Deuterium Isotope Effects in Complexation Kinetics
1275
Deuterium Isotope Effects in Complexation Kinetics. I I. Lanthanide( I I I) Sulfate Systems' Jeffry Reidler and Herbert
B. Silber"
Department of Chemistry, University of Maryland Baltimore County, Baltimore, Maryland 27228 (ReceivedJuly 27, 7972)
The complexation reaction between La(III), Nd(III), Gd(III), and Dy(II1) with so42-was investigated in HzO and D2O at 25". In all cases the complexation rate constant, k34, and the dissociation rate constant, k43, for each cation were different in the two solvents. The values of 10-%34 in HzO and DzO are La 0.84 and 0.24 sec-1, Nd 1.9 and 0.80 sec-l, Gd 8.2 and 6.9 sec-1, and Dy 0.91 and 1.1 sec-I; and 10-'k43 equals La 3.8 and 7.0 sec-1, Nd 4.8 and 7.1 sec-I, Gd 3.8 and 0.74 sec-1, and Dy 4.5 and 2.1 sec-l. These data were compared to other lanthanide deuterium isotope studies to establish that the isotope effect is not a function of the solvated cation alone, but rather of the cation-ligand complex.
Introduction Systematic kinetic studies of lanthanide complexation reactions have attempted to probe the nature of the solvated lanthanide ion in terms of the Eigen-Diebler multistep complexation reaction2,3 M"+(aq)
-+
4 2
L"-(aq)
h2 1
M"+(H,o),L~-
k34 ==
ML"-"(aq)
(1)
k43
step 12
s t e p I11
The ultrasonic absorption measurements were carried out at selected frequencies between 10 and 270 MHz, using standard techniques.9 The apparatus consisted of Matec Models 666 pulse comparator, 120 master synchronizer, 600 RF pulse generator and receiver, and 960 RF source. X-Cut quartz crystals with fundamental frequencies of 4, 5 , and 10 MHz were obtained from Valpey-Fisher. The experimental vibration frequencies were measured using a Gertsch FM3 frequency meter in conjunction with a Heath RF signal, generator. A double-wall type interferometer cell was thermostated a t 25" by a Forma circulating temperature bath. Sound velocities were not directly measured and were assumed to be that of the pure water value. The pH values of the solutions were approximately 2 or above, because solutions of high acidity reacted with the stainless steel cell, and to prevent hydrolysis of the lanthanides. The pH was measured with a Corning Digital 112 pH meter in conjunction with a semimicro Corning combination electrode and standard buffers. All calculations were carried out on a Wang 700 programmable calculator.
Although for many cations the slow process is step I11 involving slow cation-solvent exchange, the analysis of neodymium(II1) sulfate complexation in H20 and D2O indicated that the cation-ligand bond formation step was in fact the rate-determining step.l This study was undertaken for two reasons. First, to extend the results to other lanthanide(II1) sulfate systems to determine if a turnover point exists in the lanthanide series where cation solvent exchange becomes rate determining. Second, to compare a n additional ligand series of deuterium isotope effects Results with the previously determined ones using anthranilate ion and murexide as ligands4 to discover if the deuterium An excess ultrasonic absorption was measured over the effects for each cation are independent of the ligand. whole frequency range for all four lanthanide systems. For Although the lanthanide(II1)-sulfate systems have been chemical relaxation, the ultrasonic absorption, alp, is independently investigated by Purdie and coworkers~-7 given by and by Atkinson and Grecsek,S the calculated rate con+ B stants were significantly different. It was therefore neces2 f ' . = 1 (flfRJ2 sary to repeat their measurements in H2O to minimize errors in comparisons with the DzO results. The lanthanides where Ai = relaxation amplitude, f = frequency, fRi = reinvestigated include La, Gd, and Dy, as well as the prelaxation frequency, and B = background. Although at viously determined Nd. high frequencies the absorption did not come down to the
+
Experimental Section The lanthanide sulfate solutions were made by the addition of sulfuric acid to the oxides (Research Inorganics), which were a t least 99.9% pure. Standardized HzS04 (Baker analyzed reagent grade) or DzS04 (Diaprep, 99.5% isotopic purity) was added and heat was applied to form the aqueous solutions. The D2O was at least 99.8% isotopically pure (Biorad and Diaprep). The residual H20 in the D2O lanthanide solutions was less than 0.5%, as determined by nnir analysis.
(1) Part I: H. B. Silber, Chem. Commun., 731 (1971). (2) H. Diebler and M. Eigen, Z. Phys. Chem. (Frankfurt am Main), 20, 200 (1959). (3) M. Eigen and K.Tamm, Z.Electrochem., 66,93,107 (1969). (4) H. B. Silber, R. D. Farina, and J. H. Swinehart. lnorg. Chem., 8, 819 (1969). (5) N. Purdie and C. A. Vincent, Trans. Faraday Soc., 63,2745 (1967). (6) D. P. Fay, D. Litchinsky, and N. Purdie, J. Phys. Chem., 73, 544 (1969). (7) D. P. Fay and N. Purdie, J. Phys. Chem., 74,1160 (1970).
(8)J. J. Grecsek, Ph.D. Thesis, University of Maryland, College Park,
1970. (9) H . 8.Silber, N. Seheinin, G. Atkinson, and J. J . Grecsek, J. Chem. SOC., Faraday Trans. 7, 68 1200 (1972). The Journal of Physical Chemistry, Vol. 77,No. 10, 7973
Jeffry Reidler and Herbert B. Silber
1276
A. H 2 0 (K, = 4200,K12 = 1300)
0.700 I .40 1.98 2.97 3.96 4.95 7.00
0.654 1.31 3.03 4.55 6.06 7.58 6.54
2.990 2.746 2.302 2.165 2.078 2.007 2.034
1.07 1.16 1.21 1.32 1.46 1.63 1.96
1.02 2.56 3.20 4.00 4.01 5.02 5.01
1.69 4.22 5.27 6.59 6.80 1.93 8.50
3.29 3.04 2.96 2.89 2.75 2.62 2.65
1 .os 1.31 1.47 1.70 1.69 1.99 2.01
0.502 0.522 0.533 0.554 0.579 0.606 0.649
178 f 5 315 f 5 423 f 10 543 f 10 706 f 10 833 f 10 1214 f 12
13.4 f 0.4 14.2 f 0.2 14.3 f 0.1 14.6 f 0.3 14.9 f 0.3 15.0 f 0.3 15.5 f 0.3
23.5 f 0.4 23.7 f 0.4 26.6 f 0.3 28.0 f 2.1 31.3 f 3.0 36.7f 3.7 45.2 f 6.5
283 f 7 637 f 6 865 f 7 1025 f 2 958 f 13 1180 f 17 1227 f 13
13.2 f 0.3 14.1 f 0.2 13.1 f 0.5 13.7f 0.1 14.3 f 0.3 14.1 f 0.2 13.6 f 0.2
31.4 f 0.5 36.3 f 1.1 41.7 f 1.7 40.4 f 0.9 37.5 f 2.6 37.4 f 1.5 43.3f 1.7
B. D20
a
0.586 0.639 0.665 0.696 0.695 0.729 0.730
Units = nepers cm-’ sec2.
TABLE II: Gadolinium Sulfate Data
A. H 2 0 (K, = 4500,K i 2 = 200)
0.361 0.722 1.08 1.49 1.19 1.80 2.24 1.79 2.99 3.73
3.88 7.76 11.6 42.9 69.0 19.4 64.3 104 173 107
3.456 3.266 3.123 2.215 2.407 2.914 2.318 2.307 1.930 1.902
0.977 1 .I2 1 .I8 1 .I8 1.24 1.25 1.30 1.32 1.41 1.42
0.129 0.360 0.721 0.901 0.388 1.80 1.29
16.6 4.68 9.36 11.7 49.8 23.4 166
4.84 3.83 3.30 3.16 4.38 2.95 3.69
0.906 1.02 1.24 1.29 1.33 1.43 1.61
0.164 0.183 0.191 0.191 0.200 0.201 0.207 0.209 0.221 0.222
55.3 f 2.0 91.3 f 0.9 129 f 1 I58 f 2 131 f 2 190 f 2 232 f 5 176 f 1 236 f 2 333 f 3
27.0f 1.5 30.1 f 0.5 31.1 f 0.6 32.0 f 0.7 31.8 f 0.7 33.8f 0.6 32.0 f 1.0 35.0 f 0.4 35.3 4 0.7 33.9 f 0.6
21.9 f 0.6 22.6 f 0.3 23.9 f 0.6 25.7 f 0.7 24.7 f 0.6 24.8 f 0.7 29.6 f 1.5 24.4 f 0.5 29.4 f 2.6 30.7 f 1.7
48.4 Zt 23 760 f 2.4 I33 f 2 I47 f 3 91.0 f 1.6 245 f 2 237 f 2
16.2 f 0.9 20.7 f 1.1 22.6 f 0.7 22.1 f 0.9 24.0f 0.6 28.0f 0.5 27.9 f 0.5
29.5 f 0.3 31.5 It 1.1 34.2 f 1.2 42.6 k 3.0 29.8 f 0.4 37.4 f 1.8 34.3 f 1.5
B. D20
0.154 0.170 0.199 0.206 0.210 0.224 0.244
a Units = nepers c m - l sec2.
value expected if only a single relaxation was present, the amplitude of this high-frequency relaxation was usually too low to obtain useful relaxation information, consistent with the results of other workers. The HzO background value used was B = 23.0 X 10-17 nepers cm sec2 and deviations from this are attributed to the faster relaxation. The relaxation parameters obtained as described previouslys are summarized in Tables 1-111. The identification of the low-frequency peak with step I11 of reaction 1 has previously been made with nitrate,g and acetate ligands.10 The excess absorbance deviation from the background at high frequency indicates the existence of a more rapid step which is identified with step 12, especially since 0.05 M Nd2(SO& is known to have a measured B of (22.4 f 1.1) X 1 O - l 7 nepers cm-l sec2 from high-frequency measurements.ll For reaction 1 under these experimental conditions, Tamm has shown The Journalof Physical Chemistry, Vol. 77, No. 10, 1973
that the relaxation time expressions for steps 12 and 111 are12 (3)
(4)
A scheme similar to that of the nitrate system,s modified to include the presence of H S O c , was used to evaluate 8(c) and $ ( e ) . (IO) V. L. Garza and N. Purdie, J. Phys. Chem., 74, 275 (1970). (11) K. Fritsch, C. J. Montrose, J. L. Hunter, and J. F. Dill, J. Chem. Phys., 52, 2242 (1970). (12) K. Tamm, “Dispersion and Absorption of Sound by Molecular Processes,” D. Sette, Ed., Academic Press, New York, N, Y., 1963,
pp 175-222. Tamm uses dissociation constants in his expression, thus the equilibrium constants used in this paper are reciprocals of Tamm’s.
1277
Deuterium Isotope Effects in Complexation Kinetics TABLE I l l :
Dysprosium Sulfate Data
A. H20 0.250 1.oo 2.50 3.50 5.01
0.287 1.15 2.87 4.65 6.64
3.593 3.159 3.033 2.840 2.771
0.867 1.13 1.29 I .43 1.70
0.465 1.12 1.86 4.65
0.452 1.09 1.81 4.52
4.39 3.83 3.53 3.29
1.01 1.41 1.53 1.76
(K, = 4100, K12 = 1350) 0.541 0.605 0.637 0.660 0.69.8
72 227 513 650 832
f4
90 177 287 561
f9 f6
f5 f9 f8 f8
15.2 16.0 16.1 16.8 17.6
f 0.9
f 0.4 f 0.3 f 0.2 f 0.2
22.0 25.1 28.8 29.0 32.2
f 0.4 f 0.6 f 1.1 i 1 .O f 0.9
29.0 30.1 33.9 34.3
f 0.7 f 0.4 f 0.6 f 0.8
B. D20
a
0.577 0.655 0.674 0.703
f7 f7
13.4 f 14.3 f 14.0 f 16.0 f
1.4 0.5 0.4 0.2
Units = nepers cm-’ sec2,
The initial concentration of L n z ( S 0 4 ) ~is c and the excess sulfate (HzSO4 + NazS04) is d, and the association constant for LnS04+ complexation, Kc, is 1
Kc=?
=
--
(5)
By substitution, eq 6 becomes
a) + d + c(l1++ -2CH+l (1
d In n~ 8 In [Ln3+])[SO,?
K,
>)(7)
and
+ (4+ -
(1+ A small computer program was written to calculate ry, I, and 6 ( c ) for each solution. In HzO, the bisulfate p K a was assumed to be 1.9914 while in DzO pKa equals 2.34.15 The pD was obtained relative to H20 buffers by16 CY,
pD = p H The quantity d(c) is given by12
+
0.40
-----L I1
F= LnSO,+(aq)
(11)
where (12) K,, = KI(1 + KII) Step I is diffusion controlled and KI can be calculated from the Bjerrum equation. However, KII is not diffusion controlled and the magnitude is a function of the relative stabilities of two different outer-sphere complexes. Hence, in this paper Klz was not assumed but was calculated from Kc and the rate data. T o calculate the best Klz consistent with the relaxation data, eq 4 and 12 were combined and rearranged to obtain 7111-1
+
Ln3+(H,0),S0,2111
0(c) =
*,{2nc
I
Ln3+(H20),S042-
- I y
where r y is the activity coefficient ratio, CY is the degree of dissociation of LnS04+, and Ka is the bisulfate dissociation constant. The Davies equation was used to determine r y . 1 3 For this system 6 ( c ) is given by
e(c) =
Ln3+(aq) + S04*-(aq)
*
(9)
Equation 10 shows that d(c), and hence 1234 and k43, are sensitive to the choice of Klz. In terms of an Eigen-type mechanism, step 12 involves the diffusion together of a solvated cation with a solvated anion, followed by water loss from the anion solvation shell. In terms of the threestep sequence one has
=
7III-’ was plotted as a function of X with different values of KIII assumed. The KIII value which gives a zero intercept results in a n internally consistent K12 and KIII.This technique was utilized to obtain the step I11 rate constants in HzO. In DzO, the NdS04+ high-frequency absorption data were essentially the same as in HzO and thus Klz was assumed to be the same in both solvents.1 With the other lanthanides the high-frequency absorption data were slightly higher in D20 than HzO, which was attributed to the higher background found in DzO in the absence of salts. Hence Klz was also assumed to be the same in both solvents. A cyclic calculation was used to obtain the rate constants in which Kc was assumed, 6 ( c ) and d(c) calculated, least-squares analysis to obtain the step I11 rate constants, and KIIIwas carried out and a new Kc was calculated uia Kc = K12(1 + KIII).This process was continued until Kc assumed and calculated agreed to with 0.1%. The final results for the four lanthanides are shown in Table IV, and Figure 1 represents the kinetic data for Gd.
(13) C. W. Davies, “Ion Association,” Butterworth, Washington, D. C.. 1962, p 46. (14) A. A. Noyes and M. A. Stewart, J. Amer. Chem. SOC., 32, 1133 (1910): M. S. Sherrili and A. A. Noyes, ibid., 48, 1861 (1926); M. H. Lietzke and R. W. Stoughton. J . Phys. Chem., 65, 2247 (1961). (15) M. H. Lietzke and R. W. Stoughton. J. Phys. Chem., 67, 652 (1963). (16) R. G. Bates, “Determination of pH Theory and Practice.” Wlley, New York, N. Y., 1964, pp 219-220. The Journal of Physical Chemistry, Vol. 77, No. 10, 1973
Jeffry Reidler and Herbert B. Silber
1278
TABLE IV: Lanthanide Sulfate Kinetic Results at 250a
La
4200 4370 4500 4100
Nd Gd DY
0.84 dz 0.01 1.9 0.2 8.2 f 1.2 0.91 f 0.10
1300 900 200 1350
0.24 f 0.01 0.80 f 0.56 6.9 f 1.7 1.1 f 0.5
3.8 f 0.69 4.8 dz 1.0 3.8 f 2.4 4.5 f 0.91
*
7.0 f 1.5 7.1 f 3.3 0.74 f 3.6 2.1 f 3.5
‘The same value of K;z was used in both solvent systems. Kc in DzO can be calculated from K ~and Z the reported ks4 and k43results.
TABLE V: Effect of Alternate Calculation Procedures upon the Experimental Results Ratio of H20 to D20 rate constants 1O-8k34, sec -1 Lanthanide
KC
10 -7k43,sec-1
Ki2
k34’
k34’
k43’
k43’
0.57 8 28
5
~
A. Kc and
La
4200 4500 41 00
Gd DY
K12
1300 200 1350
Remain Unchanged as D20 is Substituted for H 2 0
0.35 7.8 1.4
KcC
3.8 0.089 0.16
2.4 1.1 0.65
sec-’
10-7k43, sec-1
B. Effect of K, on the Gd ( I I I ) 4,500 20,000 19,000 18,900
3.5 1.2 0.83
7.80 6.91 6.93 6.93
0.54 2 Kca
Rate Constants in D20 0.089 0.762 0.745 0.743
177,000 18,400 18,900 18,900
C. Effect of the Choice of K1? on the Gd ( I I I ) Rate Constants 4500 4500 a
200 1180
8.2 f 1.2 5.4 0.8
3.8 f 2.4 -1.2 f 0.4
*
7.8 f 2.3 4.3 f 0.4
0.089 f 4.5 -1.0 f 0.4
1.1 1.3
’
Calculated. From Table IV. Assumed.
0.
I
I.
The Journal of Physical Chemistry, Vol. 77, No. 10, 1973
I
18,900 but the range due mainly to errors in k 4 3 is from 2600 to infinity. The question must be asked if the kinetic differences in the two solvents are artifacts caused by the choice of Kc in the calculations. An alternate set of calculations carried out for La, Gd, and Dy, based on the assumption that both Kc and K12 remain invariant in the HzO and DzO solutions, is summarized in Table V. Using this procedure differences in k34 and k43 are obtained from those where Kc was permitted to vary from the H2O value in the heavy water solutions. However, the ratios of the respective rate constants do not significantly deviate from those calculated in this paper (except for the k43 ratio for Dy which reflects the large experimental errors). Thus, the reported trends cannot be explained in terms of differences in Kc between the two solvents. The lack of influence upon the choice of Kc in the D2O solutions is illustrated in Table V for Gd, where Kc was varied from 4500 to 20,000. If Kc is much lower than the H2O constant, negative numbers are calculated for k 4 3 . For any Kc between 4000 and 20,000, all of the calculated k34 rate constants agree within experimental error. Thus, differences in the k34 rate data cannot be caused by the choice of Kc in D20. Another parameter which may be responsible for the observed differences in the Gd case is that K12 = 200 is much lower than for the other rare earths. The results
Deuterium Isotope Effects in Complexation Kinetics were recalculated assuming Klz is the mean result for the other lanthanides, and this is also shown in Table V. Although k34 decreased as K ~ increased, z k43 became negative and this is not a realistic result. Furthermore, the ratio of the k34 rate constants in the two solvents remains essentially the same, and this is the quantity which is useful in establishing the trends reported in this study. Thus, even though alternate calculation techniques can be employed to obtain rate constants from the relaxation data, resulting in some modifications in the magnitudes of the rate constants, neither the relative values for each cation nor the trends within the series can be attributed to these effects. Graphical analysis of the other three lanthanide systems shows similar results to those shown for Gd in Figure 1. In all cases the ionic concentrations within the D2O and H20 solutions are similar for each cation system and hence the rate differences cannot be attributed to concentration effects.
Discussion The ultrasonic data in HzO were extended over a much larger frequency range and consisted of many more experimental frequencies than in the previous studies.5-8 This should yield more accurate relaxation data. A comparison of the calculated k34 and k43 for the sulfate systems is shown in Figure 2. The values of k43 are similar to those obtained by Grecsek8 if one assumes a similar experimental error. Grecsek's values of k34, except for GdS04+, are consistently higher. These differences arise in that Grecsek used graphical data analysis over a more restricted frequency range. Greater differences result when one compares the results of Purdie and coworkers.6 They assume KIZ to be constant for all of the lanthanides. As in this paper, eq 4 was used to calculate the rate constants. Only in the case of SmS04+ is sufficient data present to check the results.6 Substitution of the calculated rate constants k34 and k43, along with the given 4(c) into eq 4 yields calculated relaxation times ( 2 ~ f )of 438-517 MHz. However, the experimental results range from 158 to 229 MHz. Purdie made the assumption that a certain limiting concentration of complexes must be present before the excess absorption occurred. The value of 4(c) for this point was used to determine k43. However, if one assumes $(c) to be correct as given in the paper, and obtains Kl2 using eq 13, a much improved data fit is obtained. The recalculation of Purdie's data yields the following results: Kl2 = 365; k34 = 4.1 X 108 sec-l; k43 = 3.4 x 10-7 sec-1; and the calculated relaxation times range between 159 and 229 MHz, in agreement with the experimental results. When these recalculated Sm(II1) results are replotted in Figure 2, the data follow the trend established in this study. It is noteworthy that the trends within all three studies are identical. The complexation rate constants increase toward the middle of the series, reach a maximum, and then decrease again. The interpretation of this trend in terms of a change in coordination number has been made.6,8 One of the rationales for this study was to discover if a change in the rate-determining step occurred within the series. Such a change could produce the same type of rate-determination variation. In an earlier temperature jump study involving lanthanide-anthranilate complexation in HzO and D20, a modifi-
1279
t
16.0
4 0
O L ' La
"
Pr
Ce
' Prn
"
Sm
Nd
"
Eu
"
Gd
'
Ho
Tb
Oy
'
1
Tm
Er
' Lu
Yb
Figure 2. Summary of lanthanide sulfate kinetic results: A , ref 6; 0 ,ref 8; 0, this study; 0 ,recalculated from ref 6.
cation of reaction 1was observed4 Ln3'
f
An-
kf
---L
LnAn"
(14)
kb
where An- represent4 anthranilate. For the anthranilate system a t 12.5", the value of kb was essentially the same in HzO and DzO, but significant differences occurred in k f , between solvents. The deuterium isotope effect will be represented by the ratio of a given rate constant in HzO to that in DzO. The magnitude of the deuterium isotope effect increased from 1 to about 6 from La through Gd. The two rare earths after Gd, Tb, and Dy had the same rate constants in HzO and D2O. After Dy the deuterium isotope effect increased to a value approximately 5 at Lu. With murexide as the ligand, the DzO effect was approximately 3 for Gd and slightly over 1 for Dy.4 That paper had preliminary DzO results for acetate systems,*, but recent work confirms that the process observed in the acetate system involved the formation of the bisacetato complex17 and will not be considered further. The Nd system was similar to the anthranilate system in which the dissociation isotope effect was significantly different from the complexation one.1 From the data the conclusion was reached that the deuterium isotope effect occurred in the rapid portion of step 111, including cationsolvent exchange, followed by the slow step of cation-ligand bond formation. This hypothesis would also apply to the anthranilate system. The ratios of the HzO to D2O rate constants obtained here are for La, k34 ratio = 3.5 and k43 = 0.5; Nd, k34 = 2.3 and k43 = 0.7; Gd, k34 = 1.2 and k43 = 5; and Dy, k34 = 0.8 and k43 = 2. One pattern of results occurs for the two light rare earths in which there is a large complexation isotope effect and a smaller, but opposite effect in the dissociation step. For Gd and (17) M. Doyle and H. B. Silber, J. Chem Soc., Chem. Commun., 1067 (1972). The Journa/ of Physical Chemistry, V d . 77, No. 10, 1973
1280
S.
Dy, the complexation isotope ratio is close to 1, but there is a large decrease in the dissociation rate constant upon switching to D20. Thus, one pattern of results occurs for the light rare earths, a second for the heavy rare earths, and the changeover occurs near Gd. This trend is typical in systematic studies of lanthanides and has often beeu interpreted in terms of different coordination number species existing for the light and heavy lanthanides.lSJ9 A detailed interpretation of the significance of the isotope ratio for each individual lanthanide is not possible. However, the data cannot be explained by a single process occurring in step 111, but must be interpreted in terms of at least two steps. This is the case with an Eigen-type mechanism where this step consists of cation-solvent exchange followed by cation-ligand bond formation. The large isotope differences observed for the association and dissociation steps is not consistent with the rate-determining cation-solvent, as explained previously.1 The kinetic isotope data for each lanthanide in this study is considerably different than that found with anthranilate. The big difference occurs with Gd where a
Friberg, L. Rydhag, and G . Lindblom
large complexation isotope effect occurs with anthranilate but not with sulfate. A possible explanation arises in that both sets of studies were carried out at different temperatures, and that each of the systems in each solvent has considerably different activation parameters. This explanation cannot be resolved in this study. One of the questions which this study set out to answer was whether or not similar deuterium isotope effects occur for a given lanthanide cation with different ligands. The experimental answer is that the deuterium isotope effects are not a function of solvated cation alone, but instead are a function of the solvated cation-ligand complex. Acknowledgments. The authors thank Drs. Vitullo and Pollack for several helpful discussions. Partial support for this research came from the Frederick Gardner Cottrell Program of the Research Corporation. (18) J. H. Spedding, et a/., J. Phys. Chem., 70, 2423,2430,2440,2450
(1966). (19) T. Moeller, D. F . Martin, L. C. Thompson, R. Ferris, G . R. Feistel, and W. J. Randall, Chern. Rev., 65,1 (1965).
Catalysis of p-Nitrophenol Laurate Hydrolysis in Solution Showing Transition from Reversed to Normal Micelles S. Friberg," L. Rydhag, The Swedish Institute for Surface Chemistry, Stockholm. Sweden
and G. Lindblom University of Lund, Chemical Center, Lund, Sweden (Received December 12, 1972) Publication costs assisted by the Swedish Board for Technical Development
Catalysis of the hydrolysis of p-nitrophenol laurate by normal and reversed micelles was investigated in their common solution area in the system water-cetyltrimethyltetraammonium bromide-water. Maximal reaction rates were found at the critical micellization concentrations for both kinds of micelles. In the range where transformation of one kind of micelle into the other could be expected, the reaction rates were minimal. The enhancement factor in the aqueous part was 1500, which is higher than has been reported earlier for such systems.
Catalysis by surfactant micelles has been investigated for more than a decade following the pioneering work by Duynstee and Grunwald in 1959.1 Extensive reviews2-4 of progress in the field have recently been published. The investigations prior to 1970 were, however, concerned exclusively with the effect of micelles in aqueous solutions which show the well-known association structure with the hydrophobic part of the surfactant forming the central part of the micelle and the polar parts oriented toward the surrounding aqueous solution. In amphiphilic liquids such as hexanol reversed micelles may also be formed when water and an ionic surfactant are added in sufficient quantities.5 The first report on The Journaiof PhysicalChemistry, Vol. 77, No. 10, 1973
the catalytic influence of such reversed micelles on the hydrolysis of a phenol ester was published in 19716 with reference to reversed micelles of water and cetyltrimethylammonium bromide (CTAB) in hexanol. Corresponding (1) E. F . J. Duynstee and E. Grunwald, J. Amer. Chem. SOC.,81, 4540
(1959).
(2)E. H. Cordes and R. B. Dunlop, Accounts Chem. Res., 2, 329 (1969). (3)H. Morawetz, Advan. Catai., 20, 341 (1969). (4)E. J. Fendler and J. H . Fendler, Advan. Phys. Org. Chem., 8, 271 (1969). (5) P. Ekwali, L. Mandell, and P. Soiyorn, J. Colloid lnterface Sci., 35, 266 (1971). (6)S.Friberg and S. I. Ahmad, J. Phys. Chem., 75,2001 (1971).
