J. Phys. Chem. 1980, 84, 2265-2268
tertiary amines. Possible reasons for this behavior are discussed.
References and Notes (1) Pimentel, G. C.; McClellan, A. L. "The Hydrogen Bond"; Reinhold: New York, 1960. (2) Joesten, M. D.; Schaad, L. J. "Hydrogen Bonding"; Marcel Dekker: New York, 1974. (3) Farah, L.; Giles, G.; Wilson, D.; Ohno, A,; Scott, R. M. J. Phys. Chem. 1979, 83, 2455. (4) Maron, S. H.; Filisko, F. E. J . Macromol. Sci.-Phys. 1972, B6, 57. (5) Joesten, M. D.; Drago, R. S. J . Am. Chem. SOC.1962, 84, 3817.
2205
(6) Epley, T. D.; Drago, R. S. J . Am. Chem. SOC. 1967, 89, 5770. (7) Zharkov, V. V.; Zhltinkina, A. V.; Zhokhova, F. A. Fir. Khim. 1970, 44, 223. (8) Fritzsche, M. Ber. Bunsenges. Phys. Chem. 1964, 68, 459. (9) Gramstad, T. Acta Chem. Scand. 1961, 16, 807. (10) Singh, S.; Rao, C. N. R. Can. J. Chem. 1966, 4 4 , 2611. (11) Lin, M.; Scott, R. M. J . Phys. Chem. 1972, 76, 587. (12) Van Thiel, M.; Eecker, E. D.; Pimentel, G. C. J. Chem. Phys. 1957, 27, 95. (13) Liddel, U.; Becker, E. D. Spectrochim. Acta 1957, 10, 70. (14) Eecker, E. D.; Liddel, U.; Shoolery, J. N. J . Mol. Spectrosc. 1958, 2, 1. (15) Griffiths, V. S.; Socrates, G. J . Mol. Spectrosc. 1966, 21, 302.
Anion Radical Solvation Enthalpies as a Function of Cation Size Gerald R. Stevenson" and Yoh-Tz Chang Department of Chemjstty, Illinois State University, Normal, Illinois 6 176 I (Received:March IO, 1980)
Calorimetric methods have been utilized to measure the enthalpies of reaction of solvated anthracene anion radical ion pairs with water. These enthalpies were then used in a thermochemical cycle to obtain the heats of solvation of the separated ions in the gas phase (AH" for AN-., + M+, (AN-.,M+),,~v).These enthalpies were found to be more exothermic when dimethoxyethane (DME) rather than tetrarhydrofuran (THF) served as the solvent. However, for both of these solvents the heats of solvation are more exothermic for the smaller cations (increasingexothermicity Cs+ < K+ < Na+ < Li'). In THF this enthalpy varies by more than 50 kcal/mol with the size of the cation. The results are explained in terms of ion solvation and ion association, and these results, are compared to a previous study where the anion size was varied.
-
Neutral hydrocarbons can capture electrons that come from a variety of sourlces including electron beams, alkali metals, and negatively charged electrodes to yield anion radicals in the gase phase,l in solution,2 or in the solid states3 The anion radical systems are known to be much more thermodynamically stable in the two condensed states than in the gas phasea3This is due to the very strong intermolecular stabilizing effects of crystal lattice energy in the solid state3 and of ion solvation and ion association in the solution state: In the gas phase the only important consideration in the stability of the anion radical is the electron affinity of the neutral specie^,^ which is 12.7 kcal/mol for anthracenee6 If this gas-phase anion radical is generated from the transfer of an electron from sodium metal, which has an ionization potential of 118.4 kcal/mol,' the gas-phase ions lie 105.7 kcal/mol higher in energy than However, the neutral gas-phase odium and anthra~ene.~ the very large crystal lattice energy (-166.6 kcal/mol) and solvation enthalpy (including ion association) in THF (-178.7 kcal/mol)8 bring the energies of the solid material and of the THF solvakd ion pairs to 50.9 and 63 kcal/mol lower than that of the gas-phase neutral species, respectively. The immense effect that solvation has upon the thermodynamic stability and thus chemistry of anions prompted us to carry out a systematic study of the enthalpy of solvation and thermodynamic stability of the anthracene anion radical in several solvents and with a variety of alkali metal cations. Despite the obvious importance of solvation enthalpies in controlling the chemistry of organic anion radicals, only one previous report of experimental solvation enthalpies has appeared.8 Szwarc and co-workers have collected a vast amount of qualitative information concerning the solvation of anion radical and dianion ion pairs through studies of anion 0022-3654/%0/2084-2265$0 1.OO/O
radical disproporti~nation.~ For ion pairs involving the anthracene anion radical and K+, Na+, or Li+ serving as the cation in THF (tetrahydrofuran) or in DME (dimethoxyethane), cation solvation appears to decrease as the size of the cation increase^.^ Further, DME appears to have a greater capacity to solvate cations than does THF. For instance, K+ is $oorly solvated when it is associated with the anthracene anion radical in THF, but it is strongly solvated in the same ion pair in DMESga The three most commonly used solvents for alkali generation of anion radicals are THF, DME, and HMPA (hexamethylphosphoramide). In HMPA hydrocarbon anion radicals exist free of ion association.l0J1 Since it was our intention to investigate the effect of cation upon the heat of formation of solvated ion pairs, we have chosen to include DME and THF in this first study of the effect of cation upon the heats of formation of ion pairs from the separated gas-phase ions. In a previous stud9 the enthalpies of solvation of a series of polyacene anion radicals generated via sodium reduction in THF were measured, and we were surprised to find that the enthalpies of solvation of the separated gas-phase ions to form the solvated ion pairs (AH"for the reaction depicted in eq l) are within experimental error for the entire A-e,
+ Na+,
-
(A-.,Na+)THF AH" = -177 kcal/mol
(1)
series of polyacene anion radicals that were included in the study. A-. represents the anion radical of naphthalene, anthracene, tetracene, phenanthrene, pyrelene, or pyrene. The smaller anions evidently form tighter ion pairs with the sodium cation because of their more localized charge densities. These tighter ion pairs then interact more weakly with the solvent (THF). The larger anion radicals 0 1980 American Chemical Society
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The Journal of Physical Chemistry, Vol. 84, No. 18, 1980
form weaker ion pairs and are better solvated. Evidently, the stronger solvent-ion interactions formed in the larger, looser ion pairs counterbalance their more endothermic ion association enthalpies.s We thought it would be interesting to find out whether changes in the size of the cation will also have a counterbalancing effect in the total solvation enthalpy. The disproportionation work mentioned above would suggest that this is not the case, and quite different enthalpies of solvation would be obtained upon variation of the cation as opposed to variation of the anion.
Experimental Section The anion radical solutions in DME and THF were generated and sealed into small evacuated glass bulbs in the manner previously described.s The thin-walled glass bulbs were then placed into the modified cell of a Parr solution calorimeter as described earlier.12 After the bulbs were broken under 100 mL of deoxygenated water, the change in the temperature (AT) of the calorimeter due to the heat of reaction of the salt with water and the heat of solution of the solvent in water was measured. After the reaction, the contents of the calorimeter were titrated with standardized HC1 to obtain the amount of anion radical salt in the bulb. The broken glass was then collected and weighed, and the amount of solvent in the bulb was taken as the difference in the weight of the intact bulb plus its contents and the sum of the weights of the salt, neutral hydrocarbon, and glass. The anthracene was purchased from Aldrich Chemical Co. and sublimed before use. The THF and DME were distilled from benzophenone ketyl and stored under vacuum over NaK2. The solvents were then distilled directly into the apparatus used to generate the anion radical solutions. Both DME and THF are completely miscible with water, but neither anthracene nor dihydroanthracene is soluble in the water-THF or water-DME mixture left in the calorimeter after the reaction. The only organic products that were obtained from the reaction in the calorimeter were anthracene and 9,lO-dihydroanthracene, and the kinetics and mechanism of this protonation reaction have been the subject of several rep o r t ~ The . ~ ~enthalpy measured in the calorimeter is that of
- ... ,
or THF
where s = solid. For all of the calorimetry experiments, the amount of solvent (THF or DME) varied from 2.0 to 4.0 g. Thus the temperature change in the calorimeter due to the aquation of the solvent varied from 1 to 2 "C. The anion radical concentration varied from 0.2 to 1.0 M. The presence of the THF or DME in the water in the calorimeter after the bulb is broken may have a slight effect upon the enthalpy of the process. This effect was, however, too small to be detected with our technique. Since the quantity of MOH formed is very small compared to the amount of water, the reactions were considered to generate MOH,, at infinite dilution.
Results Since the heat evolved in the calorimeter is not just that for the reactions depicted in eq 2 but also that due to the
Stevenson and Chang
2.0
I
0
2.0
4.0
grams o f solvent
Figure 1. Plots of the change in the temperature of the calorimeter vs. the grams of solvent sealed into the glass bulbs. The slopes of the lines are 0.407 and 0.530 deg/g for THF and DME, respectively. The slope for THF was measured previously.
TABLE I: Enthalpies of Reaction of Solvated Anion Radicals with Water solvent anion radical A H , kcal/mol ref THF AN- ,Li+ -30.6 f 1.8 this work THF AN-.,Na+ -28.9 * 1.4 8 THF AN-*,K+ -21.2 * 1.1 this work THF -24.8 * 1.4 this workb AN-. ,CS+ DME -18.8 f 0.8 this work AN-., Li+ -23.0 * 0.7 this work DME AN-.,Na+ DME -19.4 f 1.0 this work AN-.,K+ HMPA AN-. t Na+ -27.8 2 0.5 8 AN = anthracene. Because of solubility problems, this value was obtained indirectly. See Discussion.
-
aquation of the solvent, it is first necessary to determine the amount of heat liberated from this latter process. Samples of pure DME were placed into the evacuated glass bulbs, which were crushed under 100 mL of deoxygenated water in the calorimeter. A plot of the amount of DME in the bulbs vs. the change in the temperature of the calorimeter is linear (Figure 1). Similar plots for HMPA and THF were reported previously.8 From Figure 1the temperature change in the calorimeter due to the aquation of the solvent can be determined. Subtracting this from the total observed temperature change in the calorimeter yields that due to the reactions depicted in eq 2 and Table 11. The slopes of the lines generated from plots of the change in the temperature of the calorimeter minus that due to the solvent (AT due to the anion radical reaction) vs. the millimoles of anion radical in the glass bulbs are linear (Figures 2 and 3). Multiplying these slopes by the heat capacity of the calorimeter (-119.3 cal/deg) yields the enthalpy of the reaction shown in eq 2, where THF or DME serves as the solvent (Table I). From Table I it is clear that each cation-anion radical ion pair reacts more exothermically with water when it is solvated in THF than when it is solvated in DME. This must reflect the fact that the heats of solvation are more exothermic in DME. The actual heats of solvation of the gas-phase ions can be calculated via the thermochemical cycle shown in Table 11. The cesium anthracene anion radical salt is not sufficiently soluble in THF to allow a direct determination of its enthalpy of reaction with water. For this reason the solid anion radical salt was reacted with water in the
The Journal of Physical Chemistry, Vol. 84, No. 18, 1980 2267
Anion Radical Solvation IEnthalpies
TABLE 11: Enthalpies of Reaction AH" (THF), kcal/mol
M
reactiona ~
A",, + AN, + 2MOHaq --> 2(AN-*,M+)g,h + 2H201 2M, + 2H,01-* 2MC)Haq + H,, 2Mf + 2M, 2M + 2eg -* 2M AN, + H , ~ - +AN& 2ANg 2AN, 2AN-., -+ 2AN, + 21ei AN-., t M+, (AN-*,M+),h -+
-+
a ANH, represents 9,lO-dihydroanthracene. about * 3 kcal/mol.
Li
Na
K
cs
AH" (DME), kcal/mol Li Na K ref
61.2 57.8 42.4 49.6 37.6 46.0 38.8 -106.4 -88.2 -94.0 -96.6 -106.4 -88.2 -94.0 -71.2 -51.8 -42.8 -37.6 -77.2 -51.8 -42.8 -248.6 -236.8 -200.2 -179.6 -248.6 -236.8 -200.2 -17.0 -17.0 -17.0 -17.0 -17.0 -17.0 -17.0 -47.0 -47.0 -47.0 -47.0 -47.0 -47.0 -47.0 25.4 25.4 25.4 25.4 25.4 25.4 25.4 -205 -179 -167 -151 -217 -185 -168
14 15 16 17 18 19
By propagating the standard deviations, the error in the final enthalpy is TABLE 111: Solubility of the Cesium Anthracene Anion Radical Salt in THF T,"C solubility, mol/L 22 0.067 30 0.062 40 0.058 0.050 50
Discussion The exothermicity of reaction 3 is controlled by two
A-*,
Figure 2. Plots of the change in the temperature of the calorimeter vs. the millimoles of anion radical solvated in DME and sealed into the evacuated glass bulbs: (0)sodium system: (0) potassium system.
Figure 3. Plots of the chainge in the temperature of the calorimeter vs. the millimoles of anion radical solvated in THF and sealed into the evacuated glass bulbs: (0)lithium system; (0) potassium system.
calorimeter as previourily described.' Then, the solubility of the cesium salt in THF was studied as a function of temperature to yield the enthalpy of solution. These two enthalpies were then combined as shown by 2Cs+AN-., CS'AN-.,
+ 2Hz01- ANH2, + AN, + ~CSOH,,
-
AHo = -53.4 kcallmol
(AN-*,Cs+)THF
(AN-*,Cs+)THFt HZ01
-+
'/2ANHz,
AHo = -1.9 kcallmol
+ 'IZAN, + &OH, AHo = -53.412 i1.9
The solubility data were taken in exactly the same manner as described for the salts of cyclooctatetraeneZ0and are presented in Table 111.
+ M+,
+
(A-.,M+),olv
(3)
factors: the Coulombic interaction between the cation (M+) and the anion radical (Am-)and the solvent ion interactions. When the size of the anion is diminished, the increase in ion association is accompanied by desolvation. These counterbalancing effects result in the enthalpy of solvation (AH for reaction 3) being independent of anion size.8 From Table I1 it is obvious that this cancellation effect does not take place when the size of the cation is varied. In fact, the enthalpy of solvation of the gas-phase separated ions decreases by more than 50 kcallmol when the cation is varied from Cs+ to Li+. Incidentally, this change is more than twice the magnitude of the enthalpy measured in the calorimeter for the cesium system. For both THF and DME the exothermicity of the reaction shown in eq 3 increases as the size of the cation diminishes. This is best explained in terms of the fact that the smaller cations involved in the ion pairing are better solvated than are the larger cations involved in ion pairing? For example, Li+ is well solvated in (AN--,Li+)THF while the K+ cation is poorly solvated in ( A N - S , K + ) ~Thus, ~ . ~ the effect upon total solvation enthalpies of ion association and ion solvent interactions act in the opposite direction and cancel when the anion size is varied, but they are additive when the cation size is varied. For all of the cations studied, the solvation enthalpies are more exothermic in DME than in THF. This is the expected result, since DME is known to solvate alkali metal cations more strongly than does THF. It is interesting to note that the solvation enthalpies in THF and DME are much greater for the smaller cations. Evidently the better cation solvation of DME over that of THF is diminished for the larger cations. References and Notes (1) Becker, R. S.; Wentworth, W. E. Nature(London) 1964, 203, 1268. (2) Russell, G. A. In "Determination of Organic Structures by Physical Methods"; Nachod, F. C., Kuckerman, J. J., Eds.; Academlc Press: New York, 1971; p 301. (3) Stevenson, G. R.; Wiedrich, C. R. J. Am. Chem. Soc. 1979, 101, 5092. (4) Sharp, J. H.; Symons, C. R. In "Ions and Ion Pairs In Organic Reaction": Szwarc, M., Ed.; Wiiey-Interscience: New York, 1972; Chapter 5.
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J. Phys. Chem. 1980, 84, 2268-2272
(5) Jordan, K.D.; Michejda, J. A.; Burro, P.D. J. Am. Chem. SOC.1976, 98, 1295. (6) Becker, R. S.; Chen, E. J. Chem. Phys. 1966, 45, 2403. (7) Lotz, W. J . Opt. SOC.Am. 1967, 57, 873. (8) Stevenson, G. R.; Williams, E. J. Am. Chem. SOC.1979, 701,5910. (9) (a)Rainis, A.; Szwarc, M. J . Am. Chem. Soc. 1974, 96, 3008. (b) Szwarc, M. “Ions and Ion Pairs in Organic Reactions”; Szwarc, M. Ed.; Wiley-Interscience: New York, 1974; Chapter 1. (10) Normant, H. Angew. Chem., Inf. Ed. Engl. 1967, 6 , 1046. (1 1) Cserhegyi, A.; Chaudhuri, J.; Franta, E.; Jagur-Grcdzinski, J.; Szwarc, M. J. Am. Chem. SOC. 1967, 89, 7129.
(12) Stevenson, G. R.; Williams, E.; Caklwell, G. J. Am. Chem. Soc. 1979, 701, 520. (13) Bank, S.; Bockrath, B. J . Am. Chem. SOC. 1972, 94, 6076. (14) Gunn, S. R. J. Phys. Chem. 1967, 77, 1386. (15) Hicks, W. T. J. Chem. Phys. 1963, 38, 1873. (16) Lotz, W. J . Opt. SOC. Am. 1967, 57, 873. (17) Stevens, B. J. Chem. SOC. 1953, 2973. (18) Cox, J. D.; Pilcher, G. “Thermochemistry of Organic and Organometallic Compounds”; Academic Press: London, 1970. (19) Becker, R. S.; Chen, E. J. Chem. Phys. 1966, 45, 2403. (20) Stevenson, G. R.; Ocasio, I. J . Phys. Chem. 1975, 79, 1387.
Solubilization of Methylparaben in Nonionic Surfactant Micelles in Aqueous Solution Ayako Goto,” Masumi Nihei, and Fumlo Endo Shizuoka College of Pharmacy, 2-2- 7 Oshika, Shizuoka, Japan (Received: October 9, 1979; In Final Form: April 30, 1980)
Solubilizationof methylparaben in micellar hexaoxyethylene lauryl ether (C12E6)solution was investigated by gel filtration. The results showed that the distribution of methylparaben between the micellar and aqueous phases is nonideal. Furthermore, it was observed that the cmc, which is a thermodynamic parameter of micellization, decreased by the addition of methylparaben. This indicates that the solubilization of methylparaben by C&6 micelles resulted in the stabilization of micelles. In order to explain the nonideality and the decrease in the cmc, the theory of regular solutions was applied to the solubilized system on the basis of small system thermodynamics. The results suggested that the solubilization data are compatible with a simple picture of nonideality as described by the regular solution theory. Furthermore, the degree of interaction of methylparaben with C12E6 in the mixed micelle was estimated from the regular solution model. The interaction of methylparaben and Cl2& in the mixed micelle may be closely related to the increase in size of the hydrated micelle and the lowering of the cloud point.
Introduction The solubilizing process in micellar solutions is complex and a quantitative theory has not yet been sufficiently established. The process is assumed to be closely related to interactions between the solubilizate and surfactant in the mixed mice1le.l It is generally accepted that the deactivation of preservatives in the presence of nonionic surfactants arises from the solubilization of preservatives in the However, the type and strength of their interactions are not well known, although studies of the solubilized systems are of great significance in relation to pharmaceutical applications. The authors have been interested in the solubilization of preservatives such as methylparaben in micellar solutions of nonionic polyoxyethylene surfactants. It is to be expected that solubilized systems of nonionic surfactant solutions are simpler than those of ionic surfactant solutions. That is, in the case of nonionic surfactant systems, it is not necessary to consider the change in micellar surface charge due to the addition of a solub i l i ~ a t e .However, ~ few thermodynamic studies on solubilized systems of nonionic surfactant solutions have been reported.6 Especially, the mechanism of the stabilization of a micelle by solubilization has scarcely been studied, since cmc’s (critical micelle concentration) of nonionic surfactants are extremely low compared with those of ionic surfactants, and it is also difficult to detect the cmc-decreasing effect of solubilization. The authors have investigated solubilized systems of sodium lauryl sulfate micellar solutions, solubilizing alkylparabens, by gel filtration7i8and have analyzed the solubilized systems thermodynamically.+11 Furthermore, a gas chromatographic procedure was used for analysis of hexaoxyethylene lauryl ether (Cl2EG),a nonionic surfactant.12 This procedure was used in conjunction with gel 0022-3654/80/2084-2268$01 .OO/O
filtration on Sephadex G-200 for investigation of the solution state of nonionic surfactant micelles in aqueous solution.12 The present investigation was also undertaken in order to study solubilized systems of methylparaben in micellar solutions by gel filtration, and to apply a distribution model to solubilized systems. On the basis of the analysis, the solution state of methylparaben in micellar solution was explored.
Experimental Section Materials. Polyoxyethylene n-lauryl ethers, C12H250(CH2CH20),H (C12E,, n = 5 and 6), were obtained from Nikko Chemicals Co. Ltd. Methylparaben was of reagent grade. Distilled water was used after degassing in vacuo. Preparation of Sample Solutions. The sample solutions were prepared by dissolving methylparaben in 2 or 4 mM C12E6. Gel Filtration. Sephadex G-200 was placed in a jacketed column (1.6 X 40 cm, Pharmacia, Sweden) maintained at 20.0 f 0.5 “C. The ascending gel filtration was carried out in the same way as described previously.12 The total volume of the gel bed was 60.5 mL. The void volume of the gel bed was 24.7 mL. Tail analysis for ClzEs solution solubilizing methylparaben was carried out as follows: After the gel column was preequilibrated with 160 mL of sample solution, the sample was eluted with water at a rate of about 6 f 1mL/h. Approximately 2-mL portions were collected by an automatic fraction collector. Exact volumes were obtained by measurement of the weight and specific gravity of each fraction. Analytical Methods. The concentration of methylparaben in C&6 solution was determined by measurement of the absorbance at 256 nm (Hitachi 101 Spectrophotometer) after appropriate dilution of each fraction with water. No interference due to the presence of C12E6 was 0 1980 American Chemical Society
