Physico - chemical characterization of some surfactants in aqueous solution and their interaction with α- cyclodextrin

journalof

MOLECULAR

LIQUIDS ELSEVIER

Journal of MolecularLiquids, 63 (1995) 291-302

Physico - chemical Characterization of Some Surfactants in Aqueous Solution and their Interaction with tz- Cyclodextrin O. Ortona, L. Padua.no, L. Costantino, and V. Vitagliano Dipartimento di Chimica, Universit~ di Napoli, Federico II Via Mezzocannone 4, 80134 Napoli, Italy, Fax 3981 5527771 (Received 25 March 1994; in revised form 21 November 1994)

Summary Aqueous solutions of three surfactants: (1) Sodium 1- hexylsulfonate, (2) Polyoxyethylene - 5 - hexylether, and (3) 1- Hexaminium N,N,N triethyl bromide have been characterized through viscosity, density and calorimetry runs at 25°C. The interaction of these surfactants with ot-Cyclodextrin was also studied calorimetrically. The enthalpy of inclusion and the inclusion constant were measured. .

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Keywords: Density, Viscosity, Calorimetry, Surfactant(s), Cyclodextrin .

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Introduction Cyclodextrins are oligosaccharides having 5 or more glucosidic residues bound in a cyclic chain that generates a hydrophobic cavity where a variety of compounds may be inserted [1]. The hydrophobic cavity of cyclodextrins can provide temporary asylum for hydrophobic parts of molecules dissolved in water. This can result in an increasing solubility and in a higher solution stability of hydrophobic and amphifilic molecules. The solubilizing ability of cyclodextrins is of current pharmaceutical interest in improving drugs transport through living systems. Moreover, the inclusion complexes have been widely used in the field of food industries [2,3]. The large applicative interest in cyclodextrins and their inclusion complexes has favoured the production of large literature on the properties of these systems. With the aim to offer a contribution in this field, we explored the effect of the hydrophilic residue on the cyclodextrin-ligand interaction taking a set of calorimetric measurements. In fact, it is known that molecules having both hydrophobic and hydrophilic groups interact with cyclodextrins through their hydrophobic moieties that occupy the cyclodextrin cavity [1,2]. SSDI

0167-7322 (94) 00791-8

292

On the other hand, molecules having different polar heads and the same hydrophobic residue may interact differently with cyclodextrins leading to a different transport ability of this molecule through a biological medium. This difference may be important from a pharmacological point of view. We made a calorimetric study on the interaction between o~-cyclodextrin and three surfactants having the same hydrophobic tail: R = CH3 - (CH2)5 - and different polar heads: (1) Sodium 1- hexylsulfonate (Na-HS), R- SO3 Na, an anionic surfactant (2) Polyoxyethylene - 5 - hexylether (POHE), R- [CH2 - CH2 - 0-]5 - H, a non -ionic surfactant (3) 1- Hexaminium N,N,N triethyl bromide (HA-Br), R - N(C2Hs)3 Br, a cationic surfactant. The aqueous solutions of the three surfactants were previously characterized through viscosity, density, and calorimetry measurements. The experimental results are reported and briefly discussed in the following. THE BINARY SURFACTANT SYSTEMS Experimental

Materials. a-Cyclodextrin (CD), Sodium 1-Hexyl-sulfonate (Na.HS), Triethylhexyl-ammonium bromide (HA.Br), and Polyoxyethylene-5-hexyl ether (POHE) were purchased from Sigma and used without further purification. The water content of air-equilibrated a-CD was determined by drying at 120 °C under vacuum for several days. The initial tx-CD sample was found to have a 9.501% content of water, in good agreement with the value calculated for the formula o~-CD.6H20 [4]. The molar masses used for CD, Na-HS, HA.Br, POHE, and water are 972.9, 188.22, 266.27, 322.5, and 18.016 (g tool -1) , respectively. Viscosity measurements were performed by using an Ubbelhode viscometer where dilutions were obtained by adding measured volumes of distilled water. Viscosity numbers, (t/t°-l)/m, are shown in Figure 1 as a function of surfactants molalities, t and t ° being the flow times of solutions and solvent. Density measurements were made with an A.Paar densimeter at 25°_+0.001 °C. Equations for the density of solutions are given in Table 1 as a function of molality. The agreement between experimental and computed data is within the fifth decimal figure. Apparent and partial molar volumes were also computed from the density data and fitted with polynomials in ~/m, as proposed by various authors [5], equations are also given in Table 1. Figure 2 shows the graphs of the apparent molar volumes. Calorimeric measurements on the surfactant-water solutions were performed at 25°C with a LKB flow micro-calorimeter by diluting a surfactant solution of molality mi to a final molality mf, where mf -___mi/2. The experimental heat of mixing per Kg of solvent is obtained from the expression :

293

QMIX[(JKg']) = (dQ/dt)/Pw = mf-AHDil.

(1)

where (dQ/dt) is the heat flow(J s "]) and Pw is the total mass-flow rate(Kg s -1) of water through the calorimeter. The experimental data for the three surfactants are collected on Table 2, where AnDil. are the enthalpy changes per mole of solute. 0,00 1.40

0,02

0,04

0.06

0.0e

0,10

I

I

I

I

I

m / ( m a l Kg- Ij POHE

1 ,35

•

•

1 ,30

[(t/t*)K9 B o l ' Z

0,I 2

0,14

I

0,16

I

Tc..,.c.

•

•

! ]/m I ,25, / /

0.65

Naris

°[

13

o

.

0 0

0.55 0 DD o

0.45

"

C. II111.¢.

" 0.35

I

0.0

@

@

@

I

0,5

•

HABr

@

c.m.c.

m/(mai

Kg- !)

!

I

I

1 .0

I ,5

2,0

2,5

Figure 1 Viscosity numbers computed in terms of molality. The c.m.c, molalities are also shown.

Viscosity and Volumetric Results Figure 1 shows the graphs of the viscosity numbers computed from the experimental flow times in terms of molality for the three surfactant solutions. These graphs show a discontinuity on the viscosity number trend for each surfactant. This discontinuity can be attributed to the micellization process. The discontinuity is very evident for POHE. The critical micelle concentration observed in these graphs for POHE and NA-HS agree with the literature data [6,7]. The value for HA.Br was assumed at m ___-0.8, There are no data available in the literature for comparison.

294

Table 1 Computed Expressions for Density of Surfactant Solutions a n d f o r A p p a r e n t a n d Partial M o l a r V o l u m e s o f S o l u t e s n

dcalc./g mL-1 = Z

(m/mol Kg -1 )

dk mk

(2)

k=O

do

grnL-1 Na,HS (a) Na,HS (b) POHE (c) HA,Br

0.997043 1.06043 0.997043 0.997043

(a) for m < 0.5

dl

d2

0.05495 -0.20447 0.02242 0.02968

d3

-0.01064 0.33766 -0.00692 -0.00521

d4

-0.00361 -0.15225

c.m.c.

..... .....

..........

0.01047

gmL-I

mol Kg"1

+6.4x 10 -6 +_5.6x10 -5

0.56

+_.5.8x10 -6

0.1 0.8

-0.00862 +l.6x10 -5

mol Kg-I ; (b) for m > 0.5 tool Kg-~ ; (c) for m < 0.08 tool Kg-1

n

Vapp]mL

mo1-1 = V 0 + Z

Vk mk/2 + Ov

(m/mol Kg -1)

(3)

Vk mk/2 -+ CYv

(m]m°lKg -1)

(4)

k=l n

V---z/mL mo1-1 = V ° + Z k=l

V1

V°

Vz

V3

V4

ml.,mol "1

Ov

mLmo1-1

Na.HS (a) Na°HS (b)

132.35 (_-L-0.22) 132.35 (_+0.22)

7.55 11.32

-11.74 -23.48

7.87 19.68

.... ....

_+0.03

POHE (a) POHE (b)

301.39 (_+0.20) 301.39 (_+0.20)

..... .....

25.59 51.18

.... ....

237.7 692.1

_+0.20

11.73 17.59

-27.66 -55.32

13.66 34.15

.... ....

_-t-0.06

HA-Br (a) 235.61 (-+0.20) HA-Br (b) 235.61 (-+0.20) (a) coefficients of Eq.(3) for

Vapp;

(b) coefficients of Eq.(4) for partial molar volume

295

Table 2 Calorimetric data on Surfactant Solutions at 2 5 ° C Na.HS

POHE

Sper. m i mf AHDil. tool Kg "1 0.9883 0.9428 0.8109 0.6846 0.5635 0.4432 0.3361 0.3361 0.2740 0.2132 0.1618 0.1618 0.1058 0.0850 0.0458

0.4660 0.4487 0.3832 0.3306 0.2738 0.2167 0.1657 0.1655 0.1360 0.1063 0.0809 0.0793 0.0521 0.0419 0.0269

HA.Br

Calc.

Sper.

AHDil. J tool "1

mi mf tool Kg -1

-2953.0 ...... 0.21800 -2928.0 ...... 0.10270 -2740.0 ...... 0.08090 -2200.0 ...... 0.05980 -1716.0 ...... 0.03970 -1228.0 -1243.5 0.02980 -894.0 -904.4 0.02420 -931.0 -905.5 0.01912 -737.0 -736.0 0.01419 -604.0 -589.4 0.00947 -474.0 -474.5 0.00717 -469.0 -485.0 0.00581 -347.0 -358.3 0.00459 -288.0 -309.5 -219.0 -164.8

Calc.

Sper.

AHDi1 AHDil" J tool "1

0.10122 -5739.0 .... 0.04923 -1420.2 -1309.8 0.03902 -829.0 -841.0 0.02883 -606.8 -567.8 0.01925 -387.1 -401.5 0.01447 -301.1 -327.2 0.01170 -264.3 -282.8 0.00977 -195.6 -223.3 0.00728 -151.2 -175.4 0.00484 -166.1 -125.0 0.00368 -136.0 -97.2 0.00299 -133.8 -80.1 0.00235 -128.2 -64.7

mi mf tool Kg "1 0.9864 0.7446 0.5945 0.5215 0.4184 0.3269 0.2692 0.1655 0.129 0.08125 0.06536 0.03029 0.0168

0.4576 0.3476 0.2869 0.2535 0.2059 0.1623 0.1343 0.08134 0.06364 0.04029 0.03254 0.0151 0.0084

Calc. AHDil. AHDil" J tool -1

-1455.2-1556.0 -1063.2-1040.7 -700.1 -712.6 -539.5 -564.5 -360.7 -364.9 -195.7 -204.2 -144.0 -115.0 0.0 8.3 31.5 33.3 55.9 44.0 53.2 40.3 17.0 15.6 0.0 -0.6

CoefficientsofEq.(6) A1

Na.HS POHE HA.Br

1969 ...... 1969

A2

5280 30800 -11040

A3

-8400 ........ 19180

A4

8590 285000 -8097

A5

A6

................ ........ 210000 ................

Figure 2 shows the graphs of the apparent molar volumes for the three surfactant solutions. A few comments can be made about the Vapp shown in this figure. (1) According to the Debye-Huckel theory, the following limiting expression holds for the partial molar volumes of electrolytes V2- V~ = Sv~/m

(5)

where, for a 1-1 electrolyte at 25 °C, Sv = 3.8 (mL Kg 1/2mol -3/2) [8]. However, since the limiting slope of V2 found for electrolyte solutions is generally very scattered, as compared to the value predicted by the DebyeHuckel theory, we did not impose it on the fitting equations.

296

(2) All terms in m of the Vapp. equations (see Table 1) are negative. This is indicative of an overall hydrophobic hydration of the three surfactant molecules. In fact, the overlapping of the hydration cospheres, by increasing concentration, releases structured water to the less structrured bulk, with a consequent volume decrease. The different values of these terms, Na.HS < POHE < HA.Br, are in agreement with an increasing hydrophobicity of the polar heads: -SO3- being an ionized group, (-CHz-CHz-O-)5-H a short chain of hydrophobic methylene groups and polar oxygens, and -N(CzHs)3+a positive charged nitrogen surrounded by three hydrophobic ethyl groups. (3) As can be seen in Fig.2 for the Na.HS, at high concentration the slope of Vapp. increases drastically. This change is connected with the micellization process, in fact it appears at m --- 0.56. Above the c.m.c, the hydrophobic contribution due to concentration changes should be neglegible since micelles interact with the solvent through the polar heads so that a volume contraction due to hydrophobic interactions must be expected only during the micellization process. An extrapolation of higher concentration Vapp.data at infinite dilution suggests a V~ value for the micellized Na.HS o f - 1 2 8 mL mol -~, which is lower than that of the monomeric surfactant (132.3 mL moll). We have only one point at high concentration for POHE and HA.B r, however, an increase on the experimental Vapp. is also present in these systems supporting our interpretation of Na.HS data. The c.m.c, values guessed from the Vapp. graphs of Fig.2 for POHE and HA.Br are also in good agreement with those obtained from viscosity data. Calorimetric

Results

The heat of dilution data were fitted by the equation n

AHDil./(J mo1-1) = ~ Ak(mf ~2 _ mi W2)

(m/mol Kg -1 )

(6)

k=l

where the Ak terms are given in Table 2. The odd terms are absent for the nonionic surfactant POHE. In Eq.(6) A1 = 1969 (J Kg 1/2 tool -3/2) is the limiting value predicted by the Debye-Htickel theory for aqueous solutions of 1:1 electrolytes, in terms of molality [9,10]. The fitting of experimental data was done only using compositions below the critical micelle concentration. It can be noted that the term in m is positive for Na.HS, but negative for HA-Br. This last value is in good agreement with the values for similar compounds given in the literature, 1-Hexaminium N,N,N trimethyl bromide: AE = -10680 (J Kg mol-2), and 1-Hexaminium N,N,N tributhyl bromide: A2 = -10692 (J Kg mo1-2) [10-12]. The apparent enthalpies of dilution are given by n

Lapp/(J mol -l) = • Ak m ~2 k=l

(m/mol Kg -1)

(7)

297

0,00

0,04

0,08

305

O, 12

i

0.I 6

i

0,20

0,24

I

|

304

Yapp303 mL

POHE

tool-z302 301

0

m/molK9 300

I

4

I

I

I

238 237 236 235 234 233

w

138

m

Vapp 137 mL

tool -z136 135

Na.HS

134 133 132 131 130 129 c.m.c.

128 127

~molUZKg -zCz I

0,20

I

I

0,40

0,60

',

i

I

0,80

Figure 2 A p p a r e n t molar v o l u m e s o f surfactants in aqueous solution at 25°C

I ,00

298

The partial molar enthalpies of dilution can be obtained from Eq. (7) by proper derivation. Figure 3 is a graph of the apparent enthalpies of dilution for the three surfactants. 4000 O

3500 3000

7

2500

2000 1500 1000 500 0

/ m o l t*'ZK 9 -tz2

-500

1

0,0

0,2

I

I

0,4

0,6

I

0,8

,0

Figure 3 Apparent heat of dilution of surfactants

T H E I N T E R A C T I O N OF S U R F A C T A N T S WITH ~ - C Y C L O D E X T R I N

The calorimetric runs on the ternary systems ~-CD - surfactant (L) - water were taken by using a batch calorimeter [TAM from Thermometric]. An initial solution of surfactant of molality mi.L was titrated by adding a CD solution of molality mi,CD tO obtain a solution (mcD,mL) after each CD addition. The calorimetric study was done with the aim of distinguishing, if any, the effect of hydrophilic heads on the binding constants. Assuming a 1:1 inclusion complex between CD and L, according to the equilibrium C D + L = L-CD + zM-I° (8) the experimental heat of mixing is due to the contribution of three terms: (a) the heat of CD dilution, (b) the heat of surfactant dilution, and (c) the heat of L inclusion into CD, according to equilibrium (8): QMIX[(mi.CD),(mi.L) -~ (mcD,mL)] =

(9) AHDil,CD(mi,CD--+ mCD) + AHDil,L(mi,L--+ mL) + A H *

299

The first and second term on the right hand side of Eq.(9) were computed through the proper expressions for the heat of dilution of cyclodextrin [14] (-3920 J Kg mol-2), and surfactant [Eq.(7)], assuming that the enthalpies of mixing in the ternary solution were the same as those in binary solutions. The standard molal enthalpy of inclusion, AH °, is given by AH ° = AH*/mL.CD

(10)

where mL.CDis the molality of the complex L-CD present in solution after the mixing process. This term cannot be directly measured in our calorimetric experiments. However, when a large excess of cyclodextrin molecules is present, mL----~mL.CO,SO that AH ° = (AH*)sAT/mL (11) where (AH*)sAT is the enthalpy of inclusion measured at CD saturation. In absence of any information concerning the solutes activity coefficients we may def'me only a molal binding constant: Km =

mL'CD

(12)

(rncD- mL.CD)(mL - mL.CD )

By substituting mL.CDfrom this Eq. into Eq.(10) and rearranging the terms, one obtains the following expression: m L / A H * --

1/[AH ° Km (mCD - mL.CD)] + 1/AH °

AI-I* = AH° Km(mcD - mL'CD) mL 1 + Km(mcD- mL.CD)

(13) (14)

Eqs.(13) and (14) have the form of a Langmuir isotherm (see Figs.4 and 5), the equilibrium constant and AH ° can be obtained from the plot of mtJAH* vs 1/(mCD - mL.CD) which is linear, as shown in Fig.4. The amount of free ligand (mol Kg-~ ) is given, at any molality by the expression mL.CD= mL zSJ-I*/(AH*)SAT

(15)

Actually, neither the molality of the complex nor the enthalpy change at saturation can be measured directly, AH ° and Km are therefore obtained from Eqs.(13) and (15) by iteration; their numerical values for the three surfactants are given in Table 3. This table also gives the free energy AF°. the entropy AS °, and the excess entropy ASE of inclusion.

300

Table 3 Surf. Na.HS POHE HA-Br

All °

Km

AF °

kJ tool "1

Kg tool "1

kJ tool "1

- 22.7+0.15 - 16.9_+0.6 - 16.1 + 0.02

114+2 180+__ 11 268 + 2

-11.7 -12.9 -13.9

AS °

ASE

kJ K'lmol "1

-0.037 -0.014 -0.008

-0.003 +0.02o +0.026

We found two papers available in the literature giving the binding constant of Na.HS to c~-CD. The first one [16] deals with stopped flow conductance measurements, the results are in good agreement with ours (Kc_=_100L mol-1). The second one [17], dealing with a simple conductance technique, gives an extrapolated binding constant significantly larger than ours (Kc = 450 L mol-1). The thermodynamics of inclusion process involves various entropic and enthalpic terms connected with the changes in the solvent structure, with the water release from the CD cavity, with the specific interactions between solute and solvent and between solute molecules themselves. A detailed interpretation of all these steps, in terms of the data given in Table 3, would risk to be a mere speculation. However, from the data of Table 3 although largely approximate, one can argue that the driving force of the inclusion process is mainly enthalpic. The entropy terms are all slightly negative. If the entropies are corrected for the ideal term (cratic entropy ASc = - 0.033 kJ K -1 mo1-1 ) the excess entropies increase in the order Na.HS < POHE < HA.Br. This, of course, is indicative of the fact that the inclusion affects largely the hydration shell of the polar heads releasing some organized solvent molecule to the bulk. Finally, the higher negative value of the inclusion enthalpy of Na.HS can be attributed to some hydrogen bonding between the -SO3 oxygens and the cyclodextrin -OH groups, this possibility is in agreement with the lower value of the excess entropy of inclusion. Note: Additional tables of experimental data can be obtained from the authors on request. .

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Acknowledgement This research was supported by the Italian M.U.R.S.T. and by the Italian C.N.R.

301 0

50 ,

150

I

,

I

o,o -1,00e-4

,

~

,

0,0

{

HA.Br

~

200

I

~ -- ~ . "~

--

~

-1.00e-4

- 2.00e-4

mt/AH*

(molzKg-Ij-9

-3.oo,-,_,.oo._,

I No. I/mcD

,

0

200

,

(free)/(Ko

4.00

,

-,.oo.-, mol-t., I

,

600

-5.00e-4

800

1000

Figure 4 Graph of Eq.(13) for the three SurfactantsExperimental Data and Computed Curve

-5000

AH*/m t (J xg moz-e)

-15000

-20000 0,00

m co (free)//~m01

K9 - !/

i

i

i

o ,oi

o .02

o .03

Figure 5 Enthalpy of Inclusion of Surfactants into ~-Cyclodextrin Experimental Data and Computed Curves Eq.(14)

o ,04

302

References

(1) J.Szejtli, Cyclodextrins Technology, Kluver Acad.Publ.,Dordrecht 1988 (2) New Trends in Cyclodextrins and Derivatives, D.Duchene Ed., Editions de Sant6, Paris 1991 (3) J.Szejtli, J. Inclus. Phenomena, 14, 25-36 (1992) (4) L. Paduano, R. Sartorio, V. Vitagliano, and L. Costantino, Ber. Bunsenges. Physik. Chemie, 94, 741-45 (1990) (5) J.F.Desnoyers, P.De Lisi, C.Ostiguy, and G.Perron, Solution Chemistry of Surfactants, Vol.1, K.L.Mitlal Ed., Plenum, New York 1979 (6) A.L.M.Lelong, H.V.Tartar, and E.C.Lingafelter, J.Am.Chern.Soc., 73, 5411 (1951) (7) M.Donbrorw, J.Pharm.Pharmacol., 15, 825 (1963) (8) H.S.Hamed and B.B.Owen, The Physical Chemistry of Electrolyte Solutions, Reinholds 1958, Chapt.3 (9) Ref.8, Chapt.8 (10) P.A.Leduc, J.L.Fortier, and J.F.Desnoyers, J.Phys.Chem., 78, 1217 (1974) (11) J.L.Fortier, P.A.Leduc, and J.F.Desnoyers, J.Solution Chem., 2,467 (1973) (12) S.Lindenbaum, J.Phys.Chem., 70, 814 (1965) (13) R.B.Cassel and W.Y.Wen, J.Phys.Chem., 76, 1367 (1972) (14) G.Barone, G.Castronuovo, P.Del Vecchio, V.Elia, and M.Muscetta, J.Chem. Soc. Faraday Trans., 1, 82, 2089 (1986) (15) M.Eftink and P.Biltonen in A.F.Beezer(Ed.) Biological Microcalorimetry, Acad.Press, London 1980 (16) T.Okubo, Y.Maeda, and H.Kitano, J.Phys.Chern., 93, 3721 (1989) (17) E.S.Aman and D.Serve, J.CoHoid Int. Science, 138, 365 (1990)