Journal of Colloid and Interface Science 305 (2007) 301–307www.elsevier.com/locate/jcis

Thermodynamics of aqueous solutionsof dodecyldimethylethylammonium bromide

Emilia Fisicaro ∗, Mariano Biemmi, Carlotta Compari, Elenia Duce, Monica Peroni

Dipartimento di Scienze Farmacologiche, Biologiche e Chimiche Applicate, Università di Parma, Viale G.P. Usberti 27A, 43100 Parma, Italy

Received 18 July 2006; accepted 24 September 2006

Available online 28 September 2006

Abstract

The thermodynamic properties of the aqueous solutions of dodecyldimethylethylammonium bromide (DEDAB) were determined as a functionof concentration by means of direct methods. Dilution enthalpies at 298 and 313 K, densities and sound velocities at 298 K were measured,allowing the determination of apparent and partial molar enthalpies, volumes, heat capacities and compressibilities. Changes in thermodynamicquantities upon micellization were derived using a pseudo-phase transition approach. These data allow for the determination of the effect ofthe –CH2– group, when added to the polar head of alkyltrimethylammonium bromides. The properties mainly affected by this addition are theenthalpies and, as a consequence, the entropies. The lowering of the charge density on the quaternary nitrogen due to the inductive effect of theethyl group, greater than that of the methyl one, raises the plateau value of apparent and molar enthalpy by a quantity similar to that due to theremoving of a methylene group from the hydrophobic chain. This effect does not play a great role in the value of the cmc (i.e. on the free energyof micelle formation), since the small decrease in cmc of DEDAB compared to DTAB reflects the increase in the overall hydrophobicity of themolecule. Volumes of DEDAB are greater than those of DTAB by about 15 cm3 mol−1, both at infinite dilution and at micellar phase, a value inagreement with that generally accepted for a methylene group. The trends of apparent molar heat capacities and compressibilities vs m are thesame as for DTAB: in fact, these quantities are related to the number of water molecules involved in the hydrophobic processes in solution, notvery greatly affected by the substitution of a methyl group by an ethyl one on the polar head. In summary, this substitution affects to a significantextent the first derivatives of the free energy, but does not affect the second derivatives.© 2006 Elsevier Inc. All rights reserved.

Keywords: Dodecyldimethylethylammonium bromide; Apparent and partial molar enthalpies of; Apparent molar heat capacities of; Apparent molar volumes of;Apparent molar compressibilities of; Micellization thermodynamic parameters of; Dilution heats of

1. Introduction

Quaternary ammonium surfactants are widely applied inpractical fields, as antimicrobial and antiseptic agents, as fab-ric softeners, as corrosion inhibitors, in mineral flotation, inroad paving, etc. [1–3]. The scientific and applicative interest incationic surfactants has greatly increased, partly because of theurge for cationic amphiphiles to be used as vectors in gene de-livery [4–8]. Quite recently, the use of gemini surfactants—i.e.surfactants in which at least two identical moieties are boundtogether by a spacer, generally formed by an alkyl chain, atthe polar head level—as nonviral vectors in gene therapy has

* Corresponding author.E-mail address: emilia.fisicaro@unipr.it (E. Fisicaro).

0021-9797/$ – see front matter © 2006 Elsevier Inc. All rights reserved.doi:10.1016/j.jcis.2006.09.063

been proposed [4–7], on account of the possibility of taking ad-vantage of their cationic character necessary for binding andcompacting DNA and of their superior surface activity. Fol-lowing our research studies in this direction, in order to col-lect thermodynamic information for defining structure activityrelationships and to predict thermodynamic behaviour of tailor-made cationic surfactants, we often need to know the effect ofadding a –CH2– group on the polar head of a quaternary am-monium surfactants. The substitution of a methyl group, boundto the positive nitrogen by an ethyl group in DTAB gives riseto dodecyldimethylethylammonium bromide (DEDAB), com-mercially available. Although the micellar and thermodynamicproperties of homologous series of alkyltrimethylammoniumhalides have been studied extensively, little attention has beendevoted to the effect on the solution thermodynamic properties

302 E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307

of the substitution of a methyl group by a longer alkyl chain onthe polar head, particularly by using direct measurements. Theeffect of the introduction of methyl groups on the hydrophilicmoiety of dodecylamine hydrochloride has been carefully eval-uated by direct methods some time ago [9]. More recently, theeffect on thermodynamic properties of micellization of the sub-stitution of the ethanolyl group in place of the methyl group onthe head of CTAB surfactant [10] or in place of the hydrogenon the head of cetylammonium bromide [11] has been inves-tigated by microcalorimetric, conductometric, and fluorimetrictechniques and from the dependence of the cmc on temperature,respectively. Unfortunately, the presence of the –OH, affectingthe hydrophilic/lipophilic balance of the surfactants, does notallow to separate the contribution of the methylene group alone.

Despite our efforts to find out information in the literatureabout DEDAB solution thermodynamics, we were not able tofind a complete and reliable enough description of its behav-iour in solution, the few data available being of a conflictingnature. In fact, thermodynamic parameters have been obtainedfrom the dependence on temperature of the cmc, obtained byconductivity measurements [12]. In this way also very small er-rors in the experimental data or the often present uncertainty inthe evaluation of cmc have a great effect on the value of thermo-dynamic quantities upon micellization. In this paper our resultson the solution thermodynamics of DEDAB, obtained by meansof direct methods, were collected and compared with those ofDTAB, with the goal of enriching the fundamental understand-ing on self-aggregation thermodynamics of surfactants and ofproviding a precious information for the theoretical evaluationof their properties.

2. Experimental

The surfactant dodecyldimethylethylammonium bromide(DEDAB) having purity �99% was purchased from AldrichChemical Company and used without further purification.

2.1. Apparatus

The enthalpies of dilution were measured by means ofthe Thermometric TAM (flow mixing cell) microcalorimeter,equipped with 221 Nano Amplifier, at 298 and 313 K. Thefreshly prepared surfactant solutions, kept before injection atthe experimental temperature by means of a Heto cryothermo-static bath, were diluted into the “mixing” measuring cell ofthe microcalorimeter in a ratio 1:1 by using CO2-free water.The solutions and the water were injected by means of a Gilsonperistaltic pump, Minipuls 2, and their flows were determinedby weight.

Density and sound velocity of the solutions were measuredby Paar DSA 5000, oscillating U-tube density (±0.000001 gcm−3) and sound velocity (±0.1 m s−1) meter which mea-sures to the highest accuracy in wide viscosity and temperatureranges. Based on an additional measuring cell made of stainlesssteel and high resolution electronics, the sound velocity of thefilled in sample can be determined accurately. Both measuring

cells are temperature controlled using a built-in solid state ther-mostat and two integrated Pt 100 platinum thermometers. Theinstrument was calibrated before each series of measurementsby degassed bidistilled water and dry air.

3. Results

The experimental data were expressed in terms of apparentand partial molar quantities of the solute, as usual in solutionthermodynamics, assuming infinite dilution as the referencestate. Apparent and partial molar quantities were obtained fromthe experimental data using methods stated in detail elsewhere[13–17].

For the sake of clarity, we recall that, with reference to thestate of infinite dilution, the molar enthalpy of dilution, �Hd,is given by:

(1)�Hd = LΦ,f − LΦ,i,

where LΦ is the apparent molar relative enthalpy and the in-dexes f and i stand for the final (after dilution) and initial (beforedilution) concentrations, respectively.

For ionic surfactant in the premicellar region, the relativeapparent molar enthalpy can be expressed by means of a poly-nomial of m1/2. Stopping the serial expansion at the third termwe obtain:

(2)LΦ = ALm1/2 + BLm + CLm3/2,

where AL is the limiting Debye–Huckel slope for relative en-thalpies accounting for the long range electrostatic solute–solute interactions. Parameters BL and CL are obtained from theexperimental points in the premicellar region by a least squarescurve fitting.

In the micellar region, the apparent molar enthalpies areevaluated by means of Eq. (1) and, when a value of LΦ vs m notexperimentally measured is needed, by graphical interpolation.

The partial molar enthalpies L2 are determined by drawingthe best curve for the apparent molar enthalpies vs m and thenby calculating the partial molar quantities as �(mLΦ)/�m

from points interpolated at regular intervals.Apparent and partial molar isobaric heat capacities vs m

were obtained from

(3)CΦ = [LΦ(313 K) − LΦ(298 K)

]/15

under the assumption that the heat capacities are constant in theexperimental temperature range.

In Fig. 1, the plots of the apparent and partial molar en-thalpies as a function of molality at 298 and 313 K are shown.

The apparent molar volumes, VΦ and adiabatic compress-ibility, Ks,Φ were calculated by means of the following equa-tions [18–21]:

(4)VΦ = M

d− 103(d − d0)

mdd0,

(5)Ks,Φ = −Mβs

d− 103(βs,0d − βsd0)

mdd0,

E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307 303

Fig. 1. Apparent molar relative enthalpies (", 2) and partial molar relativeenthalpies (!, 1) of DEDAB at 298 K (!, ") and 313 K (1, 2) as a functionof surfactant concentration at 298 K.

where d is the density of the solution of molality m, M is themolecular weight of the surfactants, d0 is the density of the sol-vent. βs,0 and βs are the coefficient of adiabatic compressibilityof the solvent and of the solute, respectively. The latter is cal-culated from sound velocity, u, and density data as

(6)βs = 100/(u2d

).

Tables with all the experimental data are available as supportinginformation.

4. Discussion

Only a few data have been published up to now as tomicelle formation thermodynamics of DEDAB [12,22–28],mainly derived from specific conductivity measurements. Asfar as we know, directly measured thermodynamic quantitiesfor the DEDAB–water system are not available, whereas spe-cific conductivity of DEDAB has been measured as a functionof temperature in the range 15–40 ◦C [12,22–25]. Cmc and ion-isation degree of the micelles have been obtained, allowing theevaluation of thermodynamic functions of micellization by ap-plying a mass action model.

The trend of cmc vs temperature shows the U-shaped de-pendence, typical of hydrophobic processes, with a minimumaround room temperature [22]. At 298 K the cmc resultsas being 0.0140(4) mol kg−1 and β , the degree of counte-rion binding, 0.72(2) [22,25,28] and at 313 K the values are0.0151 mol kg−1 for the cmc and 0.68 for β , very close tothose accepted for DTAB [29]. The small decrease in cmc ofDEDAB compared to DTAB reflects the increase in the over-all hydrophobicity of the molecule. DEDAB is considered inRef. [23] as an asymmetric double chain surfactant with a veryshort second alkyl chain, showing that in this case the free en-ergy to transfer from water to the micelle per methylene unitis significantly small for asymmetric double-chain surfactantswith a shorter alkyl chain.

Fig. 2. Apparent molar relative enthalpies (", Q) and partial molar relativeenthalpies (!, P) of DEDAB (!, ") and of DTAB (P, Q) from Ref. [32] as afunction of surfactant concentration at 298 K.

4.1. Apparent and partial molar enthalpies and heatcapacities

In general for ionic surfactants the curves of the apparentand partial molar enthalpies as a function of concentration, afterincreasing in the premicellar region, tend to level off at con-centrations above the cmc, where they are almost parallel. Thelowering of the curves in the micellar region, proportional tothe number of carbon atoms in the alkyl chain, is attributed tothe electrostatic interactions in micellar solutions. The compar-ison of the enthalpic trends of DEDAB and DTAB at 298 K(Fig. 2) show that, notwithstanding the greater total number ofcarbon atoms, the curves of DEDAB lie above those of DTAB,allowing us to evaluate the group contribution to the plateauvalue of a methylene when added to the polar head. It re-sults as being about +1.6 kJ mol−1 group−1 for LΦ and about+1.75 kJ mol−1 group−1 for L2. The effect is very close tothat obtained by shortening the alkylic chain by one methylenegroup [9,14]. The very short second chain present on the nitro-gen bearing the positive charge allows for both a better chargedelocalization and an increase in the size of the polar head,so that the charges are farther apart on the micelle surface ofDEDAB than on that of DTAB. This is confirmed by the valueof the degree of counterion binding, lower for DEDAB than forDTAB. On the other hand, we have compared the behaviour ofdifferent dodecylic surfactants having the same counterion, inorder to rationalize the effect of the polar head on the enthalpicproperties of their solutions [30]. We have thus established asort of thermodynamic “charge localization scale” for the po-lar head. The more delocalized the charge, the more similarthe behaviour to that of a nonionic surfactant. The lowering ofthe plateau value of the apparent and partial molar enthalpiescurves vs m in a homologous series of surfactant is related tothe increase in hydrophobicity of the alkyl chain and to the sizeof micelles; in contrast, for the same alkyl chain and counte-rion the plateau value is strictly related to the modulation ofthe charge density due to inductive or resonance effects. Theseeffects do not play a great role in the value of the cmc (i.e.

304 E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307

Fig. 3. Apparent molar relative heat capacities of DEDAB (") and of DTAB(P) from Ref. [32] as a function of surfactant concentration at 298 K.

in the free energy of micelle formation), the cmc of DEDABbeing only slightly smaller than that of DTAB, although theystrongly affect the solution enthalpies and, as a consequence,the entropies.

The trends of apparent and partial molar enthalpies vs m

(Fig. 1) are also very sensitive to the temperature, as generallyfound for processes involving the hydrophobic effect: they de-crease in micellar region with the increase in temperature [31].

Because, in general, the lowering of enthalpic curves belowcmc is nearly proportional to the temperature, it can be as-sumed that the heat capacities are constant in the temperaturerange examined. The trends of apparent molar heat capacitieswere obtained by averaging the ratios �LΦ/�T at each con-centration. In Fig. 3 the quantity (Cp,Φ − C0

p,2), where C0p,2 is

the infinite dilution value as a function of m, is reported forDEDAB and compared with that of DTAB, from Ref. [32]. Thetrends are nicely comparable so that we can assume that themethylene group added to the polar head has a minor effect onthe apparent molar heat capacities of the solutions.

4.2. Volumes and compressibilities

In Figs. 4 and 5 the trends as a function of concentrationof apparent molar volumes and of the apparent molar adiabaticcompressibilities, respectively, here reported for the first time,are shown for DEDAB, in comparison with those of DTABfrom Ref. [19]. Volumes of DEDAB are greater than those ofDTAB by about 15 cm3 mol−1, both at infinite dilution and atmicellar phase, a value in agreement with that of a methylenegroup in the alkyl chain [14]. Peculiarities, indicating the oc-currence of a phase transition in solution, were not detectedin the trends of volumes or compressibilities. It is noteworthythat apparent molar compressibilities of DEDAB and DTAB areoverlapping (Fig. 5), indicating that the addition of a methylenegroup in the polar head does not affect the compressibility ofthe surfactant in micellar solution.

The overall compressibility of the solution depends both onthe intrinsic compressibility of the solute and the compressibil-ity of the solvent. It has been suggested [19] that at infinite

Fig. 4. Apparent molar volume of DEDAB (") and of DTAB (P) from Ref. [19]as a function of surfactant concentration at 298 K. Also included are resultsfrom the fit of the experimental data by Eq. (11) (—).

Fig. 5. Apparent molar isoentropic compressibilities of DEDAB (") and ofDTAB (P) from Ref. [19] as a function of surfactant concentration at 298 K.Also included are results from the fit of the experimental data by Eq. (11) (—).

dilution the intrinsic compressibility of quaternary ammoniumsalts is a minor contribution compared to the change in thecompressibility of water, due to the electrostriction and to theformation of a shell of higher density structured water aroundthe alkyl chain. When micelles are formed, the compressibil-ity of the solution increases, mostly due to the disruption of thecavity hosting the polar moiety of the surfactants and the voidspace in the inner micelle, whereas it has been shown [33] thatthe micelle remains hydrated at the depth of two –CH2– groupsin the case of DTAC. This means that a methylene group, whenadded to the quaternary nitrogen as in DEDAB, does not changesignificantly its hydration state when micelles are forming, sothat its effect on the overall compressibility of the solution isnegligible. In this way we can explain why apparent molar com-pressibilities of DEDAB and DTAB show the same trends vs m.In contrast, when a –CH2– is added to the hydrophobic chain,it is generally accepted that its group contribution to the molarcompressibility is negative (−1.9 × 10−4 cm3 bar−1 mol−1 inRef. [34], −1.6 × 10−4 cm3 bar−1 mol−1 in Ref. [21]) at infi-

E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307 305

nite dilution. At micellar phase, the –CH2– group contributionbecomes positive: the value of 1.5 × 10−3 cm3 bar−1 mol−1

is reported in Ref. [19] for the change in adiabatic molarcompressibility upon micellization of alkyltrimethylammoniumbromides.

As outlined above, the trends of apparent molar heat capac-ities vs m are also the same: in fact, heat capacities are alsorelated to the number of water molecules involved in the hy-drophobic processes in solution [36], little affected by the sub-stitution of a methyl group by an ethyl one on the polar head.In summary, we can say that this substitution affects to a signif-icant extent the properties which are the first derivatives of thefree energy, but does not affect the second derivatives.

4.3. Change in thermodynamic properties upon micellization

The changes in thermodynamic properties for micelle for-mation of DEDAB, reported in literature, are very few andcontradictory. The thermodynamic behaviour of DEDAB wasinvestigated by Mehta et al. [12] by measuring the cmc as afunction of temperature by conductivity measurements. Ther-modynamic parameters for micelle formation were estimatedby applying the charged pseudo-phase separation model, ob-taining a positive value for �Hm (+0.14 kJ mol−1) at 298 K,so that the formation of micelles is entropy driven at this tem-perature. The same experiments have been done by Galán et al.[22], finding that the micellization process is exothermic in thewhole temperature range studied. The U shaped trend of cmcvs T suggests, in contrast, the presence of a minimum in thecurve of �Gm/T vs 1/T : this is to say that a temperature atwhich �Hm vanishes must exist in the temperature range stud-ied. We have tried to recalculate thermodynamic parameters formicelle formation (Eqs. (7)–(10)) from the values of cmc andα (degree of counterion dissociation) obtained from the plots inFig. 2 of Ref. [22], both taking into account the degree of coun-terion binding (Eq. (7)) and without considering it (Eq. (8)), as

(7)�Gm = (1 + β)RT ln cmc,

(8)�Gm = RT ln cmc,

(9)�Hm = [∂(�Gm/T )/∂(1/T )

]P,

(10)�Sm = (�Gm − �Hm)/T .

The ln cmc vs (1/T ) was interpolated by a second order poly-nomial. The results of our analysis are shown and compared inFig. 6. We can observe that in both cases the enthalpy of mi-celle formation vanishes at a temperature around 290 K and isnegative at room temperature.

In order to obtain the enthalpy change upon micellization,�Hm from our experimental data, we have applied a pseudo-phase transition model, in which the aggregation process is con-sidered as a phase transition, taking place at equilibrium. In thismodel, it is assumed that, at the cmc, the partial molar proper-ties present a discontinuity due to the formation of the pseudo-phase. The micellization parameters are obtained by extrapolat-ing at the cmc the trends of partial molar properties before andafter cmc [9,13–17]. Following this procedure, it results that�Hm = −1.8 kJ mol−1 at 298 K and �Hm = −8.2 kJ mol−1 at

Fig. 6. Changes in free energy, �Gm (1, 2), enthalpy �Hm (P, Q) and en-tropy �Sm (!, "), upon micelle formation for DEDAB, obtained from thedependence of cmc on T (data from conductivity measurements in Ref. [22])by Eqs. (7)–(10). Full symbols represent values obtained from Eq. (7), opensymbols from Eq. (8).

313 K. The value at room temperature agrees quite well withthat obtained from data in Ref. [22] starting from Eq. (8).

The micellization heat capacity, �Cp,m results as being−426 J K−1 mol−1, a little lower, as absolute value, than thatreported in Ref. [1] for DEDAB (−471 J K−1 mol−1, start-ing from Eq. (8)) and in Ref. [32] (−500 J K−1 mol−1) forDTAB, but a little greater than that we obtained from thelinear dependence of �Hm on T from the data in Fig. 6(−334 J K−1 mol−1). Once more, it is interesting to com-pare the change in enthalpy upon micellization for DEDABand DTAB. Surprisingly, considering that DTAB is one ofthe most popular cationic surfactants, literature data are quitecontroversial, the proposed values ranging from −8 [29] to−1.5 kJ mol−1 [36] at 298 K. Moulik et al. [37] have de-termined the solution properties of some cationic surfactants,DTAB included, obtaining −1.77 kJ mol−1 at 303 K, a valuerecently quoted in Ref. [38]. From the dependence of micellarparameters on temperature, reported in the paper of Mata et al.[39], a value of −2.7 kJ mol−1 at 298 K can be obtained (butwith a very small unreliable absolute value of �Cp,m), whereas−2.18 kJ mol−1 at 303 K, has been proposed from calorimetricdata in Ref. [40].

The value in Ref. [36] has been obtained from the datashown in Fig. 2 and, in our opinion, is underestimated. If weadopt the same criteria we used for DEDAB, we obtain �Hm =−2.3 kJ mol−1. In this way, we can evaluate the group contri-bution of the –CH2– group, when added to the polar head ofalkyltrimethylammonium bromides, as +0.5 kJ mol−1 group−1

at 298 K, about half of the effect due to the removal of the samegroup from the alkyl chain [14,40].

Volumetric properties, such as volumes and compressibil-ities, are reflective of the solute–solvent interactions. Theirchanges upon micellization are primarily due to the change instructure of the water molecules involved in the hydrophobiceffect [35] and to the changes in electrostriction of the polar

306 E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307

Table 1Changes in thermodynamic properties upon micellization for DEDAB and DTAB at 298 K, where not differently stated

DEDAB Reference DTAB Reference

cmc (mol l−1) 0.0140(4) [22,25,28] 0.0151 [29]

β 0.72(2) [22,25,28] 0.68 [29]

�Hm (kJ mol−1) −1.8(1) This work From −8 [29,35–40]+0.14 [7] to −1.5−2.1a [22]

�Hm (kJ mol−1) at 313 K −8.2(1) This work−6.8a [22]

�Cp,m (J K−1 mol−1) −426(10) This work −500 [32]−471 [22]−334a [22]

VΦ,cmc (cm3 mol−1) 302.8(1) This work 288.2 [19]

VΦ,m (cm3 mol−1) 309.4(1) This work 294.8 [19]

�Vm (cm3 mol−1) 6.6(1) This work 6.6 [19]

Ks,Φ,cmc (bar−1 cm3 mol−1) −0.001987(10) This work −0.0024 [19]

Ks,Φ,m (bar−1 cm3 mol−1) 0.011729(10) This work 0.0112 [19]

�Ks,m (bar−1 cm3 mol−1) 0.013716(10) This work 0.0136 [19]

a Data recalculated by us (see text).

head and of the counterion. A great change in these interac-tions occurs when the micelle formation begins. In order to ex-tract micellization parameters, the apparent molar volumes andcompressibilities vs m were analysed assuming a pseudo-phasetransition model [41,42]. Following this model, the observedtrends above the cmc can be described by the equation:

(11)XΦ = XΦ,M − (cmc · �Xm) · (1/m).

Here X stands for the volume or for the adiabatic compress-ibility. Knowing the values for the cmc, the values of XΦ,M,the property at micellar phase, and �Xm, the change in prop-erty upon micellization, can be obtained by a least square fit.In Table 1 the values thus obtained, together with the values ofXΦ,cmc, the value at the cmc, obtained by

(12)XΦ,cmc = XΦ,M − �Xm

are reported. The good agreement between experimental andcomputed data is shown in Figs. 4 and 5, in which the solidline represents the computed function. The value of Ks,Φ,cmc,the compressibility at the cmc, obtained by Eq. (12) is nega-tive, as expected. The changes in volume and compressibilityupon micellization are the same for DEDAB and DTAB, show-ing that the addition of the –CH2– group to the polar head doesnot greatly affect the hydration sphere of the molecule either asmonomer or in micelle.

5. Conclusions

Solution thermodynamic properties of DEDAB, obtained bymeans of direct methods, were collected and compared withthose of DTAB, with the goal of enriching the fundamentalunderstanding on self-aggregation thermodynamics of surfac-tants and of providing information for the theoretical evalua-tion of their properties. In fact, for instance, when we try to

predict by means of a group contribution approach the thermo-dynamic properties of Gemini surfactants as a function of thespacer (the alkyl bridge between the polar heads) length, weneed absolutely to know the group contribution of methylenegroup when added to the hydrophilic moiety. The evaluation,by means of direct methods, of the thermodynamic properties ofthe aqueous solutions of DEDAB as a function of concentrationallows for the determination of the effect of the –CH2– group,when added to the polar head of alkyltrimethylammonium bro-mides. The properties most affected by this addition are withoutdoubt the enthalpies and, as a consequence, the entropies: forthe same alkyl chain and counterion the plateau value of the ap-parent and partial molar enthalpies vs m is strictly related tothe modulation of the charge density due to inductive or reso-nance effects, greater in DEDAB than in DTAB. These effectsdo not play a great role in the value of the cmc (i.e. on the freeenergy of micelle formation), since the small decrease in cmcof DEDAB compared to that of DTAB reflects the increase inthe overall hydrophobicity of the molecule. The value of theplateau is similar to that of an alkyltrimethylammonium bro-mide surfactant having eleven carbon atoms in the alkyl chain.Volumes of DEDAB are greater than those of DTAB by about15 cm3 mol−1, both at infinite dilution and at micellar phase,a value in agreement with that of a methylene group in thealkyl chain. The trends of apparent molar heat capacities andcompressibilities vs m are the same as for DTAB: in fact, thesequantities are related to the number of water molecules in-volved in the hydrophobic processes in solution, little affectedby the substitution of a methyl group by an ethyl one on thepolar head. In summary, we can say that this substitution af-fects to a significant extent the properties which are the firstderivatives of the free energy, but does not affect the secondderivatives.

E. Fisicaro et al. / Journal of Colloid and Interface Science 305 (2007) 301–307 307

Acknowledgment

The authors are grateful to the University of Parma (FIL2005) for financial support.

Supporting information

The online version of this article contains additional support-ing information.

Please visit doi: 10.1016/j.jcis.2006.09.063.

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