Fluorocarbon and Hydrocarbon Short-Chain Nonionic Amphiphiles: A

J. Phys. Chem. B 1998, 102, 10917-10924

10917

Fluorocarbon and Hydrocarbon Short-Chain Nonionic Amphiphiles: A Comparative Study of Their Behavior in Aqueous Medium Christine Damas, Re´ gine Naejus, and Robert Coudert* Laboratoire de Physicochimie des Interfaces et des Milieux Re´ actionnels, PIMIR EA 2098, De´ partement de Chimie, UniVersite´ Franc¸ ois Rabelais, AVenue Monge, Parc de Grandmont, F-37200 Tours, France

Ce´ line Frochot,† Alain Brembilla,‡ and Marie-Laure Viriot† De´ partement de Chimie Physique des Re´ actions, UMR 7630 CNRS-INPL, and Laboratoire de Chimie Physique Macromole´ culaire, UMR 7568 CNRS-INPL, Groupe ENSIC, 1 rue GrandVille, BP 451, F-54001 Nancy Cedex, France ReceiVed: July 17, 1998; In Final Form: October 15, 1998

The aqueous solution properties and the micellar structure of two short-chain nonionic surfactants containing a hydrocarbon tail, 1,2-hexanediol (HD), and a perfluorinated tail, 3,3,4,4,5,5,6,6,6-nonafluoro-1,2-hexanediol (PFHD), have been compared by using various techniques such as pyrene fluorescence spectroscopy, vapor pressure osmometry, tensiometry, and dye solubilization. The aggregational behavior of both systems in aqueous medium has been evidenced by the polarity decrease of the pyrene microenvironment with increasing surfactant concentration. The binding coefficient of pyrene with the aggregates was calculated by application of the phase-separation model to the pyrene fluorescence results. The aggregation numbers of the HD (NH) and PFHD (NPF) micelles have been evaluated by application of the phase-separation and the mass-action law models to the osmotic coefficients measurements. The NH value (26 ( 8), which is in good accordance with previous experimental results (30 ( 10), is higher than NPF (15 ( 1). Both compounds exhibit surface-active properties with a maximum surface tension lowering of 42 and 57 mN m-1 for HD and its perfluorinated homologous compound, respectively. Their solubilizing power toward Orange OT was compared. Critical micelle concentrations (CMCs) have been determined in the temperature range 20-50 °C (30-50 °C for PFHD insoluble below 30 °C), and thermodynamic parameters such as standard enthalpy and entropy changes for micellization have been calculated.

Introduction Amphiphiles, which have been extensively studied for decades, offer a wide variety of physicochemical properties since their structure can be modified at will. As it can be seen effectively, covalent association of the hydrophobic part characterized by the presence of hydrocarbon, halogenated, siliceous, linear, branched, saturated, or unsaturated groups and of the hydrophilic part (anionic, cationic, amphoteric, or nonionic) leads to various structures going from simple molecules to globular systems. Among them, linear perfluorocarbon surfactants are of great interest. Because of their ability to lower considerably the surface tension of aqueous solutions and their high chemical and thermal stability, fluorinated surfactants offer a wide range of application fields.1 Studies of the aqueous medium behavior of such compounds are mainly concerned with anionic perfluorinated surfactants having more than seven carbon atoms in their hydrophobic tail.2 However, few studies have been reported for perfluorinated short-chain surfactants.3 Hence, the effect of chain length on the critical micelle concentrations (CMCs) of perfluorinated carboxylic acids with more than two carbon atoms has been previously investigated3 to show the linear CMC decrease with the increase in the chain length. * To whom correspondence should be addressed. Fax: (33) 02 47 36 69 60. E-mail: [email protected]. † Department de Chimie Physique des Re ´ actions. ‡ Laboratoire de Chimie Physique Mocromole ´ culaire.

Nevertheless, to the best of our knowledge, no deep investigation on the aggregational properties of the shortest compounds has been undertaken. Thus, the goal of this paper is to deal with some original thermodynamical properties that heretofore have not been systematically studied, in continuation of the study of the micellar behavior of short alkyl chain surfactants developed in our laboratory.4-6 In the present work, we examine the aqueous solution properties of a short-chain nonionic perfluorinated compound containing four carbon atoms in its hydrophobic part 3,3,4,4,5,5,6,6,6-nonafluoro-1,2-hexanediol (PFHD) in comparison with those of its homologous hydrocarbon compound. As shown in a previous paper,4 1,2-hexanediol (HD) behaves similarly to nonionic surfactants in aqueous medium. The existence of micelle-like aggregates has been evidenced, and they form over a wide range of concentrations (0.52-0.76 mol L-1). The aggregation number (20) is lower than those (in the range 50-100) of most conventional surfactants but is comparable to those reported for short-chain ionic ones.7,8 The CMCs of both compounds have been determined by using pyrene fluorescence spectroscopy, vapor pressure osmometry, surface tension measurements, and dye solubilization. The polarity dependence of vibrational band intensities in pyrene monomer fluorescence and vapor pressure osmometry were used to investigate the structure of micelles. Their surface activity was examined and their solubilizing power toward Orange OT was evaluated. In the last part of this work,

10.1021/jp9830777 CCC: $15.00 © 1998 American Chemical Society Published on Web 12/04/1998

10918 J. Phys. Chem. B, Vol. 102, No. 52, 1998

Damas et al.

the influence of temperature on their micellar behavior was investigated. Thermodynamic properties such as molar adsorption Gibbs energy and molar enthalpy and entropy changes for micellization have thus been determined and compared. Experimental Section Materials. 1,2-Hexanediol (HD), an Aldrich product (with a purity higher than 98%), was distilled under reduced pressure (77 °C/4 mmHg) before use. 3,3,4,4,5,5,6,6,6-Nonafluoro-1,2-hexanediol [CF3-(CF2)3CH(OH)-CH2OH] (PFHD) was synthesized from 3,3,4,4,5,5, 6,6,6-nonafluoro-1-hexene (Aldrich, 99%) by one of the procedures reported previously.9 After oxidation of perfluorohexene by hexadecyltrimethylammonium permanganate in methylene chloride at 0 °C for 8 h and then at room temperature for 15 h, diethyl ether was first added. Second, water was dropped until dark manganese oxide precipitated. Then, the same volume of water was added again to complete the hydrolysis reaction. The precipitates were washed several times with diethyl ether. After removal of solvents, the product was purified by three recristallizations from chloroform at -10 °C. The compound structure and the purity were confirmed by 1H NMR spectroscopy and elemental analysis, respectively. 1-(o-tolylazo)-2-naphthol (Orange OT, C. I. 12100, Aldrich) and pyrene (Reference CRM 177, Commission Europe´enne, Centre Commun de Recherche, Institut des Mate´riaux et des Mesures de Re´fe´rences IRMM) were used as received. Methods. Fluorescence Spectroscopy. Fluorescence spectra were recorded on a Spex fluorolog-2 spectrometer equipped with a thermostatically controlled cell (to (1 °C). The fluorescence emission spectra of pyrene (0.6 × 10-6 mol L-1) were obtained by exciting the samples at 332 nm (excitation slit width ) 0.5 mm, emission slit width ) 0.5 mm, ∆λ1/2 ) 2 nm). The spectra were used to determine the ratios (I1/I3) of the fluorescence emission intensities of the first (I1) and third (I3) vibrational peaks of monomeric pyrene. Experiments were performed at 30, 40, and 50 °C for HD and at 50 °C for PFHD. Vapor Pressure Osmometry. Vapor pressure osmometry measurements were carried out using a KNAUER apparatus kept at 40 ( 0.1 °C. This type of osmometer and its operating principles have been described previously.5,10 The CMC was evaluated by plotting the osmotic coefficient φ as a function of the inverse of molality (1/m in mol kg-1) at fixed temperature. The osmotic coefficient for a nonionic solute solution was calculated by the following relation:

φ ) k∆R/m where m is the molality of the solute, k is the calibration constant, and ∆R is the resistance difference between two thermistors in a water-saturated atmosphere at the experimental temperature, previously calibrated with water, when a drop of water was placed on one thermistor and a drop of solution was placed on the other one (k ) (443.3 ( 1.1) × 10-5 ohm-1 mol kg-1). Measurement of Surface Tension. The surface tension was measured by using a LAUDA tensiometer in the temperature range 30-50 °C ((0.5 °C). Measurements were done with the ring method. Solubilization of Orange OT. To prepare the saturated Orange OT solutions, excess crystalline Orange OT was added to surfactant aqueous solutions. The PFHD solutions were sonicated for a few seconds and stirred. The sample solutions were allowed to equilibrate for 1 week at 20 °C ((0.1 °C) for HD

Figure 1. Variations of the I1/I3 ratio calculated from the pyrene fluorescence spectrum (pyrene ) 0.6 × 10-6 mol L-1) as a function of HD (2) and PFHD (b) concentration C (C in mol L-1) at 50 °C. The solid lines are the calculated curves obtained by fitting eq 3 to the data points. The straight dotted lines are drawn by connection of data points at low and high concentrations (first and second plateaus) and at concentrations around the CMC.

and at 30 °C ((0.1 °C) for PFHD. They were filtered, and their absorbency was measured at 500 nm (molar extinction coefficient  ) 17 400 L mol-1 cm-1). Results and Discussion 1. Aqueous Micellar Structure of PFHD in Comparison with HD. As a first part of this work, the micellar structure of both systems has been examined by using pyrene fluorescence and vapor pressure osmometry. The fluorescence properties of pyrene have been extensively examined, and many works have been devoted to its use to study micellar systems or amphiphilic polymers.11,12 The fluorescence spectrum of pyrene depends strongly on both solvent polarity and microenvironment. Hence, correlations between the I1/I3 ratio and the empirical polarity parameters ET(30) defined by Dimroth and Reichardt13 have been established previously.14 As a hydrophobic probe, pyrene is preferentially solubilized in hydrophobic medium. It can thus detect the formation of hydrophobic aggregates in polar media. Figure 1 shows the variations of I1/I3 as a function of HD and PFHD concentrations C at 50 °C. Both plots present three parts: first a plateau corresponding to the high value of I1/I3 (1.80, according to our apparatus function) due to pyrene in water, second an abrupt decrease owing to the beginning of the formation of micelles, and third a second plateau whose low I1/I3 values (1.05 for HD and 1.35 for PFHD) reflect the less polar microenvironment sensed by pyrene in the micellar aggregates. As previously reported,15,16 the surfactant concentration where I1/I3 has dropped to a low and a nearly constant value marks the CMC. Thus, the CMC values, which are reported in Table 1 and noted CMCc, have been determined from the intercept of the straight lines obtained where the I1/I3 ratio decreases and remains approximately constant at the second plateau as shown in Figure 1. The concentrations corresponding to the onset of micelle formation and at the inflection point of the I1/I3 curve have been also listed in Table 1 (noted CMCa and CMCb, respectively) together with literature data for HD. The CMCs have also been determined analytically by application of the phase-separation model. Calculated values are listed in Table 2. Supposing that the global polarity sensed by pyrene is the mean polarity of pyrene associated with micelles (m) and of pyrene dispersed in the bulk phase (w), the I1/I3 ratio which depends linearly on

Short-Chain Nonionic Amphiphiles in Aqueous Medium

J. Phys. Chem. B, Vol. 102, No. 52, 1998 10919

TABLE 1: Critical Micelle Concentrations for HD and PFHD Determined from Fluorimetry and Tensiometry in the Temperature Range 25-50 °C 1,2-hexanediol (HD) CMC (mol L-1) T (°C) 25 30 35 40 50

PFHD CMC (10-2 mol L-1)

fluorimetry

tensiometry

0.47a 0.61b 0.76c

0.58d 0.58

0.47a 0.60b 0.73c 0.45a 0.59b 0.74c

0.51 0.49

fluorimetry

tensiometry

0.93a 1.41b 1.89c

1.82 1.86 1.97 2.15

a Determined graphically at the intercept between the first plateau and the linear part of the I1/I3 vs log C decrease. b Determined at the inflection point. c Determined graphically at the intercept of the linear decrease of I1/I3 vs log C and the second plateau. d Extrapolated from ref 4.

TABLE 2: Parameters Calculated by Application of the Phase-separation and/or the Mass-Action Models to Fluorescence and Vapor Pressure Osmometry Measurements on HD and PFHD Aqueous Solutionsd HD

PFHD

Phase-Separation Model Applied to Fluorimetry at 50 °C CMC (mol L-1) 0.46a 0.97 × 10-2 K/N 10a 600 Phase-Separation Model Applied to Osmotic Coefficients at 40 °C n 18.5 ( 0.8 13.7 ( 0.8 CMCb (mol kg-1) 0.52 4.1 × 10-3 CMCc (mol kg-1) 0.72 2.13 × 10-2 Mass-Action Model Applied to Osmotic Coefficients at 40 °C n 34 ( 2 16 ( 3 AI -0.42 ( 0.03 -44.5 ( 5.7 ∆n1 1.4 ( 0.8 0(3 ∆n2 2.5 ( 2.0 0(4 a Calculated from eq 3. b Calculated from eq 5. c Determined at the break point of φ vs 1/m. d The differences (∆n1 and ∆n2) between the calculated aggregation numbers (n′ and n′′) and the initial n value are indicated.

polarity, can be thus expressed as follows

I1/I3 ) Pm/Pt(I1/I3)m + Pw/Pt(I1/I3)w

(1)

where Pt is the total pyrene concentration (Pt ) Pm + Pw). Furthermore, pyrene is distributed between the two phases according to the following binding coefficient:

K ) Pm/[PwM]

(2)

where M is the aggregate concentration equal to M ) (C CMC)/N, where N is the aggregation number. Combining eqs 1 and 2 leads to17

I1/I3 ) (I1/I3)w + KN-1(C - CMC)/ [1 + KN-1(C - CMC)][(I1/I3)m - (I1/I3)w] (3) K/N and CMC values were obtained by a parameter adjustement using eq 3. Figure 1 plainly shows the calculated curves of I1/ I3 versus log C after micellization, which become identical to the experimental values. Table 1 reports the experimental CMC values determined by plotting I1/I3 versus log C, as in the above cited studies, at 50 °C. Notice that, for HD as well as PFHD, the values of the CMCs obtained graphically are about 2 times higher than the calculated ones from eq 3 (see Table 2), i.e., with the phase-separation model treatment. For such short-chain surfactants with fairly small micelles, one would expect that the above model is inadequate. Nevertheless, if we consider

the onset of micelle formation, i.e., at the first break of the I1/I3 plot versus log C, which occurs at 0.45 mol L-1 and 0.93 × 10-2 mol L-1 for HD and PFHD, respectively, there is now a good agreement. In our study, the choice of the CMC is not unambiguous since the change in slope occurs over a wide range of concentrations rather than at a definite concentration. Then, as previously reported5,18 for short-chain nonionic surfactants, the CMC can be probably defined as the inflection point in the micellar region of the corresponding graph. Finally, the CMCs thus calculated are at about 0.60 mol L-1 and 1.4 × 10-2 mol L-1 for HD and PFHD, respectively, at 50 °C. Although the HD micellar properties at 50 °C are not described in the literature, our results are in good accordance with the reported values, which range between 0.73 and 0.84 mol L-1 at 8 °C,15 0.6015 and 0.734 mol L-1 at 25 °C, and 0.61 mol L-1 at 40 °C.15 To our knowledge, there are no reports of PFHD in the literature. The hydrocarbon HD exhibits a much larger CMC (with a factor of 44) than its homologous perfluorinated compound. Similar observations have been noticed in the case of carboxylate salt surfactants.2a The perfluorination effect on CMC can be expressed in terms of molar free energy for micellization per CH2 or CF2 group

∆Gm (CH2 or CF2) ) ωRT

(4)

Kunieda and Shinoda2a obtained ω ) -2.21RT for n-C7F15COOK. This value is about twice the ω value (-1.08RT) found for the corresponding hydrocarbon surfactant.19 Furthermore, the CMC of PFHD with a hydrophobic chain length of four carbon atoms is on the same magnitude as the CMC of perfluorinated carboxylate salts bearing eight carbon atoms in their hydrophobic tail. This can be attributed to the less hydrophilic character of the polar diol headgroup than the carboxylate group. The diol headgroup thus favors early micellization (owing to the absence of Coulombic interactions). Despite the presence of the same hydrophilic head in both compounds (HD and PFHD), their micellar behavior is strongly affected by the chemical nature of their hydrophobic tail as is evidenced by pyrene fluorescence. First of all, the K/N ratio is found to be lower for HD than for its homologous perfuorinated compound. Moreover, with an aggregation number equal to 20 in HD micelles,4 the pyrene binding coefficient is around 200, whereas the K value for PFHD lies over 600 (with N > 1). This indicates a higher affinity for pyrene toward perfluorinated micelles than hydrocarbon ones, due to the stronger hydrophobicity of the perfluorinated tail. However, the polarity sensed by pyrene is higher in perfluorinated aggregates than in the HD micelles. This apparent discrepancy can be attributed to great differences in micellar structures. Using the polarity scale from I1/I3 versus ET(30),14 the polarity sensed by pyrene in HD micelles corresponds to an ET(30) value of 47, whereas the ET(30) value is equal to 54 for perfluorinated aggregates. This may result from the structure of fluorocarbon chains. The higher bulkiness of fluorine atoms than hydrogen atoms tends to make the fluorocarbon chains more rigid than the hydrocarbon chains. This could thus inhibit the incorporation of the bulky hydrophobic probe into the core of the fluorocarbon aggregates and could also favor water penetration. Hence, this observation suggests that an appreciable water contact of the fluorocarbon chains would exist at the solubilization site of pyrene in the perfluorinated aggregates and that pyrene would be close to the periphery of the aggregates whereas it would be located more inside the hydrocarbon micelles. Therefore, pyrene association with perfluorinated micelles does not seem to be affected by

10920 J. Phys. Chem. B, Vol. 102, No. 52, 1998

Damas et al. by Desnoyers et al.23 for alkyldimethylamine oxides with short alkyl chains, has been used in our study. It can be formulated by the following equations where the nonionic monomers M associate into a micelle Mn.

nM a Mn

(6)

The thermodynamic equilibrium constant for (6) is

K) Figure 2. Plots of the osmotic coefficient φ versus the inverse of the molality, m, at 40 °C for HD (b) and PFHD (O).

the hydrated microenvironment. Similar observations concerning the polar microenvironment have been noticed in the case of other fluorinated aggregates.20 Calculation of the osmotic coefficients by using vapor pressure osmometry measurements leads to useful information to describe thermodynamic properties and is often used to determine other parameters such as aggregation numbers in the case of self-associative systems.6,21-23 The phase-separation model generally applies very well for nonionic surfactants. Beyond the CMC, the monomer concentration is assumed to remain constant and micellar aggregation number is n. The osmotic coefficient φ is thus given by23

1 1 1 φ ) + 1 - CMC n n m

(

)

()

(5)

where m is the surfactant molality, and the plots of φ versus 1/m would give n and the CMC. Application of this model leads to n values of 13.7 ( 0.8 and 18.5 ( 0.8 for PFHD and HD respectively, and the CMC values calculated from the slope are 4.1 × 10-3 mol kg-1 for PFHD and 0.52 mol kg-1 (0.485 mol L-1) for HD at 40 °C (see Figure 2). From these results, we can notice first the close value of n to the number obtained from static light-scattering measurements on HD aqueous solutions (n ) 20 at 20 °C),4 according to a slightly temperatureindependent aggregation number, as for other nonionic surfactants, and second the CMC of HD, which is on the same order of magnitude as the CMC determined by pyrene fluorescence spectrometry. Note that the break point at 0.72 mol kg-1 (0.657 mol L-1) for φ vs 1/m is quite close to the CMC value reported in Table 1 for HD. However, the results obtained in the case of PFHD and especially the calculated CMC, which differs by a factor 2-5 from the CMC determined by pyrene fluorescence measurements or by tensiometry (see Table 1), suggest the low validity of the phase-separation model for the perfluorinated compound in contrast to the hydrocarbon surfactant. It should also be pointed out that the concentration of PFHD (2.13 × 10-2 mol kg-1) determined at the break point for the φ versus 1/m curve is in convergence with the CMC value obtained from other methods. Although the phase-separation model has proved to be useful in defining the CMC and in deriving thermodynamic functions of micellization, it is clear that large deviations from this model are noted in the case of PFHD, according to the results found for other short-chain surfactants.23 An alternative to the phase-separation model is the use of the mass-action law model, which has been used by numerous authors to investigate the micellization process.24,25 The multipleequilibrium model extended by Mukerjee26 involving a stepwise aggregation and the model based on a one-step association process are thermodynamically equivalent. This latter approach of the mass-action model, which has been avantageously used

(m(1 - R)/n)γMn (Rm)n(γM)n

(7)

in which R represents the fraction of free monomers at the experimental molality m. The activity coefficients (γMn and γM) are taken as unity at all temperatures. The osmolality mφ, which is related to the water activity a1 by

103 mφ ) - ln a1 M1

(8)

can be also expressed as the sum of all the osmolalities of each species constituting the solution

mφ ) RmφM +

m(1 - R) φMn n

(9)

Assuming an ideal behavior of the micelles, φMn is taken as unity. φM is expressed as

φM ) 1 + AIRm

(10)

where AI is the second virial coefficient related to pairwise monomer interactions. Therefore, φ is expressed as

φ ) R(1 + AIRm) +

1-R n

(11)

and

(

(1 - φ)m ) 1 -

1 [(1 - R)m] - AI(Rm)2 n

)

(12)

Since mi and Ri represent respectively the total surfactant molality and the fraction of free monomer at the CMC, the approximated expression of the constant K becomes

K)

(1 - R)m n

)

(1 - Ri)mi

n(Rm)

n(Rimi)n

(13)

where Ri depends on the aggregation number n according to23

Ri(n) )

n (1 - (2n)-1/2) (n - 1)

(14)

The following equation is then obtained by combining eqs 12 and 13

(

(1 - φ)m ) 1 -

1 mi - xi,n n (x ) - AI(xn)2 n (x )n n

)

(15)

i,n

where xi,n ) Ri(n)mi and xn ) Rm. In this expression, mi corresponds to the molal CMC and is given by the break point of the φ vs 1/m plot.

Short-Chain Nonionic Amphiphiles in Aqueous Medium

J. Phys. Chem. B, Vol. 102, No. 52, 1998 10921

Figure 3. Plots of m(1 - φ) versus molality, m, of HD (b) and PFHD (O) from vapor pressure osmometry measurements at 40 °C. The solid lines are the curves calculated by fitting eq 15 to the data points.

The xi,n and xn parameters depend on n. Thus, for a given n number, the xn values can be evaluated by using eq 13 or eq 16, which is equivalent to eq 13, at each experimental m value

mi - xi,n

m ) xn + (xn)n

(xi,n)n

(16)

With the molality m being plotted as a function of xn at a fixed n value, an xn value can thus be determined graphically or numerically for a given experimental m value. Equation 15 is then solved by successive iterations. Starting from initial values of n0, xi,n0, and AI0, a parameters adjustement is applied to eq 15 and a new set of parameters n′, xi,n′, and (AI)′ is obtained. One calculates first the difference ∆n1 ) (n′ - n0). Second, from xi,n′ and eq 14, Ri and its corresponding n′′ value are determined and the second difference ∆n2 ) (n′′ - n0) is calculated. The calculation is repeated until the differences ∆n1 and ∆n2 tend to zero as much as possible. When (1 - φ)m is plotted as a function of m (see Figure 3), we can notice a good accordance between the calculated curves by using eq 15 with the n and AI parameters obtained after iterations and the experimental points. Values of micellar aggregation numbers obtained by application of this model to vapor pressure osmometry measurements on HD and PFHD aqueous solutions are summarized in Table 2. The aggregation numbers of both systems are higher than the n values obtained by application of the phase-separation model. In the case of PFHD, this difference (of approximately 2) is not as significant as that of the n value of HD, whose micellar aggregation number is multiplied by a factor 1.8 as the massaction model is used instead of the phase-separation model. It must also be mentioned that similar aggregation numbers (40 ( 10 from light-scattering measurements and 35 ( 4 from vapor pressure measurements at 25 °C) have also been obtained15 for HD. This suggests a large uncertainty about the determination of HD micellar aggregation number, which depends not only on the model used but also on the physical property being measured. It would then lie in the range 20-40. Undoubtedly, we can conclude that the n value of PFHD micelles is lower than the n value of its hydrocarbon homologue. This may be attributed to the higher stiffness of the perfluorinated chain than the hydrocarboned one, which would be induced by the larger size of the fluorine atoms in comparison with the hydrogen atoms. This result reinforces the observations, made from pyrene fluorescence measurements, about the difference in the micellar micropolarity of the two systems. The lower micellar aggregation number of PFHD having the same tail length as HD would effectively favor water penetration in

Figure 4. Plots of the surface tension γ versus ln C at 50 °C for HD (b) and PFHD (O) aqueous solutions.

the micelle and would thus explain the higher micropolarity observed in PFHD micelles. 2. Surface-Active Properties. Figure 4 shows the concentration dependence of the surface tension at 50 °C (the plots in the temperature range 30-50 °C are not shown) for aqueous solutions of both hydrocarboned and perfluorinated compounds. There is a clear-cut break in each curve, showing that HD and PFHD self-associate in a manner similar to that of conventional surfactants. In comparison with those of HD, PFHD solutions exhibit much lower surface tension values. The minimum surface tension reached by the aqueous HD solutions is approximately 27-28 mN m-1 in the temperature range 3050 °C, whereas that of aqueous solutions of PFHD is as low as 15 mN m-1 approximately in the same temperature range. The values of CMC are estimated from the intercepts of the linear parts of γ vs ln C curves. The CMC values thus obtained are reported in Table 1 at various temperatures. In the case of HD, the CMC slightly decreases upon temperature as for polyoxyethylene n-alkyl alcohols (CmEn),27 whereas the phenomenon is inverted for PFHD. In comparison with the CMC values determined from pyrene fluorescence measurements for HD, the CMC values obtained from surface tension experiments are quite close to the CMC values determined at the inflection point in the I1/I3 vs log C plots. Despite the presence of pyrene, which may disturb the micellar properties, the fluorescence method appears to be more sensitive to self-association of HD (due to the I1/I3 decrease starting at a surfactant concentration lower than the CMC) than the surface tension one, which gives a mean value of the CMC. In the case of PFHD, the CMCs interpretation is completely different from that of HD. At 50 °C, the CMC obtained from tensiometry measurements (Table 1) is effectively on the same order of magnitude as that determined at the second break of the I1/I3 vs log C curve. Thus, unlike that for HD, the CMC value of PFHD from tensiometry would characterize the complete micellization. In addition to the determination of the CMC values and the minimum surface tension achievable, surface tension curves yield useful information such as the adsorption amount of surfactant at the air/water interface (or surface excess concentration) and the free energy of adsorption. The surface adsorption amount Γ is defined by the Gibbs equation written as follows for nonionic compounds:

1 dγ Γ)RTd ln C

(17)

where R is the gas constant (8.314 J mol-1 K-1) and T is the absolute temperature. The Γ values are thus calculated from the slope of the linear portion of the γ vs ln C curves. Under conditions of surface saturation, the area A occupied by a molecule can be calculated from the surface excess concentration

10922 J. Phys. Chem. B, Vol. 102, No. 52, 1998

Damas et al. TABLE 3: Molecular Areas AG and AS of HD and PFHD at the Air/Water Interface Calculated by Using Gibbs Equation and Szyszkowski’s Equation, Respectively, and Free Energies of Adsorption (∆G°ads) Obtained from Szyszkowski’s Equation HD

PFHD

T AG (Å2/ AS (Å2/ ∆G°ads AS (Å2/ ∆G°ads AG (Å2/ (°C) molecule) molecule) (kJ mol-1) molecule) molecule) (kJ mol-1) 30 35 40 50

Figure 5. Plots of surface tension for HD aqueous solutions (b) and PFHD aqueous solutions (O) versus concentration (C) at 50 °C, with scale expansion in the inset. The solid lines are calculated by fitting Szyszkowski’s equation to the experimental points. The dotted line is drawn by connection of data points above the CMC.

of the saturated surface

A ) 1/(NAΓ)

(18)

where NA is Avogadro’s number. Another quantitative treatment of the surface tension curves is the use of the following Szyszkowski’s equation, which arises from the combination of the Langmuir adsorption equation with the Gibbs adsorption equation.

(Cβ + 1)

γ0 - γ ) R log

(19)

where γ0 is the solvent surface tension and the constants R and β are related respectively to the maximum surface excess saturation Γm and to the standard free energy of adsorption ∆G°ads by the expressions

R ) 2.303RTΓm

(

β ) W exp

(20)

)

∆G°ads RT

(21)

W is the number of moles of water per liter of water at temperature T. The values of R and β are calculated by a parameters adjustement of Szyszkowski’s equation. The calculated curves, which are shown for HD and PFHD at 50 °C as an exemple in Figure 5, are in good agreement with the experimental points. The free energies of adsorption are thus obtained from eq 21, and the areas covered by the surfactant molecule adsorbed at the air/water interface are calculated by combining eqs 18 and 20. Their values are listed in Table 3 together with the molecular areas obtained from the Gibbs equation. The areas calculated by using Gibbs equation (AG) are on the same order of magnitude as the areas obtained from Szyszkowski’s equation (AS) (within 1-8%). The influence of a perfluorinated chain on the A values does not appear significantly. Thus, the diol polar head seems to have a preponderant rule on the way the molecules are distributed at the air/water interface. The effect of the polar headgroup on adsorption at the air/water interface has been examined in the case of fluorinated nonionic surfactants having a two-chain polyoxyethylene hydrophilic head.28 The area per surfactant

39.5

39.1

-21.6

39.9 41.6

38.4 40.8

-21.8 -22.5

41.6 41.6 42.2 45.4

42.2 42.0 42.6 44.8

-33.5 -33.8 -33.9 -34.6

molecule has been studied as a function of the number of oxyethylene units in the hydrophilic head and was compared with the area occupied by the corresponding hydrocarbon-based surfactant. It is shown that, for nonionic surfactants with a number of oxyethylene units larger than six, the A values are similar for fluorinated or nonfluorinated surfactants whereas the fluorination of the hydrophobic tail tends to reduce the molecular area as the hydrophilic group becomes smaller. For the present compounds, the nonionic diol headgroup has thus a greater influence than the hydrophobic part on their adsorption at the air/water interface. Although the tendency of a surfactant to adsorb at the liquid/ air interface parallels, in some cases, its tendency to form micelles, these two phenomena do not necessarily run parallel to each other. Then, for the compounds studied in the present paper, the influence of fluorination first on micellization is examined through the CMC values (Table 1) and second on adsorption is examined through the free energies of adsorption ∆G°ads listed in Table 3. The CMCs for complete micellization of an HD range from 0.49 to 0.58 mol L-1 and those for PFHD lie between 1.82 × 10-2 to 2.15 × 10-2 mol L-1 in the temperature range 30-50 °C. The corresponding standard Gibbs energies of micellization are obtained using the following equation27

∆G°m ) RT ln(CMC/W)

(22)

The calculated values are between -11.3 and -12.7 kJ mol-1 for HD and between -20.6 and -21.7 kJ mol-1 for PFHD, while the ∆G°ads values of HD and PFHD are comprised between -21.6 and -22.5 kJ mol-1, and -33.5 and -34.6 kJ mol-1 respectively. The fluorination of the aliphatic chain promotes both micellization and adsorption at the air/water interface, as is evidenced by the increase in the negative values of ∆G°ads and ∆G°m. However, the negative increase in ∆G°m is smaller (-9.1 ( 0.1 kJ mol-1) than that for ∆G°ads (-12.0 ( 0.1 kJ mol-1). Thus, the perfluorinated compound shows a better ability toward adsorption than micellization. It can be attributed to a steric hindrance due to the bulkiness of the perfluorinated chain, which appears to have more influence on micellization than on adsorption. 3. Solubilization of Orange OT. Solubilization is of great theoretical as well as practical interest (in drug delivery systems, detergency, and tertiary oil recovery methods). However, to the best of our knowledge, only a few studies are concerned with solubilization by fluorinated surfactants.2a,29 The maximum amount of a solubilizate that can be dissolved in micelles, its location, the interactions involved, and the effect of the solubilizate on the micellar properties depend closely on the structures of the surfactant and of the solubilizate. As part of this work, we have focused our interest on the effect of perfluorination of the HD hydrophobic tail on the amount of Orange OT, a hydrophobic dye, solubilized in

Short-Chain Nonionic Amphiphiles in Aqueous Medium

J. Phys. Chem. B, Vol. 102, No. 52, 1998 10923 TABLE 4: CMC (in mol L-1) and Thermodynamic Parameters of Micellization for HD and PFHD Such as the Free Energy for Micellization (∆G°m in kJ mol-1) and the Entropic Contribution (T∆S°m in kJ mol-1) HD T (°C)

CMC

∆G°m ((0.01)

T∆S°m ((1.1)

20

0.76a 0.715a 0.68 0.75a 0.74a 0.75a 0.76

-10.46 -10.61 -10.73 -10.56 -10.70 -10.67 -10.81

10.8 10.9 11.0 10.9 11.0 11.0 11.1

0.71a 0.66 0.73 0.74

-11.15 -11.52 -11.26 -11.57

11.5 11.8 11.6 11.9

22 25 30

Figure 6. Absorbency of Orange OT as a function of HD (b) or PFHD (O) surfactant concentration, with scale expansion in the inset. A at 500 nm is monitored for both the HD and PFHD solutions.

aqueous surfactant solutions. Many studies have been undertaken about the solubilization of Orange OT in various hydrocarbon surfactants.30 Such a dye is only slightly soluble in water and is most likely solubilized in the hydrocarbon core of the micelles. Figure 6 shows the plots of Orange OT absorbency at 500 nm versus surfactant concentration for HD at 20 °C and PFHD at 30 °C (PFHD was not sufficiently soluble, and values at 20 °C were not available). In both systems, a linear increase of the absorbency (A) with an increase of the surfactant concentration is observed. The equation of these straight lines can be represented by30

A ) b(C - CMC)

35 40 50 a

CMC

∆G°m ((0.01)

T∆S°m ((1.1)

1.82 × 10-2 2.23 × 10-2 1.86 × 10-2 1.97 × 10-2 2.13 × 10-2 2.15 × 10-2 1.89 × 10-2

-20.21 -19.70 -20.49 -20.66 -20.46 -21.08 -21.43

14.1 13.6 14.4 14.6 14.4 15.0 15.3

From the CMC in ref 4.

(23)

where b is the slope parameter. The CMC of HD or PFHD is thus determined by extrapolating the linear portion to zero absorbency (0.68 mol L-1 for HD and 2.23 × 10-2 mol L-1 for PFHD). They are close to the CMCs given by the other methods used throughout this work. The amount of dye S solubilized by the micelles can be estimated from the slope b of eq 23.

S ) b/(l)

PFHD

(24)

where  is the molar extinction coefficient of Orange OT ( ) 17 400 L mol-1 cm-1) and l is the cell length (1 cm). The values are 8 × 10-4 and 2 × 10-4 for HD and PFHD, respectively, and correspond to the number of Orange OT molecules solubilized per micellized surfactant molecule. These results show the higher ability of HD to solubilize Orange OT than its perfluorinated homologous compound. Similar observations have been noticed in previous works.31,32 Thus, the solubilizing power of ammonium perfluorooctanoate solutions and of sodium octanoate solutions toward octanol were compared, and the lower solubility of octanol in the perfluorinated surfactant solutions was attributed to a low mutual solubility of fluorocarbon and hydrocarbon chains. Similarly, the mutual phobicity between fluorocarbon and hydrocarbon groups was the explanation given for the reduction of Orange OT solubilization by the terminal perfluorination of dodecyltrimethylammonium bromide. Hence, it can also explain the differences in the solubilization of Orange OT by HD and PFHD. However, despite the reduction in the solubilizing power by the perfluorinated chains, it should be noted that the low CMC of PFHD enables the solubilization of Orange OT in small quantities but at small surfactant concentrations (lower than 0.68 mol L-1), which is impossible with HD solutions.

Figure 7. Variations of the free energy of micellization ∆G°m with temperature T for HD (+, this work; [ from ref 4) and PFHD (- - -), linear regression of all the points; s, linear regression of the points from tensiometry).

4. Influence of Temperature on the Micellar Behavior of PFHD and HD Diols. The data presented in Tables 1 and 2 on CMC include a certain amount of information such as the effect of temperature on CMC. Data on CMC as a function of temperature may be used to calculate heats and entropies of micelle formation. The standard free energies for HD and PFHD micelle formation (∆G°m) are calculated from eq 22 and are presented in Table 4. The ∆G°m values for HD are less negative than those of its perfluorinated homologue by about 10 kJ mol-1. This suggests less of an ability toward micellization exhibited by HD than PFHD, attributed to the higher hydrophobicity of the fluorocarbon chain relative to the hydrocarbon chain. In addition, ∆G°m is plotted against the temperature T as shown in Figure 7. In the case of HD, the ∆G°m values calculated from literature data4 are presented with the ∆G°m values obtained in the present study on the same plot. These two series of ∆G°m values obey quite well the same linear variation with the temperature. For PFHD, a quite linear variation of ∆G°m with T is also observed. This is particularly clear for data points obtained from surface tension measurements. Since these ∆G°m vs T curves are nearly linear, the corresponding ∆H°m and ∆S°m can be calculated from the following equation:

∆G°m ) ∆H°m - T∆S°m

(25)

The ∆H°m and ∆S°m values of HD are respectively 0.3 ( 1.1 kJ mol-1 and 36.9 ( 3.7 J mol-1 K-1. The ∆H°m and ∆S°m values of PFHD are -6.1 ( 1.1 kJ mol-1 and 46.6 ( 3.6 J

10924 J. Phys. Chem. B, Vol. 102, No. 52, 1998 mol-1 K-1 when tensiometry data points are considered separately and become equal to 0.7 ( 4.0 kJ mol-1 and 68 ( 13 J mol-1 K-1 respectively, if all the points are taken into account. The more reliable thermodynamic values, which are determined from tensiometry measurements, will be only taken into account for further discussion. The slightly endothermic behavior of HD micellization suggests weaker chain-solvent interactions than chain-chain interactions in the micelles. Thus, an increase in temperature favors micellization, and the CMC should decrease. In the case of PFHD, the opposite tendency is observed and an increase in temperature should lead to a decrease of the CMC. However, the effect of temperature on the CMC of HD and PFHD is hardly noticeable. This can be attributed to the higher contribution of the entropic term T∆S°m than enthalpy to the free energy for micellization ∆G°m. Micellization involves the transfer of the hydrophobic chain of the surfactant from an aqueous environment to the micelle, and the positive entropy change indicates an increase in randomness, which is unlikely to result from aggregation of surfactant monomers. This probably arises from the destruction of “iceberg” structured water around the hydrophobic chain. Moreover, the higher entropy of PFHD than that of its hydrocarbon homologue is attributed to higher hydrophobic interactions. As it is suggested in ref 33, the fluorocarbon chain has a higher water-structure-promoting ability than a hydrocarbon chain. Hence, the hydrogen-bonded water structure around fluorocarbon chains is more extensive, and micellization requires breaking more water structure around a fluorinated surfactant chain than around the corresponding hydrocarbon one. This would thus lead to a higher entropy for micellization of PFHD than for micellization of the hydrocarbon homologue. Conclusion The results presented in this paper clearly evidence the aggregation ability in water of a short perfluorinated chain nonionic surfactant PFHD, similar to that of its hydrocarbon homologue (HD). There is a good agreement between the critical micelle concentrations of PFHD determined by the techniques used (pyrene fluorescence, vapor pressure osmometry, tensiometry, and Orange OT solubilization). Micellization of PFHD occurs at a much lower concentration than in the case of HD and also with a lower aggregation number, which has been evaluated by application of the phase-separation and mass-action models to the vapor pressure osmometry measurements. The higher polarity sensed by pyrene fluorescence in the PFHD aggregates in comparison with the HD aggregates led us to suggest the water incorporation due to the stiffness of the perfluorinated chain. The PFHD compound exhibits a much higher surface activity than HD and shows solubilization toward Orange OT at lower surfactant concentrations than HD. Heats and entropies of micellization for HD and PFHD have been

Damas et al. determined from the data on the CMCs at various temperatures. For both compounds, the results indicate a high entropic contribution on the free energy for micellization. The more positive entropy of micellization for PFHD than for HD is attributed to the higher capacity of the fluorocarbon chain to promote the structure of water than that of the hydrocarbon one. References and Notes (1) Kissa, E. Fluorinated Surfactants. Synthesis. Properties. Applications; Surfactant Science Series; Marcel Dekker: New York, 1994; Vol. 50. (2) (a) Kunieda, H.; Shinoda, K. J. Phys. Chem. 1976, 80, 2468. (b) Turro, N. J.; Lee, P. C. C. J. Phys. Chem. 1982, 86, 3367. (3) Klevens, H. B.; Raison, M. J. Chim. Phys. 1954, 51, 1. (4) Hajji, S. M.; Errahmani, M. B.; Coudert, R.; Durand, R. R.; Cao, A.; Taillandier E.; J. Phys. Chem. 1989, 93, 4819. (5) Coudert, R.; Hajji, S. M.; Cao, A. J. Colloid Interface Sci. 1993, 155, 173. (6) Coudert, R.; Paris, J.; Cao, A. J. Colloid Interface Sci. 1994, 163, 94. (7) Lianos, P.; Zana, R. J. Colloid Interface Sci. 1981, 84, 100. (8) Hayter, J.; Zemb, T. Chem. Phys. Lett. 1982, 93, 91. (9) Ayari, A.; Szonyi, S.; Rouvier, E.; Cambon, A. J. Fluorine Chem. 1990, 50, 67. (10) Dupont, G.; Dulou, R.; Duplessis-Kergomand, A. Soc. Organ. BreVet. Fr. 1951, 979, 551. (11) Moroi, Y. Micelles: Theoretical and Applied Aspects; Plenum Press: New York, 1992; Chapter 12. (12) Binana-Limbele, W.; Zana, R. Macromolecules 1987, 20, 1331. (13) Reichardt, C. In SolVents and SolVents Effects in Organic Chemistry, 2nd ed.; Ebel, H. F., Ed.; VCH: Weinheim, 1988; pp 365-371. (14) (a) Damas, C.; Abidnejad, M.; Benjelloun, A.; Brembilla, A.; Carre´, M. C.; Viriot, M. L.; Lochon, P. Colloid Polym. Sci. 1997, 275, 364. (b) Dong, D. C.; Winnik, M. A. Can. J. Chem. 1984, 62, 2560. (15) Frindi, M.; Michels, B.; Zana, R. J. Phys. Chem. 1991, 95, 4832. (16) Kalyanasundaram, K.; Thomas, J. K. J. Am. Chem. Soc. 1977, 99, 2039. (17) Anthony, O.; Zana, R. Macromolecules 1994, 27, 3885 and references therein. (18) Coudert, R.; Durand, R. J. Chim. Phys. 1987, 84, 1021. (19) Shinoda, K. Bull. Chem. Soc. Jpn. 1953, 26, 101. (20) Asakawa, T.; Mouri, M.; Miyagishi, S.; Nishida, M. Langmuir 1989, 5, 343. (21) McGinnis, T. P.; Wooley, E. M. J. Chem. Thermodyn. 1997, 29, 401. (22) Rusanov, A. I. MendeleeV Comm. 1996, 217, 6. (23) Desnoyers, J. E.; Caron, G.; DeLisi, R.; Roberts, D.; Roux, A.; Perron, G. J. Phys. Chem. 1983, 87, 1397. (24) Shinoda, K.; Hutchinson, E. J. Phys. Chem. 1962, 66, 577. (25) Rosenholm, J. B.; Burchfield, T. E.; Hepler, L. G. J. Colloid Interface Sci. 1980, 78, 1981. (26) Mukerjee, P. J. Phys. Chem. 1972, 76, 565. (27) Meguro, K.; Ueno, M.; Esumi, K. In Nonionic Surfactants. Physical Chemistry; Surfactant Science Series; Schick, M. J., Ed.; Marcel Dekker Inc.: New York and Basel, 1987; Vol. 23, Chapter 3 and references therein. (28) Selve, C.; Ravey, J. C.; Stebe, M. J.; Moudjahid, C. El.; Moumni, E. M.; Delpuech, J. J. Tetrahedron 1991, 47, 411. (29) Take’uchi, M.; Moroi, Y. J. Colloid Interface Sci. 1998, 197, 230. (30) Schott, H. J. Phys. Chem. 1967, 71, 3611. (31) Teddy, G. J. T.; Wheeler, B. A. J. Colloid Interface Sci. 1974, 47, 59. (32) Gerry, H. E.; Jacobs, P. T.; Anaker, E. W. J. Colloid Interface Sci. 1977, 62, 556. (33) Tadros, Th. J. Colloid Interface Sci. 1980, 74, 196.