J. Phys. Chem. 1983, 87,4000-4004
4000
Calorimetric Study of the Micelllzation of Pentaethylene Oxide Dodecyl Ether Gerd Olofsson Thermochemistry Laboratory, Chemical Center, University of Lund, 5220 07 Lund, Sweden (Receiv&: January 18, 1983; I n Final Form: January 24, 1983)
The micelle formation in water of pentaethylene oxide dodecyl ether, C12E5, was studied by reaction solution calorimetry. From measurements of the enthalpy change for 100-fold dilution of a solution containing 5 x mol dm-3 of C12E5 the enthalpy of micelle formation was found to be +14.0 f 1.6 kJ mol-' at 25 "C. To study the effects of temperature and concentration, enthalpies of solution of liquid C12E5 in aqueous solutions containing varying concentrations of C12E5 were measured between 10 and 40 "C. From these and other calorimetric measurements the partial molar heat capacity of C12E5 in the micellar state was determined. The partial molar heat capacity in the monomeric state was deduced by using heat capacity group contribution schemes. Thus, a value of -685 J K-' mol-' was obtained for the heat capacity change of micelle formation at 25 "C. When the value 6.4 x 10-5 mol dm-3 was used for the cmc, the entropy of micellization was found to be 127 J K-' mol-'. The presence of 5 wt% C12E5 in the solution did not produce any measurable effect on the micellization enthalpy between 10 and 30 OC. An analysis of the entropy and heat capacity changes in terms of changes in hydrophobic hydration indicated that the contribution of the alkyl chain dominated and that only minor changes occurred in the polyethylene oxide chain.
Introduction Many polyethylene oxide alkyl ether-water systems show complex phase behavior a t ambient temperatures. The water-rich part of the phase diagram consists of an isotropic micellar solution which splits into two water-rich isotropic liquid phases as the temperature is increased. As the amphiphile concentration and temperature are changed, various liquid crystalline phases appear. At sufficiently high temperatures an upper consolute temperature is probably reached in most systems. The phase behavior has been studied in detail only for tetraethylene oxide decyl ether' but other systems show the same general features.14 An approximate diagram for the C12E5-H20 system is shown in Figure 1. It is basically the diagram reported by Harusawa e t al.5 with some modifications based on results from a preliminary study of phase changes in hyusing differential scanning ~ a l o r i m e t r y . Changes ~ dration of the alkyl chain and the polar head group with temperature and concentration give rise to marked changes in aggregation behavior. Micelle formation is the basic process and a general understanding of these systems requires thermodynamic information about this micellization process. Usually estimates of enthalpy and heat capacity changes are based on measurements of the cmc and its variation with temperature. However, there are two problems with using this approach.8 If the mean aggregation number varies with temperature, the derived enthalpy will contain a term originating from changes in micelle size. There is usually no information available about such changes. Also, changes in heat capacity accompanying micelle formation are large for both ionic and nonionic detergents. The micellization enthalpy is therefore strongly temperature dependent. This should
lead to nonlinear plots of log cmc vs. 1/T. Such nonlinearity is usually not observed, an indication that measurements of the cmc either are not accurate enough or do not cover a sufficiently large temperature range to allow the determination of accurate enthalpy values. Therefore, careful calorimetric measurements are needed to provide reliable information about enthalpy, entropy, and heat capacity changes accompanying micelle formation. Two calorimetric studies of polyoxyethylene surfactants have been made, giving enthalpy values of moderate accuracy and rather uncertain estimates of the heat capacity change~.~JOThe scarcity of calorimetric data is due in part to the low cmc of these surfactants. In order to study the thermodynamics of micellization of a typical polyethylene oxide alkyl ether system calorimetric measurements have been made on the formation of micelles of pentaethylene oxide dodecyl ether, CI2E5, a t various temperatures and amphiphile concentrations. In the process of acquiring data for the micellization process, it was of interest to see if indications of micelle growth could be observed. Results of several experimental methods have been interpreted to indicate that the size of the amphiphile aggregates increases rapidly as the cloud point is approached (cf. ref 11 and references therein). With CI2E5,micelles are observed to grow with increasing amphiphile concentrations a t low temperatures." The results found are discussed in terms of changes in hydrophobic hydration of the alkyl and ethylene oxide groups following micelle formation, using group additivity schemes correlating the thermodynamics of hydration of organic nonionic solute^.'^-'^ Materials and Methods. High-purity pentaethylene oxide dodecyl ether (C12H25(OC2H,)50H, abbreviation
~~
(1)J. C. Lang and R. D. Morgan, J. Chem. Phys., 73, 5849 (1980).
(2) J. S. Clunie, J. M. Corkill, J. F. Goodman, P. C. Symons, and J. R. Tate, Trans. Faraday Soc., 63, 2839 (1967). (3) J. S.Clunie, J. F. Goodman, and P. C. Symons, Trans. FaradaS SOC.,65, 287 (1969). (4) K. Shinoda, J. Colloid Interface Sci., 34, 278 (1970). (5) F. Harusawa, S.Nakamura, and T. Mitsui, Colloid Polym. Sci. 252, 613 (1974). (6) R. Hensch, Makromol. Chem., 182, 589 (1981). (7) B. Andersson and G. Olofsson, unpublished results. (8) D. G. Hall and B. A. Pethica in "Nonionic Surfactants", Vol. 1, M. J. Schick, Ed., Marcel Dekker, New York, 1966, p 549.
(9) J. M. Corkill, J. F. Goodman, and J. R. Tate, Trans. Faraday Soc.. 60, 996 (1964). (10) J. L. Woodhead, J. A. Lewis, G. N. Malcolm, and I. D. Watson, J . Colloid Interface Sci., 79, 454 (1981). (11)P:G. Nilsson, H. Wennerstrom, and B. Lindman, J . Phys. Chem., 87. 1377 (19831. .., I
~~
(12) N. Nichols, R. Skold, C. Spink, J. Suurkuusk, and I. Wadso, J . Chem. Thermodyn., 8, 1081 (1976). (13) S. J. Gill and I. Wadso, Proc. Natl. Acad. Sci. U.S.A. 73, 2955 1197fi). ~~- -,
(14) S.Cabani, P. Gianni, V. Mollica, and L. Lepori, J . Solution
Chem., 10, 563 (1981).
0022-3654/83/2087-4000$01.50/0C 1983 American Chemical Society
Calorimetric Study of C12E5
The Journal of Physical Chemistry, Vol. 87, No. 20, 1983 4001
TABLE I: Enthalpies of Solution of a b o u t 100 mg of Liquid C,,E, in 100 cm3of Water Containing 0.05-0.15%C,,E,
tl" c [ p ( m i c )-
loa
25
30
40b
-26.05
-24.08
-22.98
-19.98
0.26
0.12
0.04
0.32
1 5 O
-26.47
W1)1
( k J mol-l) 2 s C / ( k J m o l - l ) 0.36
+ 2ot f i fl ,
L1,
0
02
0.4
0.6
0.8
Within the t w o a Supercooled samples of C,,E,. 2s is twice t h e standard deviation of the phase region. mean calculated from series of five experiments a t 10 "C, three experiments a t 25 '(2, and four experiments a t t h e o t h e r temperatures.
1.0
Weight fraction C12E5 Figure 1. Approximate phase diagram for the C,,E5-H20 system. The hatched areas indicate two-phase regions.
C12E5) from Nikko Chemicals Co., Ltd., Tokyo, Japan, was used without further treatment. Judging from gas chromatograms provided by the manufacturer, the purity of the two samples used was better than 98 mol%. The water content of the original samples was less than 0.03% by volume as determined by GC. The water content had increased to about 0.1% after transfer to the calorimeter ampules. Measurements of surface tension as a function of concentration showed clean break points and no minima, indicating the absence of surface-active impurities. The observed cmc was in good agreement with the value of 6.4 X mol L-' a t 25 "C,15 which is judged to be the best value available. Reagent-grade water produced by a Milli-Q system (Millipore AB, Goteborg, Sweden) was used to prepare all solutions. Enthalpy measurements of dissolution of liquid C12E5 or dilution of C12E5 solutions were made by using an LKB 8700 reaction-solution calorimeter with a 100- or 25-cm3 glass vessel. Samples to be dissolved or diluted were placed in cylindrical glass ampules of l-cm3 volume. They had thin end walls and narrow necks which were sealed under low flame and detached. In the dissolution experiments the calorimeter liquid consisted of either 100 cm3 of solution containing between 0.05% and 0.15% by weight of C12E5 or 25 cm3 of solution containing between 4 % and 6% C12E5. In the dilution experiments the vessel contained 100 cm3 of pure water. Small corrections for evaporation of calorimeter liquid and heat of breaking were applied as described in ref 16. A drop heat capacity ~alorimeter'~ was used to determine the heat capacity of liquid C12E5.
Results Dissolution of Liquid e,&&.Calorimetricmeasurements were made by dissolving small amounts of liquid C12E5 in aqueous solutions containing varying concentrations of C12E5, all above the cmc. As the resulting changes in concentration were small, the added C12E5 associated to give micelles a t approximately the concentration of the calorimeter liquid. The measured enthalpy changes per mole of dissolved C12E5 are denoted R(mic) - H(1). They are good approximations of the partial molar enthalpy change for the transfer of amphiphile from net liquid to micelle. In this way liquid C12E5 provided a thermodynamic reference state for determining the effects of temperature and excess C12E5 on the partial molar enthalpy of the micellar state. (15) M. J. Rosen, A. W. Cohen, M. Dehanayake, and X.-Y. Hua, J. Phys. Chem., 86, 541 (1982). (16) N. A. Mazer and G. Olofsson, J. Phys. Chem., 86, 4584 (1982). (17) J. Suurkuusk and I. Wadso, J. Chem. Thermodyn., 6,667 (1974).
TABLE 11: Enthalpies of Solution of a b o u t 100 mg of Liquid C,,E, in 25 cm3 of Solution Containing between 5 a n d 6% C,,E.
ti" c [U(mic)- H ( l ) ] ( k J mol-') 2sb/(kJ m o l - ' )
1oa
15a
25
30
-26.88
-25.71
-23.63
-22.84
0.34
0.48
0.10
0.16
Supercooled samples o f C,,E,. 2s is twice t h e standard deviation of t h e mean calculated from series of five experiments a t 1 0 'C, four experiments a t 1 5 'C, a n d three experiments a t 25 and 3 0 "C. a
Results of experiments of dissolving 0.2-0.3 mmol of liquid C12E5 in 100 cm3 of solution containing 0.05-0.15 wt % C12E5 (0.0012-0.0037 mol dm-3) are summarized in Table I. The freezing point of C12E5 is 23 "C but about half of the ampules could be supercooled long enough (ca. 10 min) to allow dissolution measurements to be made on liquid samples a t 10 and 15 "C. However, the reproducibility of these experiments is not as good as those at higher temperatures. In the measurements a t 40 "C the calorimeter liquid consisted of two liquid phases and the added C12E5 entered the C12E5-richphase.18 An additional series of measurements a t 25 "C in which the initial concentration of C12E5 was between 0.7 and 1.5 wt% gave R(mic) - H(1) = -24.06 f 0.05 kJ mol-l ( n = 4).19 To see if micellar growth indicated by other studies" was accompanied by detectable enthalpy changes, experiments were made of dissolving liquid C12E5 in solutions containing appreciable amounts of amphiphile. An exothermal contribution to the dissolution enthalpy would be an indication of growth. In these experiments 0.18-0.28 mmol of liquid C12E5 was dissolved in 25 cm3 of solution containing between 3.8% and 4.4% C12E5at 10 "C and between 5.0% and 5.5% at the other temperatures. The results are summarized in Table 11. Comparison of the results in Tables I and I1 shows that the enthalpy changes observed in the concentrated solutions are equal within error limits, or slightly less exothermal than in the dilute solutions. Thus, the calorimetric results provide no indication of micelle growth. Measurements were also made by dissolving 0.3 mmol of liquid C12E5 in 100 cm3 of toluene at 25 "C which gave H(to1uene) - H(1) = -0.25 & 0.10 kJ mol-'. In aromatic solvents like toluene the amphiphile exists as a solvated monomer.20 (18) The water content of the Cl,E5-enriched phase in equilibrium with the aqueous surfactant phase was determined by Karl Fischer titration. Samples equilibrated a t 34.4 and 41.4 "C contained 91.3 & 0.4 and 77.5 f 0.5 w t 70 of water, respectively. At 40.1 "C (the temperature of the calorimetric measurements), the CIPES-richphase is estimated to contain 20% of surfactant. (19) In the following error limits indicate the imprecision in the results expressed as twice the standard deviation of the mean, and n indicates the number of replicate measurements. (20) H. Christensen and S. E. Friberg, J. Colloid Interface Sci., 75, 276 (1979).
The Journal of Physical Chemistty, Vol. 87, No. 20, 1983
Olofsson
TABLE 111: Enthalpy of Dilution of 1 cm3 of 0.005 mol dm-3 C,,E, in 100 cm3 of Water
1759 J K-' mol-'. The uncertainty is estimated to be less than f20 J K-' mol-'. The calculated heat capacity contribution of the OCHzCH2moiety (122 and 123 J K-' mol-', respectively) is in good agreement with the value of 124 J K-' mol-' which was derived from values of the partial molar heat capacities of ethylene glycolz2and triethylene determined from calorimetric measurements. The heat capacity change for micelle formation AC,(mic) at 25 "C can now be calculated as cp(mic)- C,(aq) = -685 f 26 J K-' mol-'. In the same way the heat capacity change for-the dissolution of liquid C12E5 to give aqueous monomer C,(aq) - CJ1) can be deduced. The heat capacity values are summarized in Table IV. Thermodynamics of Micelle Formation of C I S 5 . The value of 6.4 X mol dm-3 has been chosen as the best value of the cmc at 25 "C.15 The accuracy is not stated but the uncertainty is probably of the order of f0.5 x 10" mol dm-3. The relation AG"(mic) = R T In cmc gives AG"(mic) = -23.94 f 0.20 kJ mol-'. The present study gives AHo(mic) = 14.0 f 1.6 kJ mol-' at 25 "C which leads to ASo(mic) = 127 f 6 J K-' mol-'. From the relation AGO = AH" -TAS" it is seen that the entropy term dominates and gives rise to the large negative value of AGO and accordingly the very low cmc. When the calculated value of AC,"(mic) is used, the following relations are derived to express the temperature variation of the cmc, AH"(mic), and AS" (mic) at ambient temperatures:
4002
103ci
t/"C
(mol dm-')
24.8 24.9 9.8
5.58 5.27 5.27
[H(rfq);
1
0
3
~
~
(mol dm-3) 0.051-0.055 0.042-0.049 0.046-0.049
H(mlc)l (kJ mol-' )
n
-14.0 + 1.6 -14.0 t 1.4 - 2 5 . 7 f 3.2
5 4
6
The imprecision is indicated by twice the standard de. viation of the means.
Dilution Measurements. The enthalpy of micelle formation has been determined from experiments in which 1 cm3 of solution containing 5 mmol dm-3 of C12E5 was diluted 100-fold with pure water. In the initial solution all C1zE5 is assumed to be in the micellar state while in the final solution the concentration was below cmc and the micelles disintegrated, leaving the amphiphile as hydrated monomers. Therefore, the measured enthalpy change calculated per mole of C12E5 equals the demicellization enthalpy. It is denoted H(aq) - H(mic). The results are summarized in Table 111, where ci indicates the initial concentration of C12E5 in the ampule and cf the final concentration in the calorimeter liquid. The heat evolved in the experiments is small, about 0.1 J, which leads to somewhat imprecise results. It is, therefore, not possible to determine precisely the temperature dependence of H(aq) - H(mic) in this way. However, the observed decrease when lowering the temperature to 10 "C is of the right order of magnitude (see below). The formation of micelles from the hydrated monomers is the reverse process of demicellization and the micellization enthalpy H(mic) - H(aq) is the negative of the measured enthalpy change. The formation of micelles is therefore endothermic, H(mic) - H(aq) being 25.7 f 3.6 kJ mol-' at 10 "C and 14.0 f 1.6 kJ mol-' a t 25 "C. The measured enthalpy change H(mic) - H(aq) is assumed to equal the standard enthalpy of micellization AH" (mic). Deductions of Partial Molar Heat Capacities of the Monomeric and Micellar State. The heat capacity change for the dissolution process, Cg(mic) - CJl), was calculated from the measured values of H(mic) - H(1)at low detergent concentration. The results a t 15, 25, and 30 "C give C,(mic) - C,(l) = 204 J K-' mol-' at 25 "C. The uncertainty is estimated to be f15 J K-' mol-'. The calculated value represents, to a good approximation, the partial molar heat capacity of C12E5 in the micellar state relative to the liquid. The heat capacity of liquid C12E5, CJI), was found from drop heat capacity measurement^'^ to be 886 f 3 J K-' mol-' at 30.36 "C. The temperature variation is estimated to be 0.10% per degree2' which gives C,(1) = 881 f 5 J K-I mol-' at 25 "C. The partial molar heat capacity of C12EJin the micellar state c,(mic) is found by adding Cp(l) to C (mic) - Cp(l). This gives Cp(mic) = 1085 f 17 J K-' mol- f. Due to the low cmc it was not possible to make measurements from which the molar heat capacity of the aqueous C12E5 monomer Cp(aq) could be deduced. However, it can be estimated with good acc-uracy by using group c o n t r i b u t i ~ n s . ' ~ JEstimates ~ ~ ~ ~ of Cp(aq) using the parameter values reported by Nichols et al.12 and by Cabani et al.14differ by only 3 J K-' mol-', the mean value being (21) The estimate is based on the observation that C, for liquid n-Clo to n-C12alkanes varies approximately 0.10% per degree at ambient temperatures (H. M. Huffman, G. S. Parks, and M. Barmore, J. Am. Chem. SOC.,53, 3876 (1931)). The same temperature variation of C, has been observed for pentaethylene glycol (M. A. Stephens and W. S. Tomplin, J . Chem. Eng. Data, 24, 81 (1979).) (22) C. Jolicoeur and G. Lacroix, Can. J . Chem., 54, 624 (1976).
In cmc = 26247.6/T
+ 82.3872 In T - 567.100
(1)
AH"(mic)/(kJ mol-') = 14.0 - 0.685 (T - 298.15)
(2)
AS"(mic)/(J K-' mol-') = 127 - 685 In (T/298.15)
(3)
In eq 2, AH" = 0 a t around 45 "C and accordingly eq 1 has a minimum at this temperature.
Discussion The formation of micelles in aqueous solution from liquid CI2E5is exothermal by -24 kJ mol-' at 25 "C while the micellization from aqueous monomers is endothermal by 14 kJ mol-'. Thus, the enthalpy of solution of C12E5 in water below the cmc is strongly exothermic, R(aq) - H(1) = -38 k J mol-'. Also, the formation of micelles from aqueous monomers is entropy driven, the endothermal formation enthalpy counteracting aggregation. The dissolution enthalpy is strongly temperature dependent, as shown by the value C,(aq) - C,(1) = 878 J K-' mol-'. On the other hand, the enthalpy of dissolution of C12E5in an organic solvent like toluene where only weak interactions are expected is close to zero. The heat capacity change can here safely be assumed to be small. The strong interaction between the monomeric amphiphile and water and its strong dependence on temperature are fundamental to the aggregation behavior of nonionic surfactants. It is generally agreed that the interior of the micelle in many respects resembles a hydrocarbon liquid as originally suggested by H a r t l e ~ . The ~ ~ experimental evidence for a close similarity between the micelle core and liquid hydrocarbons has been reviewed by Lindman and W e n n e r ~ t r o m .The ~ ~ reversal of micellization, that is, the breaking up of aqueous, nonionic micelles, resembles in some important aspects the dissolution of hydrocarbons in water. The change in solvation of the alkyl chain is very similar to the change experienced by a hydrocarbon solute (23) K. Bystrom and G. Olofsson, unpublished results. (24) G. S. Hartley, J . Chem. SOC.,1968 (1938). (25) (a) H. Wennerstrom and B. Lindman, Phys. Rep., 52, 1 (1979); (b) B. Lindman and H. Wennerstrom in "Topics in Current Chemistry", Vol. 87, F. L. Boschke, Ed., Springer-Verlag, West Berlin, 1980.
The Journal of Physical Chemistry, Vol. 87, No. 20, 7983 4003
Calorimetric Study of C,2E5
TABLE IV: Partial Molar Heat Capacities of C,,E, in Liquid, Monomer, and Micellar States and Heat Capacity Changes for Dissolution of Liquid C,,E, To Give Monomers Respective Micelles and for Micelle FormationQ CpV)=
-
881
*
5
c p (aq) -
c8 f 8(1)= ? 21 a
Values a r e
-
-
-
-
Cp(as) = 1 7 5 9 f 20
C, (mic) =
Cp(mic)-
Cp(mic)-
C p (1) = 204 ? 1 5
1085
?
17
C p ( a s )=
- 6 8 5 + 26
given in J K-' mol-' a n d refer to 2 5 "C.
upon dissolution. The interactions between hydrocarbon solutes and water show some distinct features, which have led to the introduction of the concept "hydrophobic hydration" and "hydrophobic interactions.26 The dissolution of liquid hydrocarbons in water at room temperature is characterized by a small enthalpy change but a large and negative entropy change leading to a large and positive free energy change, which means low solubility. The enthalpy of solution varies strongly with temperature, the heat capacity change being large and positive. The volume change is negative. These changes in the thermodynamic properties upon dissolution indicate a profound influence of hydrocarbon solutes on water. Several properties have been found to vary linearly with the number of hydrogen atoms in hydrocarbon s01utes.l~ On the basis of this observation Gill and WadsO13 have derived an equation of state for the hydrophobic effect correlating the Gibbs energy change, the enthalpy change, and the heat capacity change for the dissolution of liquid hydrocarbons in water with the number of hydrogens in the molecule. Cabani et al.14 have recently made an extensive analysis of existing data on the hydration of gaseous organic nonelectrolytes and derives sets of group contributions that can be used to describe the Gibbs energy, enthalpy, heat capacity, and volume changes for the transfer of organic compounds from the gaseous state to aqueous solution a t 25 "C. Group contributions to the partial molar heat capacity and partial molar volume at infinite dilution were also derived. If one uses the relation derived by Gill and WadsO13to express the enthalpy of solution of liquid hydrocarbons in water, the contribution of the Clz alkyl chain to the solution enthalpy of liquid C12E5 at 25 "C (H(aq) - H(1) = -38 kJ mol-') is estimated to be 3.5 kJ mol-'. From enthalpies of solution in water of triethylene glycol and of polyethylene oxideB together with enthalpy of melting and heat capacities of solid and liquid polyethylene oxide,27the contribution of the pentaethylene oxide group to the solution enthalpy is estimated to be about -35 kJ mol-'. The two different estimates agree in that they show that the alkyl chain contribution to the enthalpy of solution of liquid C12E5 is small a t 25 "C and that the exothermal interaction between the ethylene oxide groups and water dominates. The corresponding separation of the hydrophobic and hydrophilic group contributions to the heat capacity change for dissolution of liquid C12E5requires the separate evaluation of the contributions to C,(aq) and Cp(l). The alkyl chain contribytes 1140 J K-' mol-' and the E5group 619 J K-' mol-' to C,(aq) of C12E5.12The alkyl chain contribution to Cp(l) will be close to C, for dodecane (375 J K-' mol-')2s and the contribution from the E, group close to C, for pentaethylene glycol (516 J K-' mol").29 (26) F. Franks in 'Water-A Comprehensive Treatise", Vol. 2, F. Franks, Ed., Plenum Press, New York, 1972, Chapter 1. (27) U. Gaur and B. Wunderlich, J . Phys. Chem. Ref. Data, 10, 1001 (1981). (28) R. C. Weast and M. J. Astle, Eds., "CRC Handbook of Chemistry and Physics", 62nd ed., CRC Press, Boca Raton, FL, 1981.
Thus, of the heat capacity change of 889 J K-' mol-' for the hydration of liquid C12E5, about 770 J K-' mol-' can be attributed to the alkyl chain while the contribution of the E5 group is only about 120 J K-l mol-'. The large increase in the heat capacity of the alkyl chain indicates strong interactions between the hydrocarbon groups and the surrounding water. For the pentaethylene oxide group this interaction is to a large extent offset by the interjacent oxygen atoms which give a large negative heat capacity contribution possibly stemming from hydrogen bond interaction between water and the ether groups. In forming micelles, the alkyl groups leave the water milieu to enter the interior of the micelles. This should result in a heat capacity contribution approximately equal to the heat capacity change for the transfer of the corresponding alkane from water to the net liquid. The polar head group will be much less affected by micellization and will give only a minor heat capacity contribution. Thus, the heat capacity change for micelle formation of C12E5can be expected to be close to -770 J K-' mol-'. The experimental value is -685 f 26 J K-1 mol-', which is in satisfactory agreement with the expected value considering the approximations made. The equation of state for the hydrophobic effect13 also permits the estimation of the part of the micellization entropy that can be attributed to the alkyl chain. The entropy of transfer of a Clz-hydrocarbonchain from water to a hydrocarbon environment is estimated to be 132 J K-' mol-' with concentrations expressed in mol dm-3. The entropy change for the micelle formation of C12E5 is 127 f 6 J K-' mol-'. The close agreement between the two values indicates that the large positive entropy of micellization stems mainly from the dehydration of the alkyl chain. This is also indicated by a comparison of estimated alkyl chain contributions with the entropies of micellization for C8E6, and C10E6.' Suggestions that the desolvation of ethylene oxide groups should be the major contributing factor to the positive entropy of micellization of polyoxyethylene oxide ethers9J5~30 do not find any support from this analysis. On the contrary, thermochemical evidence points to only minor changes of solvation of the polar head group upon micelle formation. However, the interaction between water and the ethylene oxide groups will change with temperature and amphiphile concentrati0n.l~~'The positive heat capacity change for the hydration of the ethylene oxide chain will give a positive contribution to the Gibbs energy change with increasing temperature which can be seen as a decrease in solubility. The polar chains will be less effectively hydrated due to decreasing availability of water as the surfactant concentration is increased and interactions between chains in the same micelle as well as in different micelles will increase. However, in a recent NMR study of water self-diffusion, it was found that changes of hydration with temperature and concentration in this type of system are not drastic3' For instance, the surfactant aggregates remain strongly hydrated close to the cloud point and there is no direct relation between dehydration and the liquid-liquid phase separation. Sensitive and accurate enthalpy of dilution and heat capacity measurements are needed in order to (29) M. A. Stephens and W. S. Tomplin, J . Chem. Eng. Data, 24,81 (1979). (30) M. J. Schick, J. Phys. Chem., 67, 1796 (1963). (31) A theoretical discussion of the thermodynamic properties of polyethylene oxide-water solutions has recently been presented by R. Kjellander and E. Florin, J . Chem. SOC.,Faraday Trans. 2, 77, 2053 (1981). (32) P.-G. Nilsson and B. Lindman, J . Phys. Chem., in press.
J. Phys. Chem. 1983, 87, 4004-4007
4004
explore further the thermochemical effects of changes in hydration of the polyethylene oxide chain. Acknowledgment. This work was supported by grants from the Swedish Natural Science Research Council. The skillful assistance of Mrs. Stina Bergstrom is gratefully acknowledged. Note Added in Proof. An extensive study has recently
been made of the phase behavior of polyoxyethylene surfactants with water by Mitchell et al.33 A more complete phase for C~2E5-water can be found in their paper* Registry No. Pentaethylene oxide dodecyl ether, 3055-95-6.
I
(33) D. J. Mitchell, G. J. I. Tiddy, L. Waring, T. Bostock, and M. P. McDonald, J. Chem. Soc., Faraday Trans. I , 79, 975 (1983).
Solvation Enthalpies and Crystal Lattice Energies of Polyacene Anion Radicals Gerald R. Stevenson,' Laurel E. Schock, and Rlchard C. Relter Department of Chemistry, Illinois State University, Normal, Illinois 61761 (Received: January 24, 1983; I n Final Form: April 26, 1983)
-
-
Calorimetric techniques coupled with thermochemical cycles have been utilized to measure both the solvation (A-., + Na+, A-*,Na+THFenthalpies and crystal lattice energies (-AHo for A--, + Na+, Na+A--,) of several polyacene anion radical sodium salts. These new data have been combined with previously published results, and a correlation has been found between both AHosolv and the crystal lattice energy (U,) and the total a energy of the polyacene. Surprisingly, these two linear correlations (plots of both -AHosolvand U,, vs. total a energy are nearly linear) have similar slopes. These results are used to predict the solvation enthalpy and crystal lattice energy of the sodium benzene anion radical. The results are explained in terms of ion association and crystal packing. An improved estimate of the EA of pentacene is suggested.
Introduction The chemistry of anions in the condensed phases has long been known to be very different from that in the gas phase. Statistical thermodynamic calculations show that the simple dissociation of the HBr to Br- and H+ is hundreds of orders of magnitude smaller in the gas phase than that measured in any solvent system. This fact has been more recently demonstrated by the observation that the order of basicities of the alkoxide anion is reversed in the gas phase as compared to those in solution.' This large difference in the chemistry of anions in the gas and condensed phases is due to the very large heats of solvation and crystal lattice energies that are present in the condensed phases. Either gas-phase or condensed-phase anions can be generated by two different methods from the neutral systems, either by the extraction of a proton or by the addition of an electron. We have recently developed an experimental technique for the determination of both crystal lattice energies2 and solvation enthalpies3 for organic anions generated by the addition of an electron. The method involves the measurement of the heat of reaction of either the solvated (reaction 1)or the solid (reaction 2) A-',Na+THF + H2O1Na+A-.,
+ HzOl
+ '/2AHZs + NaOH,, '/zAs + '/2AH2s + NaOH,,
l/pA,
-
(1) (2)
anion radical salt with water to produce the neutral hydrocarbon, hydrogenated hydrocarbon, and aquated alkali-metal hydroxide, where THF represents tetrahydrofuran. (1) Brauman, J. I.; Blair, L. K. J. Am. Chem. SOC. 1970, 92, 5986. 1979,101, (2) (a) Stevenson, G. R.; Wiedrich, C. R. J. Am. Chem. SOC. 5094. (b) Stevenson, G. R.; Wiedrich, C. R.; Clark, G. J . Phys. Chem. 1981, 85, 374. (3) (a) Stevenson, G. R.; Williams, E. J . Am. Chem. SOC.1979, 101, 5910. (b) Stevenson, G. R.; Chang, Y. J. Phys. Chem. 1980, 84, 2265.
0022-3654/83/2087-4004$0 1.50/0
TABLE I : Thermodynamic Parameters (in kcal/mol) for Aromatic Hydrocarbons
EAc
hydrocarbon
benzene
naphthalene anthracene tetracene
pentacene pyrene perylene
26.6 +3.5 i 12.7 A 16.0 t 19 -
113.3 - 21.1 t7.1
1@Hlb
t
1.1 0.37
i
1.4 1.0
* t
+8.09 -17.4 t23.4 -28.9 35 -7
i
ref for EA
1.5
-22.5 -30 -21.7
7 8 8 8
see text 8 9
AH",^^'^ 5.6 -3.2 -17 -12.5 -15 i
- 2.7 -15
phenanthrene t 1.7 8 7.7 Of the parameters reported in this table, the EA'S have by far the greatest error. The EA of benzene was determined via electron transmission spectroscopy. Those taken from ref 8 were obtained by using the method of electron capture with a determined intercept. The heat of vaporization of benzene was taken from Osborne, N. S.; Ginnings, D. C. J. Res. N a f l . Bur. S t a n d . (U.S.) 1 9 4 7 , 39, 453-77. The heat of hydrogenation of naphthalene to give 1,4-dihydronaphthalene was obtained from: Shaw, R . ; Golden, D. M.; Benson, S. W. J. P h y s . Chem. 1 9 7 7 , 81, 1 7 1 6 . This reference is in agreement with the value reported i n : Cox, J. D. ; Pilcher, G. "Thermochemistry of Organic and Organometallic Compounds"; Academic Press: New York, 1 9 7 0 ; p 170. The remaining heats of hydrogenation and heats of sublimation are listed in ref 3. The dihydro compounds referred to are those resulting from the reaction of the anion with water as described in the Appendix.
The enthalpies of the reactions listed in eq 1 and 2 can be combined with the heat of reaction of sodium with water (-44.1 kcal/mol)? the heat of sublimation of sodium (+25.9 kcal/mol)? the ionization potential of sodium (118.4 kcal/mol),6 and some constants for the hydrocarbon (A) (4) Gunn, S. R. J. Phys. Chem. 1967, 71, 1386. ( 5 ) Hicks, W. T. J . Chem. Phys. 1963, 38, 1873.
@ 1983 American Chemical Society
