentley and G. E. Carter, J. Chem. SOC.,Furuduy Trans. I, 1982, 1633. 24 J. L. Kurz and G. J. Ehrhardt, J. Am. Chem. SOC.,1975, 97, 2259. Received 5th July 1982; Paper 211123