J. Org. Chem.
Thin-layer chromatography of the remaining mixture of 59 and 54 (silica gel/hexane) leads then to the isolation of a small amount of pure 59. The physical data of 59 were given previously.16 For the pyrolysis experiments, samples of about 50 mg per run were subjected to flash vacuum pyrolysis at temperatures ranging from 500 to 900 “C at a pressure of 0.1 torr. In general, the product composition determined from GLC and NMR corresponded well and the results from GLC analysis were only used as a check on the values calculated from NMR integration. Some discrepancies between GLC and NMR data could be ascribed to interconversions on the GLC column: endo- and exo-25 epimerize on the
1984, 49, 106-109 GLC column, the divinyl benzenes 31,42, and cis- and trans-49 and the 1,4-dihydronaphthalenes 38,45, and 59 appeared to be thermolabile under GLC conditions. Registry No. 21,15677-15-3; 22,8264520-3; exo-23,85803-90-3; 24, 85803-91-4; endo-25, 67504-58-9; exo-25, 67504-57-8; 42, 87729-00-8; 59, 40650-73-5. Supplementary Material Available: Tables of the results of the analyses of the pyrolysis mixtures from which the plots of Figures 1-5 have been constructed (2 pages). Ordering information is given on any current masthead page.
Substituent Effects and Mechanism in the Micellar Hydrolysis of Hydroxamic Acids’ Donald C. Berndt,* Nop Utrapiromsuk, and Douglas E. Conran Department of Chemistry, Western Michigan University, Kalamazoo, Michigan 49008 Received December 10, 1982 The rates of hydrolysis of octanohydroxamic and N-methyloctanohydroxamic acids under acidic conditions with sodium 1-dodecanesulfonate as surfactant and under alkaline conditions with cetyltrimethylammonium bromide as surfactant have been measured. Normal reaction rate orders were obtained except for the alkaline hydrolysis of octanohydroxamic acid above the critical micelle concentration of the surfactant. The latter yielded the novel result of pseudo-zero-order kinetics. The rates of acidic hydrolysis of a series of meta- and para-substituted benzohydroxamic acids in aqueous solution with sodium 1-dodecanesulfonate as surfactant were also measured. The substituent effects indicate specific micellar influences on the rates and a difference in mechanism between the bulk aqueous phase and the micellar phase.
Micellar enhancement of reaction rates of bimolecular reactions could result from one or both of two factors: the concentration of reactants by the micellar phase and the relative stabilization of transition states and/or destabilization of reactant states by the micellar environment relative to the bulk aqueous phase.,v3 We report in this paper evidence that the latter factor has a significant role in the micellar catalysis of hydroxamic acid hydrolysis.
Results and Discussion Equation 1, derived by Romsted: is applicable to bimolecular reactions in which one reactant is an ion. The k2 = k,’ PSKa(C,- cmc) k,’ (1) Ka(C, cmc) + 1 [Ka(C,- cmc) + l][I,+ XtKI] overall second-order rate constant and the rate constants in the aqueous and micellar phases are k,, k,’, and k,‘, respectively, is the degree of counterion binding to the Stern layer, S the molar density of the micellar phase, K , the association constant of the organic substrate with the micellar phase, C, the total surfactant concentration, cmc the critical micelle concentration, Itthe total concentration of the hydrophilic reactant ion, X , the total concentration of the surfactant counterion, and K I is the ion exchange constant given by eq 2 in which the subscripts w and m refer to the aqueous and micellar phases, respectively.
IuXm KI = ImX,
(1)Abstracted from the Ph.D. Dissertation of N.U., Westem Michigan University, 1982,and the M.A. Thesis of D.C., Western Michigan University, 1983. (2) Bunton, C. A. Tech. Chem. (N.Y.) 1976, 20,731. (3) Cordes, E. H. Pure Appl. Chem. 1978,50, 717. (4)Romsted, L. S. In ‘Micellization, Solubilization and Microemulsions”; Mittal, K. L., Ed.; Plenum Press: New York, 1977;Vol. 2,p 509.
Table I. Kinetic Data for Acidic Hydrolysis in 0.09279 N HCI at 50.0 i 0.11 “C as a Function of Sodium 1-Dodecanesulfonate Concentration octanohydroxamic N-methyloctanohydroxamic 103C, acid, 1 0 5 k 0 b d ,s-’ acid, 1 0 5 k 0 b d ,s-’ 0.0 0.060 0.485 3.01 4.996 7.996 9.990 11.99 15.00 20.40 30.01 40.00 60.07
2.07 1.99 2.06 5.69 11.5 17.4 21.9 23.3 26.2 29.6 32.5 34.7 34.2
4.94 4.85 4.99 12.5 22.0 31.0 34.4 37.1 39.2 42.6 44.4 44.3 44.2
The constants of eq 1 (with k,’PS considered as one constant) may be estimated as follows. In the absence of added common ions, X , = Ct.In the range (C, - cmc) N C,, eq 1 may be treated after the manner of Martinek et al.5 to yield eq 3,
in which kobsd = kzI,, k , = k,’It, and p , q, and r are constants which are functions of the constants in eq 1. A t sufficiently low enough values of C,, the C,2 term is negligible and a graph of the left side of eq 3 vs. will yield values for p and r. A t the rate maximum eq 4 (5)Martinek, K.;Yataimirski, A. K.; Levashov, A. V.; Berezin, I. V. In “Micelliiation, Solubilization and Microemulsions”; Mittal, K. L., Ed.; Plenum Press: New York, 1977;Vol. 2, p 489.
0022-3263/84/1949-0106$01.50/00 1984 American Chemical Society
J. Org. Chem., Vol. 49, No. 1, 1984 107
Micellar Hydrolysis of Hydroxamic Acids
Table 11. Kinetic Data for Alkaline Hydrolysis in 0.1111 N NaOH at 50.0 f 0.11 "C as a Function of Cetyltrimethylammonium Bromide Concentration octanohydroxamic 0
I I 10
102 C t
Figure 1. Experimental (circles) and calculated (squares) fmt-orderrate constants w. C, for octanohydroxamic acid (closed
symbols) and N-methylhydroxamicacid (open symbols).
0.0 0.27 1.5 4.92 10.0 12.00 30.04 50.01 201.9 399.9 600.1
(1.96)a (2.20)a (2.78)a 0.731 1.06 1.28 2.14 2.53 2.89 3.15 3.16
20.1 50.05 200.1 399.8
N-methyloctanohydroxamic acid, 1O6hobd,s-'
1.84 1.84 1.87 2.80 3.67 4.19 4.83 7.80 8.08 6.92 7.07
Pseudo first order, 1O6h,bd, s-'.
' NaOH, 0.04539 M.
results, which may be used with eq 3 to evaluate the constants of eq 1. 1 - 2qkzmaxCtmax - pk2" =0 (4) Kinetic data for the acidic hydrolysis of octanohydroxamic and N-methyloctanohydroxamic acids is listed in Table I. Figure 1shows the observed and calculated k o b d for these compounds. The calculated kobsdvalues were obtained from the constants estimated by the above procedure. Values for k,'PS, K,, and KI are 4.55 X lo4, 86.4, and 0.319 and 5.21 X lo", 190, and 0.247 for octanhydroxamic acid and N-methyloctanohydroxamic acid hydrolyses, respectively. The cmc under the conditions of M by the rate measurements was evaluated as 2 X graphical interpolation of the observed rate constant vs. surfactant concentration for acidic hydrolysis of the octanohydroxamic acids (Table I) and p-ethylbenzohydroxamic acid (Table 111). The results show satisfactory agreement between calculated and observed values. k,' was evaluated in the absence of surfactant and no corrections for changes in ionic strength have been made (ionic strength effeds are expected to be small6). The two values for KI are in reasonable agreement. An unusual result was obtained for the alkaline hydrolysis of octanohydroxamic and N-methyloctanohydroxamic acids (Table 11) in the presence of cetyltrimethylammonium bromide. The N-methyl derivative exhibited the usual pattern of pseudo-first-order rate behavior; however, octanohydroxamic acid hydrolysis above the cmc displayed pseudo-zero-order behavior. The cmc was determined to be 2 X lo4 M as described above. The same rate order pattern was observed a t two additional base strengths as is shown in Table 11. Bimolecular eq 1reduces to the unimolecular eq 5 when the factor (I, X&I) N It with X,= Ct.This will be the
(5) case for the data in Table I and in Table I11 (acidic hydrolysis of substituted benzohydroxamic acids) with It = [H+] and K,,as determined above, in the C, range below the rate maximum. k, and K I N in eq 5 correspond to (6) Berndt, D. C.; Sendelbach, L. E.
acid, 10'%obad, mol L-'s-
J. Org. Chem. 1977, 42, 3305.
NaOH, 0.1829 M.
k,'@S and K, in eq 1,respectively. The k, and K I N values determined from eq 5 (used in the appropriate Ct range) will be more accurate than the corresponding values from eq 1 and are the quantities of interest to determine substrate structural effects in micellar catalysis. Equation 5 is used in the analysis of substituent effects which follows. The accepted mechanism7-10of acidic hydrolysis of hydroxamic acids is as shown below. We have determined RC(=O)NHOH
+ H30+ == RC+(OH)NHOH + HzO
substituent effects on the hydrolysis of meta- and parasubstituted benzohydroxamic acids in the presence and absence of the surfactant sodium 1-dodecanesulfonate. These effects indicate that a change occurs ill the relative importance of the first and second steps in the mechanism when the surfactant is added to the reaction system. The kinetic data is given in Table 111. The data for all compounds, except for the m-nitro derivative, yield a good fit to eq 5. K I N and k, values were obtained by least squares analysis of the linear relationships between l / ( k w - kobsd) and 1/c,- cmc) from eq 5. The results are in Table IV. The fit of eq 5 is excellent in all cases (correlation coefficients 0.997-0.999) except for the m-nitro derivative in which case a linear relationship between l / ( k w- kobsd)and l/(Ct - cmc) was not obtained. Rate data of previous studies6-" of hydroxamic acid hydrolysis with surfactants were correlated by eq 5. Equation 6 shows the parameters obtained for the correlation of rate constants in the absence of surfactant with the Hammett substituent parameter,12 a, for all ten compounds in Table I11 (the correlation coeficient is 0.8879 and the F-test13shows significance within the 0.1% level; (7) Berndt, D.C.; Fuller, R. L. J. Org. Chem. 1966, 31, 3312. (8) Tillet, J. G.;Hudson, K.; Buglass, A. J. J.Chem. SOC.E 1971,123. (9) Berndt, D.C.; Sharp, J. K. J. Org. Chem. 1973, 38, 396. (IO) Berndt, D.C.; Ward, I. E. J. Org. Chem. 1974, 39,841. (11) Berndt, D.C.;Utrapiromsuk, N.; Jaglan, S.S.J.Org. Chem. 1979, 44, 136. (12) Exner, 0.In "Advances in Linear Free Energy Relationships"; Chapman, N. B., Shorter, J., Eds.; Plenum Press: New York, 1972; Chapter 1. (13) Leabo, D.A. 'Basic Statistics", 4th ed.; Richard D. Irwin, Inc.: Homewood, IL, 1972; Chapter 16.
J. Org. Chem., Vol. 49, No. I, 1984
Berndt, Utrapiromsuk, and Conran
Table 111. Kinetic Data for Substituted Benzohydroxamic Acid Hydrolyses in HCl Solutiona a t 67.9-68.2 "C as a Function of Sodium 1-Dodecanesulfonate Concentration 105kobsd, 5 - l by substituent Ctb,M
0 0.001 0.01 0.03 0.05 0.07 0.10
1.51 1.52 1.83 2.43 3.12 3.30 3.61
1.57 1.60 2.37 3.49 4.80 4.99 5.76
m-CH, p-CH,CH,' 1.73 1.45 2.38 3.86 5.83 6.38 6.36
1.34 1.30 1.72 2.35 2.91 3.06 3.36
1.32d 1.40 1.71 2.88 3.42 4.17 4.20
1.01 0.99 1.27 1.89 2.18 2.45 3.05
0.8Bd 0.90 1.37 1.5Bd 1.68 1.86 2.62
1.22 1.18 2.24 4.04 5.47 6.07 6.41
1.14 1.02 2.29 3.62 4.79d 5.41 5.73d
1.62 1.46 3.60 5.84 7.45 7.89 7.80
Surfactant concentration. ' Additional values for In 0.160 M hydrochloric acid solution, unless otherwise indicated. In 0.158 M hydrochloric p-CH,CH, are 1.74 and 2.83 at surfactant concentrations of 0.003 and 0.007 M, respectively. acid solution. Table IV. Correlation of t h e Rate Data by Equation 5 in C+ Range 0.01-0.10 M 105kw,
Table V. Correlation of log 12, b y Equation 7
H P-CH, m-CH, p-CH,CH, p-OCH, m-OCH, P-NO, p-c1 m-C1
1.51 1.57 1.73 1.62 1.34 1.32 1.01 1.22 1.14
9.2 18.2 3.8 39.0 16.0 29.1 7.2 15.1 28.7
6.20 7.84 23.75 9.97 4.71 18.72 5.78 10.68 7.23
4.1 5.0 14.0 6.2 3.5 14.0 5.7 8.8 6.3
-0.999 -0.998 -0.998 -0.998 -0.999 -0.997 -0.999 -0.999 -0.997
(14)Hansch, C.; Leo,A.; Unger, S.H.;Kim,K.H.; Nikaitani, D.; Lien, E.J. J.Med. Chem. 1973, 16,1207.
P-H P-CH, p-CH,CH, p-OCH, P-NO,
-4.207 -4.106 -4.001 -4.327 -4.238 -3.972
-4.227 -4.103 -3.970 -4.324 -4.271 -4.011
From values in Table IV. Table VI. Correlation of log ( K I N ) b y Equation 8
the calculated value, -4.850, compares well to the measured value, -4.829, for the reference substituent, hydrogen). log k , = - 0 . 2 2 7 ~- 4.850 (6)
constant for hydrogen as substituent. Multiple regression
a Correlation coefficient.
The negative value of p (-0.227) for reaction in dilute acid solution is consistent with previously reported studies of meta- and para-substituted phenylacetohydroxamic" acids (substituent parameter, go, for "insulated" systems), aliphaticg hydroxamic acids (substituent parameter, u*), and ortho-substituted benzohydroxamic10acids (substituent parameter, The negative value of p for the observed rate indicates that the substituents have a larger influence on the first step than on the second step in the mechanism in the absence of surfactant. In our earlier study," a correlation of log ( K I N ) with the lipophilicity constant, P,was reported for meta- and para-substituted phenylacetohydroxamic acid hydrolysis in dilute hydrochloric acid solution with sodium l-dodecanesulfonate as surfactant. Since the micellar aggregation number, N , presumably is constant under constant experimental conditions for the series of substrates, the correlation indicated that the association constants, K , reflect the perturbations on solubilities produced by substituents in the micellar phase relative to the aqueous phase. The group of substituents in the earlier study did not include any with oxygen or nitrogen atoms. In the present work, satisfactory correlations of log (KIN) or log k , do not result if both meta and para substituents are included; however, meaningful correlations do exist for the para series. Equation 7 assumes that lipophilicity and polar effects as represented by the 7r14 and u+12 parameters are separable. a is a susceptibility constant and kmH is the rate log k , = pu+ + CYK + log kmH (7)
P-H P-CH, p-CH,CH, p-OCH, P-NO, p-c1
0.964 1.259 1.591 1.204 0.860 1.180
1.037 1.304 1.479 1.174 0.787 1.283
From values in Table IV. Table VII. Substituted Benzohydroxamic Acids mp, "Ca substituent
P-CH, m-CH, p-CH,CH, p-OCH,
140 95 99 160
obsd 142.5-143 97-98dec 106.5-107 160-160.5
lit. 141-143 161-162 160.5-162'
m-OCH, 70 76.8-77 P-NO, 170 178-179 dec 176.2-177d m-NO, 140 143-144 145e p-c1 175 182-183 182-184 m-C1 170 177-178 dec 169-171' Insert temperature is the temperature of melting point Reference 8. bath when sample is placed into bath. Buraczewski, K.; Czerwinska, E.; Eckstein, Z.; Grochowski, E.; Kowalik, R. ; Plenkiewicz. Bull. Acad. Pol. Sei., Ser. Sei. Chim. 1964, 12, 773. Kornblum, N.; Brown, R. A. J. Am. Chem. SOC.1 9 6 5 , 8 7 , 1742. e Exner, 0.; Simon, W. Collect. Czech. Chem. Commun. 1 9 6 5 , 3 0 , 4078.
analysis (the correlation coefficient is 0.982 and the F-test shows significance a t the 1%level) yields p = 0.117, a = 0.375, and log kmH = -4.228. The last value compares well with the value -4.207 determined by eq 5. The values of log k , from eq 5 and those predicted by eq 7 are in Table V. Equation 7 is for reaction within the micelle while eq 6 is for reaction in the bulk aqueous phase. They differ in three significant ways: First, the best correlation in the micellar phase is obtained with U+ rather than with U. Second, the sign of p is reversed in the micellar phase. Third, the lipophilicity of the substituents has a significant
J. Org. Chem., Vol. 49, No. 1, 1984 109
Micellar Hydrolysis of Hydroxamic Acids hydroxamic acids m-methvlbenzop-ethylbenzom-methoxybenzoN-methyloctano-
Table VIII. Elemental Analysis of New Hydroxamic Acids analysis %C %H %N 6.03 9.49 obsd 63.58 calcd 63.60 5.96 9.27 obsd calcd obsd calcd obsd calcd
65.47 57.29 57.51
role in the total substituent effect in the micellar phase. Comparison of pa+ with (YT shows that the polar effect (pa+) predominates only with the oxygen- and nitrogencontaining substituents. These differences indicate specific micellar effects upon the reaction and that a change in the relative importance of the first and second steps in the mechanism occurs when reaction takes place in the micellar phase compared to the aqueous phase. In a related system a difference in the mechanism of alkaline hydrolysis of N,N-disubstituted amides has been reported for reaction in the micellar phase as compared to the bulk phase.15J6 The lipophilicity parameters measure substituent effects on the relative solubility in octanol compared to water and therefore reflect various solvation factors such as dipole interactions, London dispersion force interactions, and hydrogen bonding in different degrees of importance with different substituents. The presence of the (YTterm in eq 7 indicates the importance of these solvation effects as part of the substituent effect in the micellar microenvironment. Equation 8 gives the best correlation for log (K N); the correlation coefficient is 0.941, and the F-testL shows significance a t the 5% level. Calculated values are p = log (K/N) = pu+
+ cia + log (KH/N)
-0.185, ci = 0.375, and log(KH/N) = 1.037, the value for hydrogen as substituent. Table VI compares log (KIN) values predicted by eq 8 to those from eq 5. Comparison of the pa+ and cia terms for eq 8 shows that the u+ contribution to log (KIN) is only predominant for the oxygen-containing substituents. These results are consistent with our earlier observation" of a correlation of log (KIN) with only the T parameter in the substituted phenylacetohydroxamic acid series in which there were no oxygenor nitrogen-containing substituents and in which there is a saturated carbon between the substituted phenyl group and the reaction center. The positive correlation of log (KIN) with the lipophilicity parameter is consistent with (15)Broxton, T.J.; Fernando, D. R.; Rowe, J. E. J. Org. Chem. 1981, 46, 3522. (16) Broxton, T. J.; Duddy, N. W. Aust. J. Chem. 1979, 32, 1717.
8.42 8.48 8.54
8.38 8.32 8.08
formula C,H,NOZ C9H
the lower polarity of the micellar environment? compared to that of the aqueous phase. Separate correlations of log k, and log ( K I N ) for the meta series similar to those of the para series might be possible. Solubility problems prevented the study of a sufficient number of compounds to establish statistical reliability. Different correlations for the meta and para series could arise from different micellar microenvironments for the two series of compounds. The large difference in the KIN values (Table IV) between the meta and para compounds for the same substituent is evidence for this possibility.
Experimental Section Octanohydroxamic acid was described previously.6 NMethyloctanohydroxamic acid was prepared by a previous method" and crystallized from a 2:l water-ethanol mixture in a freezer, mp 15.0-17.3 "C. Table VI1 lists substituted benzohydroxamic acids prepared in this investigation. (Benzohydroxamic acid was described previously.7) Previous methods were used for their preparation." They were crystallized from water, methanol, ethanol, or aqueous ethanol. The IR and 'H N M R spectra were consistent with the indicated structures. New compounds have satisfadoryC, H, N analyses (MidwestMicrolab, Ltd.,Indianapolis, IN, or Galbraith Laboratories, Inc., Knoxville, TN) . Cetyltrimethylammoniumbromide was purified by the method of Bunton and Wolfe.l* Sodium 1-dodecanesulfonatewas prepared" as reported previously and the kinetic measurements were obtained by the method used previously." Registry No. Octanohydroxamic acid, 7377-03-9; N methyloctanohydroxamic acid, 65789-45-9;benzohydroxamicacid, 495-18-1;p-methylbenzohydroxamicacid, 2318-82-3;m-methylbenzohydroxamicacid, 10335-81-6;p-ethylbenzohydroxamicacid, 87828-95-3;p-methoxybenzohydroxamic acid, 10507-69-4; mmethoxybenzohydroxamic acid, 31791-98-7; p-nitrobenzohydroxamic acid, 1613-76-9; m-nitrobenzohydroxamic acid, 7335-34-4;p-chlorobenzohydroxamicacid, 1613-88-3;m-chlorobenzohydroxamic acid, 4070-53-5;sodium 1-dodecanesulfonate, 13419-61-9;cetyltrimethylammoniumbromide, 57-09-0. (17) Berndt, D. C.; Ward, I. E. J. Org. Chem. 1976,41, 3297. (18)Bunton, C.A.;Wolfe, B. J. Am. Chem. SOC.1973, 95,3742.