Thermodynamics of self-assembly of sodium octanoate: comparison with a fully fluorinated counterpart

Thermodynamics of self-assembly of sodium octanoate:

comparison with a fully fluorinated counterpart

ALFREDO GONZALEZ-PEREZ, GERARDO PRIETO, JUAN M. RUSOand FELIX SARMIENTO*

Group of Biophysics and Interfaces, Department of Applied Physics, Faculty of Physics,University of Santiago de Compostela, E-15782 Santiago de Compostela, Spain

(Received 2 June 2003; revised version accepted 3 September 2003)

The isotherms of conductivity of sodium octanoate were measured and the critical micelleconcentration (cmc) and degree of ionization of the micelles, �, determined in a range oftemperatures (273–343 K) above the Krafft point. The thermodynamic parameters, Gibbs freeenergy �G0

m, enthalpy �H0m, and entropy �S0

m of micelle formation, were determined frompolynomial adjustments of the temperature dependence of cmc and from a proposedthermodynamic model based on the works of Muller [1993, Langmuir, 9, 96] and Rodrıguezet al. [2002, J. Colloid Interface Sci., 250, 438]. The increase in heat capacity uponmicellization, �C0

pm, was estimated from the parameters of the model and the enthalpy–

entropy compensation phenomena discussed. Finally, for information on their structuraldifferences, hence to understand their different behaviours, thermodynamic parameters arediscussed, comparing the corresponding fluorocarbon compound. A remarkable shift inminimum temperature in the U-shaped curve of cmc versus temperature was found whenhydrogen was substituted by fluorine in the hydrophobic chain of the surfactant. Thisbehaviour is a consequence of the special characteristics of the fluorine substituent in thehydrophobic tail and was reflected in the thermodynamic parameters and in the enthalpy–entropy compensation parameters, presenting different intercepts at the same compensationtemperature.

1. Introduction

The amphiphilic character of ionic salts, called sur-factants, in aqueous solutions is well known. Thesemolecules have received much attention because they canform spontaneous self-organized structures, which areof interest from a theoretical and practical point of view.The simplest of these structures, and hence the bestknown, is called ‘‘micelle’’ and the process of formation isnamely micellization. In general, micellization isa consequence of the dual character of these moleculesin solution, consisting of an ionic head group and a non-polar tail. When the salt is dissolved in water and theconcentration increases, the behaviour of this simple saltin solution is transformed to that of a self-assemblystructure. The concentration at which this process occursis called the critical micelle concentration (cmc) and thisconcentration depends on numerous parameters such aschemical structure of the surfactant, characteristics ofthe solvent, temperature and added compounds, e.g.salts, alcohols and oils, etc.

The self-assembly process of sodium octanoate inaqueous solutions has received much attention in thelast 40 years because this is a limiting case of micelle

formation, given that it has a large critical micelleconcentration and a very low aggregation number. Themicellization of sodium octanoate has been studiedextensively by Ekwall et al. [1–7] who applied differentphysical techniques such as density, viscosity and vapourpressure. Solubilization of alcohols in sodium octanoatemicelles [8, 9], as well as the formation of mesophases[10–14], were studied by the same authors.Sodium octanoate was frequently used for molecular

dynamics simulation studies due to the small aggrega-tion number of micelles that form in solution. The earlierstudies come from Watanabe et al. [15], Laaksonen andRosenholm [16] and Shelley et al. [17]. More recently, theworks of Khun and co-workers [18–20] have been addedto this information. Their studies include modelling of asolubilized alcohol in a sodium octanoate micelle.Relevant experimental data, such as aggregation num-ber, can be used to check the appropriateness ofthe molecular dynamics simulations. However, moredata are needed in the literature to better understandself-assembly processes.In the studies developed by Zemb et al. [21] aggrega-

tion numbers of sodium octanoate were determined bylight scattering at surfactant concentrations� 1.8M. Thenumbers go from 11 to 26, depending on the surfactant*Author for correspondence. e-mail: [email protected]

MOLECULAR PHYSICS, 10 NOVEMBER 2003, VOL. 101, NO. 21, 3185–3195

Molecular Physics ISSN 0026–8976 print/ISSN 1362–3028 online # 2003 Taylor & Francis Ltdhttp://www.tandf.co.uk/journals

DOI: 10.1080/00268970310001624515

concentration. The results are in good agreement withthose reported from neutron scattering, giving similaraggregation numbers. Other relevant studies includework by Lindman et al. [22] on tracer self-diffusion ofmicelle formation of sodium octanoate and the workof Persson et al. [23] on NMR studies of micellarsolutions. The solubilization of alcohols in sodiumoctanoate micelles has been studied by SANS [24] andRaman spectroscopy [25]. Some preliminary studies onapparent molar volumes and heat capacities of sodiumoctanoate at 25�C have been determined by Leduc andDesnoyers [26].

A model based on previous works of Ruckenstein andNagarajan [27], Israelachvili et al. [28] and Tanford [29]using three parameters: interfacial tension betweenwater and a liquid hydrocarbon, the standard free energyof transfer of hydrocarbon from water to liquidhydrocarbon and the net charge of the polar groups,has been applied by Eriksson et al. [30] to determinethe thermodynamics of micelle formation of sodiumoctanoate in aqueous solutions. Their results are in goodagreement with the experimental data. More recently,the aqueous solution of sodium octanoate has beenstudied by D’Angelo et al. [31] using a sound velocitytechnique to determine the temperature dependence ofthe cmc. These data were applied to estimate thethermodynamic parameters of micellization using thepseudo-phase separation model and the small contribu-tions to the standard Gibbs free energy of micellizationhave been discussed as a consequence of a compensatingeffect between enthalpy and entropy dependences ontemperature.

The aim of this work is to study experimentally theself-assembly process of sodium octanoate in aqueoussolution and determine the thermodynamics of micelli-zation using a proposed model based on previous worksof Muller [32] and Rodrıguez et al. [33] that take intoaccount the contribution of the degree of ionization ofmicelles to thermodynamic parameters. To carry outthis aim the electrical conductivity technique was usedover a wide range of temperatures and the results arediscussed for comparative purposes with the estimatedparameters obtained from polynomial determinations.Finally, for a better understanding of fluorocarbon/hydrocarbon compounds, the system was compared withthe fluorinated counterpart.

2. Experimental

2.1. MaterialsSodium octanoate was obtained from Lancaster

Synthesis Ltd (No. 10241, 97%) and was used asreceived. All measurements were performed usingdistilled water with conductivity below 3�S cm�1 at

298.15K. Sodium perfluorooctanoate from LancasterSynthesis Ltd (No. 16988, 97%) was utilized for thecomparative study.

2.2. InstrumentationConductivities were measured using a Kyoto

Electronics conductometer model CM-117 with aK-121 cell type. The cell constant was determined usingKCl solutions following the procedure suggested byMonk [34]. All measurements were performed in aPolyScience Model PS9105 thermostatted waterbath,at a constant temperature within �0.05K. The deter-mination of the isotherms of conductivity was carriedout by continuous dilution of a concentrated sampleprepared by weight.

3. Results and discussion

It is well known that ionic surfactants have a criticaltemperature at which the solid compound can bedissolved in aqueous solution and this temperaturedepends on the nature of the surfactants and solvent.This temperature is known as the Krafft point and isan important parameter used to study the appro-priate temperature range of micellization. The termKrafft point was introduced by Lawrence in 1935 [35]and interpreted as a phase transition. Later, Alexanderand Johnson [36] interpreted the Krafft point phenom-enon as the point at which the transfer of soap moleculesis effective from the solid phase to the micelle, since theconcentration of single molecules only increases quiteslowly once micelles are present. More recently, theconcept has been reviewed and discussed by Moroi [37]and finally it was concluded that the Krafft point is thetemperature at which the solubility of surfactants asmonomers becomes high enough for the monomers tocommence aggregation or micellization.To find the appropriate temperature range for

research the conductivity technique was applied todetermine the Krafft point. From a concentration of0.8mol kg�1, which is approximately twice that abovethe cmc, the temperature was increased and conduc-tivity measured simultaneously for each point untilequilibrium was reached. Figure 1 shows the curve ofconductivity versus temperature, with the Krafft pointfor the sodium octanoate close to 273.15K.With this information the temperature range chosen to

work with was 298 to 343K in which the system showsclear solubility. Figure 2 represents the isotherms ofmolality dependence of conductivity at 298, 313 and333K. Similar plots were found at other temperatures,not shown for clarity. The breaks found in everyisotherm are an experimental confirmation of the self-assembly process and the formation of a micellar phase.

3186 A. Gonzalez-Perez et al.

To determine the cmc a least-squares analysis wascarried out for the linear fragments of the pre-micellarand post-micellar regions and the point of intersection ofthe two extrapolated lines was taken as the cmc. As seenin figure 2, breaks occur in a large range of molalities.This phenomenon is characteristic of surfactants withshorter chains, i.e. the breaks are more abrupt when the

hydrophobicity of the surfactant increases. This presentsan added difficulty in the determination of cmc becausethe value depends on the number of points used in thelinear fits. As criteria, only values far from the curvaturezone and appropriate linear fits having a correlationat least of 0.9998 were chosen. Figure 3 shows thetemperature dependence of the slopes above and below

270 280 290 300 310 320 330 340

10

20

30

40

50

60

70

80

κ (m

S c

m−1)

T (K)

Figure 1. Specific conductivity, �, as a function of temperature for aqueous solutions of sodium octanoate at a concentration twicethe critical micelle concentration.

0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.95

10

15

20

25

30

35

40

κ (m

S c

m−1

)

m (mol kg−1)

Figure 2. Specific conductivity, �, as a function of molality, m, for aqueous solutions of sodium octanoate at differenttemperatures: (h) 298K and (4) 313K, (�) 333K.

Thermodynamics of self-assembly of sodium octanoate 3187

cmc and indicates an increase in the slope below cmc at313K. To understand this phenomenon, the tempera-ture dependence of conductivity was measured for afew molalities of sodium octanoate below cmc and forcorresponding NaCl solutions. The main contribution ofthe change in conductivity with temperature was due tothe viscosity of the solvent. To eliminate the viscosity

dependence, specific conductivity was multiplied by theviscosity of water. This result is known as Walden’s rule(�� ¼ constant) and it is well understood for electrolytesolutions.In figure 4 the results of the temperature dependence

of �� for sodium octanoate and NaCl are shown. Theresults indicate an increase in the values to 313K that

295 300 305 310 315 320 325 330 335

30

35

40

45

50

55

60

dκ/d

m (

mS

cm

−1 m

ol−1 K

)

T (K)

Figure 3. Slopes of specific conductivity versus molality plots, d�=dm, for aqueous solutions of sodium octanoate, as a function oftemperature: (�) above and (�) below the critical micelle concentration.

290 300 310 320 330 340 3500.50

1.00

1.50

2.00

2.50

290 295 300 305 310 315 320 325 3301.86

1.88

1.90

1.92

1.94

T (K)

κ η

(mS

cm

−1 P

a s)

Figure 4. Specific conductivity multiplied by water viscosity, ��, as a function of temperature for: (�) aqueous solution of sodiumoctanoate and (�) sodium chloride.


can be attributed to loss of water molecules surroundingthe Naþ cation which is more hydrated at lowertemperatures.

The ratio of the slopes calculated for the two linearfragments in each isotherm gives an estimation of thedegree of ionization of the micelles [38]. To eliminate thecontribution of the water surrounding the Naþ cationto the general value of conductivity below cmc, theslopes from the conductivity data above 313K wererecalculated. The results indicate that negligible tem-perature dependence was found for the ionization degreeof the micelles in the temperature range studied. Theresults of these cmc values together with those pre-viously reported by D’Angelo et al. [31] are shown atvarious temperatures in figure 5.

To compare with the low values of cmc the isotherm at288 K was determined and the value included in figure 5.This value is comparable to that reported by D’Angelo etal. [31]. The cmc data are in good agreement with theprevious data reported by D’Angelo et al. [31] obtainedby measuring the sound velocity above and below cmc.

The curve shows a typical U-shape. This behaviour isinterpreted as a consequence of two opposing effects: (a)the hydration of the head group and (b) the structuredwater molecules surrounding the hydrophobic alkylchain. The increase in temperature produces a decreasein the hydration of the head group and an increase in thebreakdown of the structured water molecules surround-ing the hydrophobic alkyl chain. The first favours micelleformation while the second does not. The first effect isdominant at lower temperatures and the second athigher, so the minimum represents the compensation of

both effects. The second effect becomes greater whenthe hydrophobic chain length increases. The decrease inthe minimum when the hydrophobicity increases is aconfirmation of this effect, as reflected in the works ofZielinski et al. [39], Chen et al. [40] and Rodrıguez et al.[33]. The effect of the counterion on the temperatureðT *Þ minimum is also large and depends on differenthydration capabilities and probably must follow theHofmeister series [41, 42] as suggested by Chen et al. [42].To know more about the process of micellization

the thermodynamic parameters associated with theself-assembly process of sodium octanoate can bedetermined. Classical methods used to estimate thermo-dynamic parameters are based on both the chargedphase separation model [43] and the mass action model[44]. From both models the Gibbs free energy ofmicellization, �G0

m, can be obtained as follows

�G0m ¼ 2� �ð ÞRT ln �cmc, ð1Þ

where � is the degree of ionization of the micelle and�cmc is the mole fraction of the surfactant at the cmc.From the temperature dependence of the cmc, �H0

m and�S0

m the enthalpy and entropy of micellization, respec-tively, can be estimated and therefore the increase inheat capacity upon micellization, �C0

pm, can be easily

determined knowing the temperature dependence of�H0

m or �S0m. The relations are as follows

�H0m ¼ �T2 @ð�G0

m=TÞP=@T� �

, ð2Þ

�S0m ¼ ð�H0

m ��G0mÞ=T , ð3Þ

�C0pm

¼ ð@�H0m=@TÞP ¼ Tð@�S0

m=@TÞP: ð4Þ

285 300 315 330 345 360 375

0.34

0.35

0.36

0.37

0.38

0.39

0.40

0.41

cmc

(mol

kg−1

)

T(K)

Figure 5. Temperature dependence of the critical micelle concentration for aqueous solutions of sodium octanoate: (g) data ofpresent work; (�) data from [31].


The usual procedure starts with the tempera-ture dependence of ln �cmc that can be obtained bycalculating the polynomial

ln �cmc ¼ Aþ BT þ CT2 þDT3 þ � � � , ð5Þ

which best fits the data. This method allows thedetermination of all the quantities from the parametersof the polynomial fit. Depending on the number ofparameters, the values of �H0

m, �S0m and �C0

pm canchange arbitrarily. Parameters A, B, C, D, . . . have noobvious physical meaning and in some cases depend onthe number and accuracy of the experimental data usedfor their determination.

This procedure was applied by D’Angelo [31] withthe temperature dependence of the cmc values of sodiumoctanoate obtained from sound velocity data and usedwithout the parameter ð2� �Þ in equation (1). Thisestimation neglects the contribution of ionization degreeto the values of Gibbs free energy of micellization.Another procedure, used earlier by Eriksson et al. [30],considersotherparameters, as commentedon inSection1,giving good concordance with the experiments.

3.1. Thermodynamic modelIt is well known that the temperature dependence of

the cmc shows a U-shape with a minimum that dependson the chain length of the surfactant. This phenomenonhas been reviewed by La Mesa [45], Muller [32] andChen et al. [40]. To obtain an appropriate method todetermine the thermodynamic parameters associatedwith micellization, Muller [32] proposed a thermody-namic relation derived from the hypothesis that � and�C0

pmare constants. More recently, Rodrıguez et al. [33]

developed a similar relation to the hypothesis that � and�C0

pmare temperature dependent. In the case presented

here, it is asssumed that � is not temperature dependent,a finding that is consistent with the experimental data,and that the standard change in heat capacity, �C0

pm,

varies linearly with temperature [46–48].The starting point in the derivation of the equation

describing the U-shaped curves is the consideration thatthe change in heat capacity upon micellization, �C0

pm,

exhibits a linear dependence on temperature:

�C0pm

¼ �C0*pm

þ �ðT � T *Þ ð6Þ

where �C0*pm

is the value of �C0pm

at T ¼ T* and � is theslope. Since

�C0pm

¼ ð@�H0m=@TÞP ¼ Tð@�S0

m=@TÞP, ð7Þ

it is straightforward to obtain

�H0m ¼�H0*

m þ�C0*pmðT �T*Þ þ ð�=2ÞðT �T *Þ

2, ð8Þ

�S0m ¼�S0*

m þ�C0*pm

ln ðT=T *Þ

þ � T �T * �T * ln ðT=T *Þ� �

, ð9Þ

where �H0*m and �S0*

m are, respectively, the enthalpyand the entropy of micellization at the minimum, that is,for T ¼ T *. Taking into account that

�G0m ¼ �H0

m � T�S0m ¼ ð2� �ÞRT ln �cmc, ð10Þ

an explicit equation for ln �cmc can be obtained

R ln�cmc

�*cmc

¼�C0*

pm

2� �1�

T*

T� ln

T

T *

� �

þ�

2� �

T*2 � T2

2Tþ T* ln

T

T*

� �

þ�H0*m

1

ð2� �ÞT�

1

ð2� �ÞT *

� �

��S0*m

1

2� ��

1

2� �*

� �: ð11Þ

The above equation includes the value of �S0*m related

to R ln �*cmc through the expression:

�H0*m

ð2� �*ÞT*�

�S0*m

2� �*¼ R ln �*

cmc, ð12Þ

where �* is the corresponding value of � at T *. Since thederivative of the lengthy equation at T * equals zero, thefollowing relation holds:

�S0*m ¼ R ln �*

cmcð�� 2Þ, ð13Þ

which enables elimination of the value of�S0m at T * from

the previous equation. Since

�H0m ¼�T2 @ð�G0

m=TÞ

@T

� �P

¼ RT2 ln �cmcd�

dT� ð2� �ÞRT2 d ln �cmc

dTð14Þ

at the minimum,

�H0*m ¼ 0: ð15Þ


Substituting this expression for �H0m in the lengthy

equation, finally the following equation is obtained:

ln �cmc ¼ ln �*cmc þ

�C0*pm

2� �ð ÞR1�

T*

T� ln

T

T *

� �

þ�

2� �ð ÞR

T*2 �T2

2TþT * ln

T

T*

� �: ð16Þ

This equation is applied to fit the ln �cmc versus Tdata. This equation relates in a formal way the entropy,enthalpy and heat capacity of micellization with thecorresponding Gibbs free energy as a function oftemperature. The parameters in this equation are: T *,ln �cmc,�C0*

pm, � and �. Since � is a quantity which is not

inherent to thermodynamics, and also because of theexcessive number of parameters used, first � is extractedfrom the experimental data, thus reducing the number offitting parameters to four. By introducing a set of initialparameters and by applying a nonlinear regression, the

values of the parameters for the best fit can be estimated.The present work shows that in all cases, independentlyof the initial approximations, the program alwaysconverges to an identical set of final fitted parameters.Table 1 presents the values of T *, ln �cmc,�C0*

pm, � and

� obtained by fitting the experimentally determinedvalues of ln �cmc to equation (16).To compare the resulting temperature dependences of

the thermodynamic parameters of micellization, figure 6(a) shows the temperature dependence of �G0

m. Somesmall discrepancies between the two methods appear butonly at low temperatures. Figures 6 (b) and (c) show thetemperature dependence of �H0

mand �S0m, respectively.

It can be seen that at low temperature ranges thedifferences between the methods increase quickly. Thedata reported by D’Angelo et al. [31] show higher �G0

m

values in all ranges studied of �2 kJ mol�1 and thevalues of �H0

m here presented are higher by 14 kJ mol�1

in the case of the thermodynamic model and 4 kJ mol�1

in the case of the polynomial fit.

290 300 310 320 330 340 350

60

80

100

120

290 300 310 320 330 340 350-21000

-20000

-19000

-18000

-17000

290 300 310 320 330 340 350-6000

0

6000

12000

18000

(c)

∆Sm0 (

J m

ol−1 K

−1)

T (K)

(a)

∆Gm0 (

J m

ol−1)

(b)

∆Hm0 (

J m

ol−1)

Figure 6. (a) Standard free energy, �G0m, (b) enthalpy, �H0

m, and (c) entropy, �S0m, of micelle formation of sodium octanoate as a

function of temperature. (- - -) Results obtained by polynomial method, equations (1), (2) and (3). (—) Results obtained byapplication of the model, equations (8), (9) and (10).

Table 1. Parameters from the fitting of ln �cmc versus temperature T values to equation (16).

System T*/K ln��cmc �C0

pm�/R/K�1 b

Sodium octanoate 333.4 �5.0965 �37.0127 1.05106 0.58672

Sodium perfluorooctanoate 316.2 �7.5479 �54.6651 0.16346 0.57622


The study of the enthalpy–entropy compensationphenomena can give a measure of the desolvation partof the micellization process through the temperatureof compensation, Tc. This parameter is a characteristicof solute–solute and solute–solvent interactions, as sug-gested by Chen et al. [40]. In general the compensationphenomenon between the enthalpy change and theentropy change appears in various processes [49, 50]and can be described as

�H0m ¼�H*

m þTc�S0m, ð17Þ

when �H*m is the intercept on the �H0

m axis whichcharacterizes the solute–solute interactions and can beconsidered as an index of the chemical part of the processof micellization. Figure 7 shows the plots of �H0

m versus�S0

m for the data estimated using both methods.Figure 7 indicates that essentially the two methods

give the same results for the enthalpy–entropy compen-sation process. The linear fit gives a Tc of 321� 2K and316� 2K for the classical polynomial determination andthe proposed model, respectively. The values are close tothose reported by Chen et al. [40] for sodium alkyl-sulphates (n- and 2-decyl, n- and 2-dodecyl, n- and2-tetradecyl), 304� 3K.

3.2. Comparative study of hydrogenatedand fluorinated compounds

Dramatic changes have been reported in hydrogenatedchains when the hydrogen was substituted by fluorine

[51]. These special properties of fluorinated chains comefrom the distinct properties of fluorine, i.e. higherbulkiness and lower polarizability than the hydrogenatom [52–54]. Fluorinated amphiphiles have a greatertendency to water self-assembly than their correspondinghydrogenated ones, given their contribution to themicellization of 1.6 times more per CF2 than per CH2

group [55]. This behaviour is reflected in the big decreasein the change in free energy of adsorption to transfera CF2 group from water to the air–water interface incomparison with that by the CH2 group. Micelles offluorinated amphiphiles tend to be rod-shaped,as predicted by geometric determinations confirmedexperimentally [56, 57]. Fluorinated surfactants are ofgreat interest not only from a theoretical point of viewbut also for their wide variety of applications [53].The comparative study of fluorinated and non-

fluorinated surfactant systems is interesting from atheoretical point of view to study the hydrophobic effect.Studies on fluorocarbon and hydrocarbon surfactantsas well as their comparative discussion have beenperformed by many authors such as De Lisi et al. [58],Ravey and Stebe [59] and Fukada et al. [60]. To studythe effect of substituting the hydrocarbon chain by afluorocarbon one, the same thermodynamic treatmentwas applied to both the present hydrocarbon compoundas well as to the corresponding fluorocarbon. Usingthe data of the cmc at different temperatures obtainedfrom conductivity, density and ultrasound velocitymeasurements for sodium perfluorooctanoate [61], theparameters obtained by equation (16) are listed in table 1.

50 60 70 80 90 100 110

-4000

0

4000

8000

12000

16000

∆Hm0 (

J m

ol−1)

∆Sm

0 (J mol−1 K−1)

Figure 7. Enthalpy–entropy compensation plot for aqueous solutions of sodium octanoate. (�) Values obtained by polynomialmethod, equations (2) and (3); (h) values obtained by application of the model, equations (8) and (9).


Figure 8 shows the ln �cmc= ln �*cmc as a function of

temperature for the hydrocarbon compound and thefluorocarbon. The U-shaped curve shifted to a lowertemperature in the case of the fluorocarbon compound.This behaviour is a consequence of the higher hydro-phobicity of the fluorocarbon tail. The minimum shifted

by 117.2 K. The hydrophobicity of the fluorocarbonsurfactant is comparable to that found in the corre-sponding hydrocarbon surfactant with a hydrophobicchain that is 1.5 times longer [62]. An analogous situationoccurs in the series of homologues with different alkylchain length when the minimum is shifted to lower

290 300 310 320 330 340 350

1.000

1.005

1.010

1.015

1.020

1.025

1.030

ln χ

cmc* / l

n χ cm

c

T (K)

Figure 8. Temperature dependence of the natural logarithm of the critical micelle concentration at the critical temperature/naturallogarithm of the critical micelle concentration (both in mole fractions) relationship for aqueous solutions of: (�) sodiumoctanoate and (h) sodium perfluorooctanoate.

290 295 300 305 310 315 320 325 330 335 340 345-30000

-25000

-20000

-15000

-10000

-5000

0

5000

10000

15000

20000

∆Gm 0 , ∆

Hm0 (

J m

ol−1

)

T (K)Figure 9. Temperature dependence of standard parameters of micelle formation in aqueous solution: (h) �G0

m of sodiumoctanoate, (g) �G0

m of sodium perfluorooctanoate, (�) �H0m of sodium octanoate and (�) �H0

m of sodium perfluorooctanoate.


temperatures when the chain length increases, asreported by Zielinski et al. [39] and Rodrıguez et al.[33]. In figure 9 the temperature dependence of �G0

m

and �H0m for the fluorocarbon and the correspond-

ing hydrocarbon compound is shown. �G0m at 298K

decreases by �9 kJ mol�1 for the fluorocarbon com-pound and �H0

m is higher by 7.8 kJ mol�1 for thehydrocarbon compound. The critical temperature with�H0

m ¼ 0 shifted 16.5 K for the sodium perfluoro-octanoate. These results indicate that the process of self-assembly was entropically driven and occurs at lowtemperature for the fluorocarbon salt as a consequenceof its higher hydrophobicity.

The �C0pm

associated with the process of micellizationfor both hydrocarbon and fluorocarbon compounds areshown in figure 10. The results indicate a higher slope forthe temperature dependence of �C0

pmfor the hydro-

carbon compound. This can be interpreted to mean thatthe substitution of hydrogen by fluorine in the carbonchain, which leads to lower energy, is needed to breakdown the water structure around a fluorocarbon chainrather than the hydrocarbon chain as seen in figure 9where the temperature dependence of the thermody-namic parameters of micellization of the two types ofsystems is shown.

Finally the enthalpy–entropy compensation processindicates that both compounds show a similar Tc andthe difference appears in the intercept of the linear

dependence with the �H0m axis. The value of sodium

perfluorooctanoate is lower, �8:8 kJ mol�1.

4. Conclusions

Sodium octanoate in aqueous solution was beenstudied by conductometry. Thermodynamic parameterswere determined by applying a proposed model based onthe linearity of�C0

pmand taking into account the value of

�. The model provides a procedure without arbitraryselection and the results are comparable with thosereported by the usual polynomial fit. The effect of sub-stituting the hydrocarbon tail by the fluorocarbon oneis discussed. The decrease in T* and ln �*

cmc was attrib-uted to the higher hydrophobicity of the fluorocarboncompound and the thermodynamic results given from themodel seem to confirm this finding. The enthalpy–entropy compensation phenomena give similar slopesyet different intercepts for fluorinated and hydrogena-ted compounds. This difference in intercepts again is aconsequence of their distinct hydrophobicity. The similarTc shows that only small differences can be attributed tothese compounds with respect to the desolvation partof the micellization process and indicates that at thislevel the same Tc is characteristic of the homologoushydrogenated and fluorinated compounds.This research was funded by the Spanish Ministry of

Science and Technology, through the Project MAT2002-00608 (European FEDER support included). Alfredo

290 295 300 305 310 315 320 325 330 335 340 345-80

-70

-60

-50

-40

-30

-20

∆Cp,

m

0 (

J m

ol−1 K

−1)

T (K)Figure 10. Temperature dependence of the increase of heat capacity upon micellization, �C0

pm, for aqueous solutions of: (�)

sodium octanoate and (h) sodium perfluorooctanoate.


Gonzalez-Perez is grateful to the University of Santiagode Compostela for his postdoctoral grant.

References[1] EKWALL, P., EIKREM, H., and MANDELL, L., 1963,

Acta Chem. Scand., 17, 111.[2] EKWALL, P., and HOLMBERG, P., 1965, Acta Chem.

Scand., 19, 455.[3] EKWALL, P., and HOLMBERG, P., 1965, Acta Chem.

Scand., 19, 573.[4] EKWALL, P., LEMSTROM, K. E., EIKREM, H., and

HOLMBERG, P., 1967, Acta Chem. Scand., 21, 1401.[5] EKWALL, P., EIKREM, H., and STENIUS, P., 1967,

Acta Chem. Scand., 21, 1639.[6] STENIUS, P., and EKWALL, P., 1967, Acta Chem. Scand.,

21, 1643.[7] EKWALL, P., and STENIUS, P., 1967, Acta Chem. Scand.,

21, 1767.[8] EKWALL, P., MANDELL, L., and FONTELL, K., 1977, J.

colloid interface Sci., 61, 519.[9] EKWALL, P., and MANDELL, L., 1979, J. colloid interface

Sci., 69, 384.[10] EKWALL, P., MANDELL, L., and FONTELL, K., 1968,

Acta Chem. Scand., 22, 365.[11] EKWALL, P., MANDELL, L., and FONTELL, K., 1968, Acta

Chem. Scand., 22, 697.[12] FONTELL, K., MANDELL, L., and EKWALL, P., 1968,

Acta Chem. Scand., 22, 3209.[13] EKWALL, P., 1988, Colloid polym. Sci., 266, 1150.[14] EKWALL, P., 1989, Colloid polym. Sci., 267, 607.[15] WATANABE, K., FERRARIO, M., and KLEIN, M. L., 1988,

J. phys. Chem., 92, 819.[16] LAAKSONEN, L., and ROSENHOLM, J. B., 1993, Chem.

Phys. Lett., 216, 429.[17] SHELLEY, J. C., SPRIK, M., and KLEIN, M., 1993,

Langmuir, 9, 916.[18] KUHN, H., and REHAGE, H., 1997, Ber. Bunsenges. phys.

Chem., 101, 1485.[19] KUHN, H., and REHAGE, H., 1997, Ber. Bunsenges. phys.

Chem., 101, 1493.[20] KUHN, H., BREITZKE, B., and REHAGE, H., 2002, J.

colloid interface Sci., 249, 152.[21] ZEMB, T., DRIFFORD, M., HAYOUN, M., and JEHANNO, A.,

1983, J. phys. Chem., 87, 4524.[22] LINDMAN, B., KAMENKA, N., PUYAL, M.-C., BRUN, B.,

and JONSSON, B., 1984, J. phys. Chem., 88, 53.[23] PERSSON, B.-O., DRAKENBERG, T., and LINDMAN, B.,

1979, J. phys. Chem., 83, 3011.[24] ROSENHOLM, J. B., LARSSON, K., and DINH-NGUYEN, N.,

1977, Colloid polym. Sci., 255, 1098.[25] HAYTER, B., HAYOUN, M., and ZEMB, T., 1984, Colloid

polym. Sci., 262, 798.[26] LEDUC, P.-A., and DESNOYERS, J. E., 1973, Can. J.

Chem., 51, 2993.[27] RUCKENSTEIN, E., and NAGARAJAN, R., 1975, J. phys.

Chem., 79, 2622.[28] ISRAELACHVILI, J., MITCHELL, D., and NINHAM, B., 1976,

Trans. Faraday Soc., 72, 1525.[29] TANFORD, C., 1980, The Hydrophobic Effect: Formation

of Micelles and Biological Membranes, 2nd Edn(New York: Wiley).

[30] ERIKSSON, F., ERIKSSON, J. C., and STENIUS, P., 1979,Solution Chemistry of Surfactants (New York: PlenumPress).

[31] D’ANGELO, M., ONORI, G., and SANTUCCI, A., 1994,Colloid polym. Sci., 97, 154.

[32] MULLER, N., 1993, Langmuir, 9, 96.[33] RODRIGUEZ, J. R., GONZALEZ-PEREZ, A., DEL CASTILLO,

J. L., and CZAPKIEWICZ, J., 2002, J. colloid interface Sci.,250, 438.

[34] MONK, C. B., 1961, Electrolytic Dissociation (London:Academic Press).

[35] LAWRENCE, A. S. C., 1935, Trans. Faraday Soc., 31, 206.[36] ALEXANDER, A. E., and JOHNSON, P., 1949, Colloid

Science (London: Oxford University Press).[37] MOROI, Y., 1992, Micelles Theoretical and Applied

Aspects (New York: Plenum Press).[38] HOFFMANN, H., and ULBRIGHT, W., 1977, Z. phys.

Chem. NF, 106, 1359.[39] ZIELINSKI, R., IKEDA, S., NOMURA, H., and KATO, S.,

1989, J. colloid interface Sci., 129, 175.[40] CHEN, L.-J., LIN, S.-Y., and HUANG, C.-C., 1998, J. phys.

Chem. B, 102, 4350.[41] KHALWEIT, M., LESSNER, E., and STREY, R., 1984, J.

phys. Chem., 88, 1937.[42] CHEN, L.-J., HSU, M.-C., LIN, S.-T., and YANG, S.-Y.,

1995, J. phys. Chem., 99, 4687.[43] SHINODA, K., and HUTCHINSON, E., 1962, J. phys. Chem.,

66, 577.[44] PHILIPS, J. N., 1955, Trans. Faraday Soc., 51, 561.[45] LA MESA, C., 1990, J. phys. Chem., 94, 323.[46] DESNOYERS, J. E., and PERRON, G., 1996, Langmuir, 12,

4044.[47] MCGINNIS, T. P., and WOOLLEY, E. M., 1997, J. Chem.

Therm., 29, 401.[48] KRESHECK, G. C., 2001, J. phys. Chem. B, 105, 4380.[49] SUGIHARA, G., and HISATOMI, M., 1999, J. colloid

interface Sci., 219, 31.[50] LUMBRY, R., and RAJENDER, S., 1974, Biopolymers, 9,

1125.[51] KRAFFT, M. P., 2001, Adv. Drug Deliv. Rew., 47, 209.[52] TIDDY, G. J. T., 1985, Modern Trends of Colloid Science

in Chemistry and Biology, edited by H. F. Eicke (Basel:Birkhauser Verlag).

[53] KISSA, E., 1994, Fluorinated Surfactants, Synthesis,Properties, Applications, Surfactant Science Series,Vol. 50 (Marcel Dekker: New York).

[54] BONDI, A., 1964, J. phys. Chem., 68, 441.[55] SCHNEIDER, J., ERDELEN, C., RINGSDORF, H., and

RABOLT, J. F., 1989, Macromolecules, 22, 3475.[56] FUNG, B. M., MAMROSH, D. L., O’REAR, E. A., FRESH,

C. B., and AFZAL, J., 1988, J. phys. Chem., 92, 4405.[57] GUO, W., BROWN, T. A., and FUNG, B. M., 1991, J. phys.

Chem., 95, 1829.[58] De LISI, R., INGLESE, A., MILIOTO, S., and PELLERITO, A.,

1997, Langmuir, 13, 192.[59] RAVEY, J. C., and STEBE, M. J., 1994, Colloids Surf. A:

Physicochem. Eng. Aspects, 84, 11.[60] FUKADA,K.,KOBAYASHI,Y.,OTA,Y., FUJII,M.,KATO,T.,

and SEIMIYA, T., 2000, Thermochim. Acta, 352, 189.[61] GONZALEZ-PEREZ, A., RUSO, J. M., PRIETO, G., and

SARMIENTO, F., 2003, J. colloid interface Sci., submitted.[62] SHINODA, K., HATO, M., and HAYASHI, T., 1972, J.

phys. Chem., 76, 909.


Thermodynamics of self-assembly of sodium octanoate: comparison with a fully fluorinated counterpart

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