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Solubilization of 2-propanol and pentanol in sodium dodecyl sulfate micelles: A thermochemical study
SolubiUzation of 2-Propanol and Pentanol in Sodium Dodecyl Sulfate Micelles: A Thermochemical Study INGER
JOHNSON
AND G E R D
OLOFSSON
Division of Thermochemistry, ChemicalCenter, Universityof Lund, Lund, Sweden Received October 7, 1985; accepted February 2, 1986 The solubilization of 2-propanol and pentanol in sodium dodecyl sulfate (SDS) solutions has been studied by titration calorimetry over a wide range of alcohol concentrations in three different SDS solutions. Measurements were made at 25°C and in addition at 35°C for pentanol. The measured enthalpy changes are close approximations of the partial molar enthalpies of solution of the alcohols in surfactant solution. Reported values of the partition coefficient/~ derived from self-diffusionmeasurements using the Fourier transform NMR pulsed gradient spin-echo method were used to interpret the results. In the calorimetric measurements pentanol dissolves predominantly in the micellar phase, while a larger part of added 2-propanol remains in the aqueous phase. To separate the solubilization effects partial molar enthalpies of transfer A~ic/7of alcohol from pure liquid to micellar phase were derived assuming the distribution of alcohol between aqueous and micellar phase to be a simple partition equilibrium. The limiting value of A~iC/7at low alcohol content is -4.4 _+ 0.2 kJ mole-1 for pentanol and -6.5 _ 0.5 kJ mole-1 for 2-propanol. The derived values of A]~ic/1for pentanol show a pronounced variation with the composition of the micellar phase. The increase is almost linear up to a molar ratio 2:1 of alcohol:SDS in the micellar phase, indicating a smooth change in the milieu for the sohibilizate from discrete micelles with low alcohol content to alcohol-rich apolar domains. The heat capacity change for the transfer of liquid pentanol to the micellar phase is zero or close to it over the concentration range studied, which indicates that the alkyl chains are dissolved in a nonaqueous environment. © 1987Academic Press, Inc.
INTRODUCTION
tively. A l t h o u g h the solubilization p h e n o m e n o n has been extensively s t u d i e d b y different m e t h o d s , few techniques are available for t h e study o f the d i s t r i b u t i o n o f t h e solubilizate. R e c e n t l y a m e t h o d has b e e n d e v e l o p e d w h i c h p r o v i d e s direct i n f o r m a t i o n a b o u t the partit i o n i n g o f solubilizate b e t w e e n micelles a n d water (3-6). It is b a s e d o n m e a s u r e m e n t s o f i n d i v i d u a l self-diffusion coefficients in m u l t i c o m p o n e n t systems using a n N M R F o u r i e r t r a n s f o r m p u l s e d - g r a d i e n t spin-echo m e t h o d . T h e solubilization in s o d i u m d o d e c y l sulfate (SDS) solutions o f a n u m b e r o f solutes such as alcohols, ketones, a n d a r o m a t i c h y d r o c a r b o n s was studied b y this m e t h o d (4). T h e twostate p a r t i t i o n e q u i l i b r i u m m o d e l was f o u n d to give a satisfactory d e s c r i p t i o n o f the distrib u t i o n a n d values o f Kc were derived. T h e dist r i b u t i o n o f alcohols in SDS solutions has also
T h e solubility o f n o r m a l l y o n l y slightly soluble organic c o m p o u n d s in a q u e o u s s o l u t i o n is substantially increased w h e n surfactants are p r e s e n t in c o n c e n t r a t i o n s a b o v e their cmc. This solubilization p h e n o m e n o n is o f basic interest in the study o f m i c e l l a r a n d m i c r o e m u l s i o n solutions a n d it also has n u m e r o u s practical a p p l i c a t i o n s (1, 2). W i t h i n t h e phase s e p a r a t i o n m o d e l for micellization, solubilization is regarded as a partitioning o f the solute b e t w e e n the m i c e l l a r a n d a q u e o u s phases. A t low c o n c e n t r a t i o n s , a d i s t r i b u t i o n coefficient is a s s u m e d to describe the solubilization behavior: KC~ S S CmiffC aq [ 1] where c~i¢ a n d c~q denote solute c o n c e n t r a t i o n in the m i c e l l a r a n d a q u e o u s phases, respec56 0021-9797/87 $3.00 Copyright © 1987 by Academic Press, Inc. All fights of reproduction in any form reserved.
Journal of Colloid and Interface Science, Vol, 115, No. 1, January 1987
SOLUBILIZATION OF ALCOHOLS been derived from vapor pressure measurements (7). Values of K¢ for pentanol, hexanol, and heptanol determined by the two methods are in good agreement. We have chosen to study the solubilization of 2-propanol and pentanol in SDS solution by titration calorimetry to gain further information about the thermodynamics of the solubilization process. Solution calorimetric measurements have been made previously on the solubilization of alcohols in surfactant solutions, but they have been limited to only one alcohol concentration (8-11). Thus the effect of increasing solute concentration and accompanying change in composition of the micellar phase has not been studied. To get as detailed information as possible, we have measured differential enthalpies of solution of 2-propanol and pentanol from 0 up to 1 mole alcohol/kg water in SDS solutions of varying molality. The amount of alcohol added in each titration step was small and the measured enthalpy changes are good approximations of the partial molar enthalpies of solution. Therefore, they reflect the conditions of the solute in the solution at the composition of the experiment. Values of the partial molar enthalpies of transfer of liquid alcohol to the micellar phase were derived assuming the distribution of alcohol between the aqueous and micellar phases to be described by partition coefficients derived from the NMR self-diffusion measurements. EXPERIMENTAL
Materials. Specially pure (>99%) sodium dodecyl sulfate (SDS) (BDH, Poole, UK) was used without further treatment. Solutions of SDS were prepared by weight using reagentgrade water produced by a Milli-Q System (Millipore AB, Grteborg, Sweden). Pentanol (BDH, Poole, UK) was fractionally distilled and purity checked by analytical gas chromatography. The sample used was judged to be better than 99.8% pure, with a water content less than 0.05%. Analytical grade 2-propanol of purity better than 99.7% (Merck, Darmstadt, Germany) and analytical grade sodium
57
perchlorate (NaC104. H20) from Fluka AG (Buchs, Switzerland) was used without further treatment. Calorimetry. Titration calorimetric measurements were made using either the LKB Batch Microcalorimeter modified for titration (LKB Model 2107-350) or an LKB 8700 reaction-solution calorimeter with 25- or 95-cm 3 glass reaction vessels. In the batch microcalorimetric experiments, 5 cm 3 of SDS solution was put into the reaction cell (made of 18carat gold) and the same amount of pure water in the reference cell. In each titration step l0 mm 3 of alcohol was injected in the reaction cell and simultaneously l0 mm 3 of water in the reference cell. In the experiments using the reaction-solution calorimeter, the sample was titrated into the glass vessel by means of a thin steel tube (length 1 m, i.d. 0.15 mm) fastened by epoxy to a Hamilton 100-mm 3 gas-tight syringe (1710 LT). The capillary tube passes through a tube (i.d. 1 ram) in a slit at the inside of the stirrer holder and ends about 30 mm below the surface of the calorimeter liquid. The syringe is motor-driven, and the samples were injected at a rate of 0.10 mm 3 s-1. The sample size varied between 12 and 20 mm 3. To prevent the small drop of pentanol introduced into the calorimeter from creeping up to the surface before it dissolves, a small Teflon cap was mounted on the capillary tube 5 mm from the end. The droplet is trapped under the cap and is completely dissolved in the aqueous solution without evaporation loss, etc. RESULTS Titration experiments consisting of successive additions of 10 mm 3 of pure pentanol, PeOH, to 5 cm 3 SDS solution have been made at 25.0 and 35.5°C using the batch microcalorimeter. Three different SDS solutions having molalities 0.089, 0.182, and 0.370 mole kg-1 were used. Each titration series consisted of 60 or 80 consecutive injections which increased the alcohol content of the calorimeter solution from 0 to 0.9-1.0 mole (kg water) -x. The change in alcohol content was small in Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
58
JOHNSON AND OLOFSSON
each step, being about 0.02 mole kg -1. Therefore the measured enthalpy changes A/l(obs) calculated per mole of added alcohol are close approximations of the partial molar enthalpies of solution of pentanol in the SDS solutions. The results are summarized in Figs. 1a and b where 2x/l(obs) is plotted against the molality of pentanol at the two temperatures. Results of measurements in 0.089 m SDS have been deleted for pentanol content above 0.4 m as it is uncertain if equilibrium was reached under the conditions of the experiments. Visual inspection of samples containing 0.089 m SDS to which varying amounts of pentanol were added showed that the dissolution of added alcohol became slow and took up to 1 h after shaking to become clear when the alcohol content exceeded about 0.5 m. At lower alcohol content and in solutions with higher SDS content, the dissolution was rapid, giving
2 o E
b
t -35.5 °C
1
0.182
o
0mi 2
• ~=~A-===~
i"~ .A¢a ~ a ' *
t
0.4
,, ," "
l*
=•=A
'
0 p ~ ~ ' ~ m"
,eeee
tel•
~'1 m0
112
m/mol kg-1
f~ -1 -2
\ " II||~" lie•= •
°° ~0.370
ill•el e
_51
a
t =25.0 °C , , , = A , , ,`,= 0.2
-2
0.4
~=~A=ee
0.6
A,-=" ,~'== ee 0.182 ,-= e•=e , ' = •e •e" e
1.0 1.2 rn/mol kg-1
0.370
.=i"e=,. " ='=~ 0.089 -5 -(]
FIG. 1. Measured enthalpy changes per mole of added pentanol A/l(obs), as a function of final amount of pentanol in each titration step (a) at 25.0°C, (b) at 35.5°C. The molalities of the different SDS solutions are indicated in the figures. Journal of Colloid and lnterface Science, Vol. 115, No. 1, January 1987
-8 0,381
E -10
•eeeeeeeelleeeeeee*•eeeee
eee•ee
e•ee••
• • •
o
0.150
-12
o12
'
o15
'
11o
' m Imol kg -1
FIG. 2. Measured enthalpy changes per mole added 2-propanol, A//(obs), as a function of final amount of 2-propanolin each titration step at 25.0°C. The molalities of the different SDS solutions are indicated in the figure.
clear solutions over the whole range of pentanol concentrations. Measurements were made at low alcohol content using the reaction-solution calorimeter with a 25 cm 3 vessel. There is good agreement between results from the reaction-solution calorimetric and batch microcalorimetric measurements in overlapping concentration regions. Batch microcalorimetric results of step-wise additions of 2-propanol to SDS solutions of molalities 0.150, 0.250, and 0.381 mole kg -1 at 25.0°C are summarized in Fig. 2. Least-squares fitting of second-order polynomials in alcohol molality mA to observed enthalpy changes gives a satisfactory representation of the experimental results. Coefficients a, b, and c from fits of A/~(obs) = a + bmA + Cm2Ato results of the various titration series are shown in Table I. The standard errors of the fits were 0.05 kJ mole -1 or less, which is the same as the reproducibility of the calorimetric measurements. The empirical relations reproduce the calorimetric results of alcohol contents up to molalities mA(max) indicated in the last column of Table I. Experiments were also made on dissolving the alcohols in pure NaCIO4 solutions of varying molality. Each series of measurements consisted of three consecutive additions of 26.68 m m 3 alcohol to 95.5 cm 3 of water or NaC104 solution. The resulting solution enthalpies are shown in Table II. The measured enthalpy changes can be assumed to refer to infinitely dilute solution, since the alcohol molality did not exceed 0.001 mole kg -~. The
59
SOLUBILIZATION OF ALCOHOLS TABLE I Values of the Coefficients in the Relations A/t(obs)/kJ mole -1 = a + bmA + Cm2A
Solute Pentanol
Temperature (°C)
SDS (mole/kgH20)
a
25.0
0.0890 0.182 0.370 0.0890 0.182 0.370
-5.67 -5.04 -4.69 -4.38 -4.27 -4.26
0.150 0.250 0.381
- 11.52 -10.81 -9.99
35.5
2-Propanol
25.0
value of the solution enthalpy in pure water is in excellent agreement with values determined using two other calorimetric methods at our laboratory and with previously reported values (12).
Calculation of Transfer Enthalpies If the micelles are regarded as a separate phase, the alcohol A added in the calorimetric
c
mh(max) (mole/kgH20)
0.939 -0.439 -0.441 -8.64 -4.22 -1.76
0.38 0.70 0.76 0.30 0.70 0.80
b 7.09 6.31 4.93 11.56 9.50 6.07 0.242 0.180 0.0588
0.616 0.501 0.641
1.2 1.2 1.2
experiments can be assumed to be distributed between the micellar and aqueous phases: A(1) --~ aA(mic) + (1 - .)A(aq)
[2]
where c~ denotes the fraction of alcohol solubilized in the micelles. A distribution coefficient Kc [1] is assumed to describe the distribution of the solute between the two phases. Stilbs found this two-state model to give a sat-
TABLE II Enthalpies of Solution of Pentanol and 2-Propanol in Pure Water and in NaC104 Solutions at Infinite Dilution of Alcohol m(NaCIO4) Solute Pentanol 25.0°C
35.0°C
2-Propanol 25.0°C
0
-7.93 + 0.009 a -7.99 b -7.93 b -4.76 + 0.14 -4.82 b
-13.16 + 0.06 -13.10 c -13.09 a -12.98 e -12.88 f
0.090
0.184
0.377
-7.53 + 0.15
-7.14 ± 0.15
- 6 . 1 4 + 0.15
-4.46 + 0.10
-4.23 + 0.10
-3.33 ___0.10
-12.68 + 0.10
-12.44 + 0.10
-11.67 + 0.10
a Error limits are expressed as twice the calculated standard deviation of the mean. b Ref. (12). *Ref. (13). dRef. (14). *Ref. (15). fRef. (16). Journal of Colloid and Interface Science, Vol. 115,No. 1, January 1987
60
JOHNSON AND OLOFSSON
isfactory description of the partitioning of various solutes in SDS solutions from N M R self-diffusion measurements (4). Four different solutes including pentanol were studied in solutions containing between 0.04 and 0.25 mole liter -1 of SDS and the concentration of two solutes, benzene and benzyl alcohol, was varied between 0.03 and 0.24 mole liter -1, while the SDS concentration was kept at 0.24 mole liter-1. Stilbs has also studied the solubilization of n-alcohols at higher solubilizate and SDS concentrations (5). For n-butanol there is a good agreement between the value of Kc derived at low concentrations and values derived for solutions containing up to 0.46 mole liter-~ of SDS and about 1 mole liter -1 of butanol. The self-diffusion behavior of the different species indicates that up to these concentrations the system is best described as a micellar solution (5). The value ofK~ = 49 for pentanol derived for a solution containing 0.34 mole liter -~ ofpentanol and 0.25 mole liter -1 of SDS is in satisfactory agreement with the value of 55 _+ 3 found at low alcohol concentration (4). At higher alcohol and SDS concentrations the self-diffusion results for SDS indicate a breakdown of the micelles. However, the self-diffusion of the solubilizate remains constantly slow irrespective of solution composition, which can be seen as indicating relatively constant alcohol partition equilibria between polar and apolar domains (5). Thus it appears that the simple two-state equilibrium model can describe the solubilization of butanol and pentanol in solutions containing up to at least 0.46 and 0.25 mole liter-~ of SDS, respectively, as long as the molar ratio of alcohol:SDS in the mixed micelles does not much exceed 1:1.1 At higher concentrations the alcohol is still solubilized in apolar domains with a distribution resembling that at lower concentrations. Assuming the two-state partition equilibrium model to be applicable, the observed en-
thalpy change A/l(obs) in our experiments can be expressed as
A/~(obs) = aA[nic/-I-t- (1 - a)A~q/~.
[31
The partial molar enthalpy of solubilization of A into the micellar phase from the liquid state is denoted ~{nic/~ and the partial molar enthalpy of solution of liquid A in the aqueous phase is denoted A~q/1. The degree of solubilization a was calculated using values for the partition coefficient Ko determined by Stilbs (4) from N M R self-diffusion measurements. The reported value for pentanol is 5 5 + 3 and for 2-propanol J.J-l.stZ t:+2.5. Values of a were calculated according to a/(1 - a) = Kc. Vmio/Vaq using the value of 250 cm 3 (mole SDS) -l for Vmi~ (4, 20) and 18.0 cm 3 (mole H20) -1 for Vaq. This means that Vmio and V,q are proportional to the amount of SDS and water in the solutions and may differ from the actual volumes, particularly in the more concentrated solutions. The concentration of free SDS is assumed to be equal to the cmc, 8 × 10-3 mole kg -1. This is an overestimate of the monomer concentration, but since it is not known how it varies with alcohol and SDS concentration, this value has been used throughout. The error introduced in Vmi~ is small and will not significantly influence the results. Changes in monomer concentration with alcohol content could give an enthalpy contribution to the measured A/t(obs). This effect is considered to be less than - 0 . 5 ld mole -1. The estimate was made assuming the monomer concentration to decrease to the same extent as has been observed close to the cmc (7) and using an enthalpy of micelle formation o f - 5 . 0 kJ/mole. 2 The self-diffusion measurements were made in D20 but a study of the solubilization of nalcohols in sodium decanoate micelles in both D20 and H 2 0 indicated that partition coefficients are the same in the two solvents (23). The effect of temperature on K~ was found to
I The abrupt changes in solubility and K~ observed 2 The enthalpy of micelle formation of SDS in pure for hexanol and heptanol at certain SDS concentrations (17-19) has not been observedfor butanol and pentanol water is closeto zero at 25°C and -5.1 kJ mole-1 at 35°C (21, 22). (5, 18). Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
61
SOLUBILIZATION OF ALCOHOLS
be small (4) and for pentanol we have used the same value at 25 and 35°C. The solution enthalpy in water, A~q/~, will vary with the composition of the aqueous phase. In evaluating A~i~/laccording to relation [3], we have assumed that the ionic medium effect on A~q/~ can be estimated from the enthalpies of solution of the alcohol in NaC104 solutions having the same molality as the SDS solutions. Further, it is assumed that the variation of A~q/-I with alcohol concentration in the aqueous phase is the same as in pure water. Estimates of the partial molar enthalpies of solution in water were based on results from dilution experiments on C1-C4 alcohols (12). Values of the pairwise and triplet interaction coefficients for 2-propanol were assumed to be the same as for 1-propanol, 557 J mole -~ (mole kg-1) and 180 J mole -1 (mole kg-l) -2, respectively. Extrapolation of the values for the coefficients for the C1-C4 alcohols gave for pentanol 1930 J mole -1 (mole kg-1)-1 for the pairwise and 740 J mole -1 (mole kg-1)-2 for the triplet interaction coefficient. Values of the solubilization enthalpy A~i~/lfor pentanol derived in this manner are summarized in Fig. 3, where A~i¢/7 is plotted against the molar ratio pentanol/SDS in the micellar phase, np~oH/nSDS. As can be seen, A~nic/tis closely the same in the three systems at the same composition of the micellar phase. However, A~i¢/7 is not constant but becomes increasingly more endothermic as the alcohol content of the micellar phase increases. The effect of the 10°C change of temperature in A~i~/t is negligible. Values of A~i¢/1 for the solubilization of 2-propanol are summarized in Fig. 4. The degree of solubilization of this alcohol is low and the value of K~ has large error limits. Therefore, the derived values of A~'CH are uncertain. The difference between A~°/t for the different SDS solutions seen in the figure is not significant, as a small change of K¢ will shift the relative position of the curves and can also make them coincide. We have defined the solubilization enthalpy as the partial molar enthalpy of solution of liquid alcohol into the micellar phase A~nic/~.
b t=35.5°C 0.18: ~
P
~-ii/~ a t=25.0°C "7,
0.182 1
.o
~
J
~
<1
3 . 89
J J__ 4 5 rl PeOH/FISDS
-2
FIG. 3. Partial molar enthalpy of solubilization of pentanol into the micellar phase from the liquid state, A[~ic/1, as a function of calculated molar ratio pentanol/ SDS in the micellar phase (a) at 25.0°C, (b) at 35.5°C. The molalities of the different SDS solutions are indicated in the figures.
It is the difference between the partial molar enthalpy content of the alcohol in the micellar phase and the enthalpy content of liquid alcohol: A~ic/~ =/tmic -- HI. Only/~mic will vary with the composition of the system. The solubilization enthalpy can also be defined as the enthalpy of transfer of alcohol from aqueous to micellar phase: Aaq micH- =/tmi¢ --/l~a. Values of A~ ~/lcan be calculated by subtracting from A~i°/1 the appropriate value for the solution enthalpy in aqueous solution, A~q~. The partial molar content of alcohol will vary both in the micellar and aqueous phase with composition. Values of enthalpy changes for the transfer ofpentanol from aqueous solution to the SDS micellar phase, A~¢H have been reported previously (8, 10, 11). These values have been derived from calorimetric measurements at only one alcohol concentration. The value reported by Lisi et aL (10) of 3.44 + 0.06 kJ Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
62
JOHNSON AND OLOFSSON o
0.381 ~
E<~
~..~177JJ.~. 0.250.
-8 t
\0.150
0'5
1'o n iPrOH / rlSDS
FIG. 4. Partial molar enthalpy of solubilization of 2-propanol into the micellarphase fromthe liquid state, A[nie/~, as a functionofcalculatedmolarratio2-propanol/ SDSin the micellarphaseat 25.0°C. The molalitiesofthe differentSDS solutionsare indicatedin the figure.
mole -1 is in good agreement with the value we calculate by subtracting the enthalpy of solution ofpentanol in water (Table II) from the limiting value of A ~ n i c / t , s e e Fig. 3. Other reported values of A~qi~/4do not agree so well, the discrepancy stemming mainly from the use of a different and, as we believe, incorrect value for the enthalpy of solution of pentanol in water. Measurements have been made at 25°C of densities and heat capacities of the ternary systems water-SDS-butanol (24) and waterSDS-2-propanol (25) over the complete alcohol mole fraction (solubility) range. Apparent and partial molar volumes and heat capacities of the alcohols were derived. Assuming that the partition equilibria at low alcohol content can be described by the partition coefficients determined by Stilbs (4), partial molar heat capacities of the solubilized alcohols can be estimated. Values derived from the reported Cp,~ at low alcohol and SDS concentrations are close to the values for the heat capacities of the liquid alcohols, indicating that the heat capacity change for the transfer of liquid alcohol to micellar phase is close to zero for both 2-propanol and butanol. DISCUSSION The observed enthalpies of solution of alcohols in SDS solutions are less exothermic than in pure water or NaC104 solutions due to partial solubilization of the alcohols in the micellar phase. Pentanol and 2-propanol differ Journal of Colloid and Interface Science, Vo]. 115, No. 1, January 1987
significantly in their solubilization behavior, with pentanol showing a strong preference for the micellar phase as indicated by the value 55 for the partition coefficient, while 2-propanol is solubilized to a much less extent, Kc being about 5.5. This means that in the calorimetric measurements on pentanol in the 0.18 and 0.37 mole liter-1 SDS solutions, a major part ofpentanol added in each experiment enters the micellar phase and the observed enthalpy change is dominated by the enthalpy contribution from the solubilization of liquid pentanol into the micelles, A~ic/1. As can be seen from Fig. 1, there is a smooth change in A/4(obs) as the pentanol concentration increases, indicating a smooth change in A~i¢/-I with composition up to high alcohol content, where micellar breakdown occurs (5). The incremental solution enthalpies become steadily more endothermic, indicating that the alcohol continues to dissolve in a nonaqueous environment and that the milieu for the solubilizate changes smoothly from discrete micelles with low alcohol content to alcohol-rich apolar domains. The change of A/4(obs) with temperature is small, indicating that the heat capacity change for the transfer of liquid pentanol to the micellar phase is small. In the experiments on 2-propanol, a larger part of each added portion will remain in the aqueous phase, but also for 2-propanol, A/l(obs) is significantly less exothermic than the solution enthalpy in pure water. To get a clearer picture of the solubilization process, a separation of the enthalpy contributions from solubilization and hydration has been attempted by the derivation of differential solubilization enthalpies A{nic/1as described previously. The calculated values are tentative due to the approximations made, but at least at low alcohol content, where the two-state partition equilibrium model is applicable, the derived values are considered to give a fair representation of the partial molar enthalpies of transfer of liquid alcohol to micellar phase. This assumption is supported by the simplified features of the plots of A~iCH against composition of the micellar phase, Figs. 3 and 4.
SOLUBILIZATION OF ALCOHOLS
63
They show, for instance, that the derived val- for the incorporation of liquid pentanol into ues of A~nic/~ do not vary with SDS concen- the miceUar phase indicates that the solubitration for molar ratios alcohol/SDS of up to lized alcohol is not exposed to water, but that the alkyl chain is in a nonaqueous environat least 1:1 in the micellar phase. As seen from Fig. 3, the partial molar en- ment. This conclusion may also apply to buthalpy of transfer, A~n'cH, of liquid pentanol tanol and 2-propanol, since the heat capacity to micellar phase varies with the composition results from the measurements of the ternary of the micellar phase. At low alcohol content, solutions (24, 25) are fully compatible with A~ic/1 is exothermic, being -4.4 ___ 0.2 kJ very small or zero heat capacity changes for mole-1 in the limit of np~oH/nSDSequal to zero. the transfer of liquid alcohol to miceUar phase The corresponding value for 2-propanol is at low alcohol content. -6.5 + 0.5 kJ mole -1. At higher concentraIt is generally assumed that an additive like tions the values of A~iCH for 2-propanol are pentanol is solubilized in the micelles with the too uncertain to give any information about polar group enclosed at the surface among the their variation with the composition of the surfactant head groups and with the alkyl micellar phase. We refrain from calling the chain immersed in the hydrocarbon case. The limiting values "standard state" values. In our observed zero heat capacity change for the opinion, the concept of standard state is not transfer of liquid pentanol into SDS micelles particularly useful in this situation, as it is not supports this model. The incorporation ofsolpossible to derive standard state values from ubilizate in the micelles will decrease the surexperimental results without using a particular face charge density, the decrease for SDS mimodel for the solubilization process. The re- celles being at least to a first approximation suiting values will therefore depend on the linear with the mole fraction of additive in the model used. Besides, even within the frame- micellar phase (28). As can be seen from Fig. work of a particular model, in the present case 3, A[nic/t increases almost linearly with penthe partition equilibrium model, it is not ev- tanol content in the micellar phase from -4.4 ident how to define meaningful standard kJ mole -1 at very low molar ratio pentanol/ state(s). SDS to zero at a molar ratio of about 2. A striking feature of the results for pentanol In the partition process A(aq) ~ A(mic), is that Arai~/ldoes not vary with temperature, the transfer ofpentanol and 2-propanol from at least not for npeoH/nsosless than 1. If Figs. the aqueous to the micellar phase is endo3a and b are superimposed, the curves coincide thermic at all concentrations studied. The to within +0.2 kJ mole -1 at low molar ratios, positive transfer enthalpies indicate that the which is remarkable considering the simpli- transfer entropy is positive and the driving fying assumptions made. At higher molar ra- factor for the solubilization process. Positive tios, the curves for the two more dilute SDS entropy contributions could stem from the solutions diverge somewhat, but they still lie dehydration of the alkyl chain of the additive within +0.5 kJ mole -1. Thus the heat capacity and a decrease in ion binding to the micelle change for the transfer of liquid pentanol to as the surface charge density decreases. For micellar phase is zero or close to it. This can the solubilization of a series of additives like be compared with the heat capacity change of 1-alkanols, the dehydration entropy can be 300 J K -l mole -1 for the transfer of liquid expected to be proportional to the chain length pentanol to aqueous solution, see Table II. The (29) while the change in surface charge density large positive value of the heat capacity of so- will not vary. The increase in partition coeflution in water stems from the strong temper- ficient with chain length observed for the solature dependence of the interaction between ubilization of 1-alkanols in SDS solution (4, the alkyl groups and the surrounding water 7) is probably governed by the increase of the layer (26, 27). The zero heat capacity change dehydration entropy with chain length. Journal of Colloid and Interface Science, Vol. 115, No. l, January 1987
64
JOHNSON AND OLOFSSON REFERENCES
1. Wennerstr6m, H., and Lindman, B., Phys. Rep. 52, 1 (1979). 2. Lindman, B., and Wennerstr6m, H., Topics Curr. Chem. 87, 1 (1980). 3. Stilbs, P., J. Colloidlnterface Sci. 80, 608 (1981). 4. Sfilbs, P., J. Colloidlnterface Sci. 87, 385 (1982). 5. Stilbs, P., J. Colloidlnterface ScL 89, 547 (1982). 6. Stilbs, P., J. Colloidlnterface Sci. 94, 463 (1983). 7. Hayase, K., and Hayano, S., Bull. Chem. Soc. Jpn. 50, 83 (1977). 8. Aveyard, R., and Lawrance, A. S. C., J. Chem. Soc. Trans. Faraday Soc. 60, 2265 (1964). 9. Larsen, J. W., and Tepley, L. B., J. Colloid Interface Sci. 53, 332 (1975). 10. De Lisi, R., Genova, C., and Livera, V. T.,J. Colloid Interface Sci. 95, 428 (1983). 11. Bury, R., and Treiner, C., J. Colloid Interface Sci. 103, 1 (1985). 12. Hall6n, D., Nilsson, S.-O., Rothschild, W., and Wads6, I., J. Chem. Thermodyn. 18, 429 (1986). 13. Rouw, A. C., and Somsen, G., J. Chem. Thermodyn. 13, 67 (1981). 14. Arnett, E. M., Kover, W. B., and Carter, J. V., J. Am. Chem. Soc. 91, 4028 (1969). 15. Hill, D. J. T., Ph.D. Thesis, University of Queensland, 1965.
Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
16. Krishnan, C. V., and Friedman, H. L., J. Phys. Chem. 73, 1572 (1969). 17. Abuin, E. B., and Lissi, E. A., J. Colloid Interface Sci. 95, 198 (1983). 18. Hojland, H., Ljosland, E., and Backlund, S., J. Colloid Interface Sci. 101, 467 (1984). 19. Lianos, P., and Zana, R., J. Colloid Interface Sci. 101, 587 (1984). 20. De Lisi, R., Genova, C., Testa, R., and Liveri, V. T., J. Solution. Chem. 13, 121 (1984). 21. Pilcher, G., Jones, M. N., Espada, L., and Skinner, H. A., J. Chem. Thermodyn. 1,381 (1969). 22. Mazer, N. A., and Olofsson, G., J. Phys. Chem. 86, 4584 (1982). 23. Carlfors, J., and Stilbs, P., J. Colloid Interface Sci. 104, 489 (1985). 24. Roux-Desgranges, G., Roux, A.-H., Grolier, J.-P. E., and Viallard, A.,J. Solution Chem. 11, 357 (1982). 25. Majer, V., Roux, A.-H., Roux-Desgranges, G., and Viallard, A., Can. J. Chem. 61, 139 (1983). 26. Tanford, C., "The Hydrophobic Effect," 2nd ed. Wiley, New York, 1980. 27. Gill, S. J., Dec, S. F., Olofsson, G., and Wads6, I., J. Phys. Chem. 89, 758 (1985). 28. Almgren, M., and Swarup, S., J. Colloid Interface Sci. 91, 256 (1983). 29. Gill, S. J., and Wads6, I., Proc. Natl. Acad. Sci. USA 73, 2955 (1976).
JOHNSON
AND G E R D
OLOFSSON
Division of Thermochemistry, ChemicalCenter, Universityof Lund, Lund, Sweden Received October 7, 1985; accepted February 2, 1986 The solubilization of 2-propanol and pentanol in sodium dodecyl sulfate (SDS) solutions has been studied by titration calorimetry over a wide range of alcohol concentrations in three different SDS solutions. Measurements were made at 25°C and in addition at 35°C for pentanol. The measured enthalpy changes are close approximations of the partial molar enthalpies of solution of the alcohols in surfactant solution. Reported values of the partition coefficient/~ derived from self-diffusionmeasurements using the Fourier transform NMR pulsed gradient spin-echo method were used to interpret the results. In the calorimetric measurements pentanol dissolves predominantly in the micellar phase, while a larger part of added 2-propanol remains in the aqueous phase. To separate the solubilization effects partial molar enthalpies of transfer A~ic/7of alcohol from pure liquid to micellar phase were derived assuming the distribution of alcohol between aqueous and micellar phase to be a simple partition equilibrium. The limiting value of A~iC/7at low alcohol content is -4.4 _+ 0.2 kJ mole-1 for pentanol and -6.5 _ 0.5 kJ mole-1 for 2-propanol. The derived values of A]~ic/1for pentanol show a pronounced variation with the composition of the micellar phase. The increase is almost linear up to a molar ratio 2:1 of alcohol:SDS in the micellar phase, indicating a smooth change in the milieu for the sohibilizate from discrete micelles with low alcohol content to alcohol-rich apolar domains. The heat capacity change for the transfer of liquid pentanol to the micellar phase is zero or close to it over the concentration range studied, which indicates that the alkyl chains are dissolved in a nonaqueous environment. © 1987Academic Press, Inc.
INTRODUCTION
tively. A l t h o u g h the solubilization p h e n o m e n o n has been extensively s t u d i e d b y different m e t h o d s , few techniques are available for t h e study o f the d i s t r i b u t i o n o f t h e solubilizate. R e c e n t l y a m e t h o d has b e e n d e v e l o p e d w h i c h p r o v i d e s direct i n f o r m a t i o n a b o u t the partit i o n i n g o f solubilizate b e t w e e n micelles a n d water (3-6). It is b a s e d o n m e a s u r e m e n t s o f i n d i v i d u a l self-diffusion coefficients in m u l t i c o m p o n e n t systems using a n N M R F o u r i e r t r a n s f o r m p u l s e d - g r a d i e n t spin-echo m e t h o d . T h e solubilization in s o d i u m d o d e c y l sulfate (SDS) solutions o f a n u m b e r o f solutes such as alcohols, ketones, a n d a r o m a t i c h y d r o c a r b o n s was studied b y this m e t h o d (4). T h e twostate p a r t i t i o n e q u i l i b r i u m m o d e l was f o u n d to give a satisfactory d e s c r i p t i o n o f the distrib u t i o n a n d values o f Kc were derived. T h e dist r i b u t i o n o f alcohols in SDS solutions has also
T h e solubility o f n o r m a l l y o n l y slightly soluble organic c o m p o u n d s in a q u e o u s s o l u t i o n is substantially increased w h e n surfactants are p r e s e n t in c o n c e n t r a t i o n s a b o v e their cmc. This solubilization p h e n o m e n o n is o f basic interest in the study o f m i c e l l a r a n d m i c r o e m u l s i o n solutions a n d it also has n u m e r o u s practical a p p l i c a t i o n s (1, 2). W i t h i n t h e phase s e p a r a t i o n m o d e l for micellization, solubilization is regarded as a partitioning o f the solute b e t w e e n the m i c e l l a r a n d a q u e o u s phases. A t low c o n c e n t r a t i o n s , a d i s t r i b u t i o n coefficient is a s s u m e d to describe the solubilization behavior: KC~ S S CmiffC aq [ 1] where c~i¢ a n d c~q denote solute c o n c e n t r a t i o n in the m i c e l l a r a n d a q u e o u s phases, respec56 0021-9797/87 $3.00 Copyright © 1987 by Academic Press, Inc. All fights of reproduction in any form reserved.
Journal of Colloid and Interface Science, Vol, 115, No. 1, January 1987
SOLUBILIZATION OF ALCOHOLS been derived from vapor pressure measurements (7). Values of K¢ for pentanol, hexanol, and heptanol determined by the two methods are in good agreement. We have chosen to study the solubilization of 2-propanol and pentanol in SDS solution by titration calorimetry to gain further information about the thermodynamics of the solubilization process. Solution calorimetric measurements have been made previously on the solubilization of alcohols in surfactant solutions, but they have been limited to only one alcohol concentration (8-11). Thus the effect of increasing solute concentration and accompanying change in composition of the micellar phase has not been studied. To get as detailed information as possible, we have measured differential enthalpies of solution of 2-propanol and pentanol from 0 up to 1 mole alcohol/kg water in SDS solutions of varying molality. The amount of alcohol added in each titration step was small and the measured enthalpy changes are good approximations of the partial molar enthalpies of solution. Therefore, they reflect the conditions of the solute in the solution at the composition of the experiment. Values of the partial molar enthalpies of transfer of liquid alcohol to the micellar phase were derived assuming the distribution of alcohol between the aqueous and micellar phases to be described by partition coefficients derived from the NMR self-diffusion measurements. EXPERIMENTAL
Materials. Specially pure (>99%) sodium dodecyl sulfate (SDS) (BDH, Poole, UK) was used without further treatment. Solutions of SDS were prepared by weight using reagentgrade water produced by a Milli-Q System (Millipore AB, Grteborg, Sweden). Pentanol (BDH, Poole, UK) was fractionally distilled and purity checked by analytical gas chromatography. The sample used was judged to be better than 99.8% pure, with a water content less than 0.05%. Analytical grade 2-propanol of purity better than 99.7% (Merck, Darmstadt, Germany) and analytical grade sodium
57
perchlorate (NaC104. H20) from Fluka AG (Buchs, Switzerland) was used without further treatment. Calorimetry. Titration calorimetric measurements were made using either the LKB Batch Microcalorimeter modified for titration (LKB Model 2107-350) or an LKB 8700 reaction-solution calorimeter with 25- or 95-cm 3 glass reaction vessels. In the batch microcalorimetric experiments, 5 cm 3 of SDS solution was put into the reaction cell (made of 18carat gold) and the same amount of pure water in the reference cell. In each titration step l0 mm 3 of alcohol was injected in the reaction cell and simultaneously l0 mm 3 of water in the reference cell. In the experiments using the reaction-solution calorimeter, the sample was titrated into the glass vessel by means of a thin steel tube (length 1 m, i.d. 0.15 mm) fastened by epoxy to a Hamilton 100-mm 3 gas-tight syringe (1710 LT). The capillary tube passes through a tube (i.d. 1 ram) in a slit at the inside of the stirrer holder and ends about 30 mm below the surface of the calorimeter liquid. The syringe is motor-driven, and the samples were injected at a rate of 0.10 mm 3 s-1. The sample size varied between 12 and 20 mm 3. To prevent the small drop of pentanol introduced into the calorimeter from creeping up to the surface before it dissolves, a small Teflon cap was mounted on the capillary tube 5 mm from the end. The droplet is trapped under the cap and is completely dissolved in the aqueous solution without evaporation loss, etc. RESULTS Titration experiments consisting of successive additions of 10 mm 3 of pure pentanol, PeOH, to 5 cm 3 SDS solution have been made at 25.0 and 35.5°C using the batch microcalorimeter. Three different SDS solutions having molalities 0.089, 0.182, and 0.370 mole kg-1 were used. Each titration series consisted of 60 or 80 consecutive injections which increased the alcohol content of the calorimeter solution from 0 to 0.9-1.0 mole (kg water) -x. The change in alcohol content was small in Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
58
JOHNSON AND OLOFSSON
each step, being about 0.02 mole kg -1. Therefore the measured enthalpy changes A/l(obs) calculated per mole of added alcohol are close approximations of the partial molar enthalpies of solution of pentanol in the SDS solutions. The results are summarized in Figs. 1a and b where 2x/l(obs) is plotted against the molality of pentanol at the two temperatures. Results of measurements in 0.089 m SDS have been deleted for pentanol content above 0.4 m as it is uncertain if equilibrium was reached under the conditions of the experiments. Visual inspection of samples containing 0.089 m SDS to which varying amounts of pentanol were added showed that the dissolution of added alcohol became slow and took up to 1 h after shaking to become clear when the alcohol content exceeded about 0.5 m. At lower alcohol content and in solutions with higher SDS content, the dissolution was rapid, giving
2 o E
b
t -35.5 °C
1
0.182
o
0mi 2
• ~=~A-===~
i"~ .A¢a ~ a ' *
t
0.4
,, ," "
l*
=•=A
'
0 p ~ ~ ' ~ m"
,eeee
tel•
~'1 m0
112
m/mol kg-1
f~ -1 -2
\ " II||~" lie•= •
°° ~0.370
ill•el e
_51
a
t =25.0 °C , , , = A , , ,`,= 0.2
-2
0.4
~=~A=ee
0.6
A,-=" ,~'== ee 0.182 ,-= e•=e , ' = •e •e" e
1.0 1.2 rn/mol kg-1
0.370
.=i"e=,. " ='=~ 0.089 -5 -(]
FIG. 1. Measured enthalpy changes per mole of added pentanol A/l(obs), as a function of final amount of pentanol in each titration step (a) at 25.0°C, (b) at 35.5°C. The molalities of the different SDS solutions are indicated in the figures. Journal of Colloid and lnterface Science, Vol. 115, No. 1, January 1987
-8 0,381
E -10
•eeeeeeeelleeeeeee*•eeeee
eee•ee
e•ee••
• • •
o
0.150
-12
o12
'
o15
'
11o
' m Imol kg -1
FIG. 2. Measured enthalpy changes per mole added 2-propanol, A//(obs), as a function of final amount of 2-propanolin each titration step at 25.0°C. The molalities of the different SDS solutions are indicated in the figure.
clear solutions over the whole range of pentanol concentrations. Measurements were made at low alcohol content using the reaction-solution calorimeter with a 25 cm 3 vessel. There is good agreement between results from the reaction-solution calorimetric and batch microcalorimetric measurements in overlapping concentration regions. Batch microcalorimetric results of step-wise additions of 2-propanol to SDS solutions of molalities 0.150, 0.250, and 0.381 mole kg -1 at 25.0°C are summarized in Fig. 2. Least-squares fitting of second-order polynomials in alcohol molality mA to observed enthalpy changes gives a satisfactory representation of the experimental results. Coefficients a, b, and c from fits of A/~(obs) = a + bmA + Cm2Ato results of the various titration series are shown in Table I. The standard errors of the fits were 0.05 kJ mole -1 or less, which is the same as the reproducibility of the calorimetric measurements. The empirical relations reproduce the calorimetric results of alcohol contents up to molalities mA(max) indicated in the last column of Table I. Experiments were also made on dissolving the alcohols in pure NaCIO4 solutions of varying molality. Each series of measurements consisted of three consecutive additions of 26.68 m m 3 alcohol to 95.5 cm 3 of water or NaC104 solution. The resulting solution enthalpies are shown in Table II. The measured enthalpy changes can be assumed to refer to infinitely dilute solution, since the alcohol molality did not exceed 0.001 mole kg -~. The
59
SOLUBILIZATION OF ALCOHOLS TABLE I Values of the Coefficients in the Relations A/t(obs)/kJ mole -1 = a + bmA + Cm2A
Solute Pentanol
Temperature (°C)
SDS (mole/kgH20)
a
25.0
0.0890 0.182 0.370 0.0890 0.182 0.370
-5.67 -5.04 -4.69 -4.38 -4.27 -4.26
0.150 0.250 0.381
- 11.52 -10.81 -9.99
35.5
2-Propanol
25.0
value of the solution enthalpy in pure water is in excellent agreement with values determined using two other calorimetric methods at our laboratory and with previously reported values (12).
Calculation of Transfer Enthalpies If the micelles are regarded as a separate phase, the alcohol A added in the calorimetric
c
mh(max) (mole/kgH20)
0.939 -0.439 -0.441 -8.64 -4.22 -1.76
0.38 0.70 0.76 0.30 0.70 0.80
b 7.09 6.31 4.93 11.56 9.50 6.07 0.242 0.180 0.0588
0.616 0.501 0.641
1.2 1.2 1.2
experiments can be assumed to be distributed between the micellar and aqueous phases: A(1) --~ aA(mic) + (1 - .)A(aq)
[2]
where c~ denotes the fraction of alcohol solubilized in the micelles. A distribution coefficient Kc [1] is assumed to describe the distribution of the solute between the two phases. Stilbs found this two-state model to give a sat-
TABLE II Enthalpies of Solution of Pentanol and 2-Propanol in Pure Water and in NaC104 Solutions at Infinite Dilution of Alcohol m(NaCIO4) Solute Pentanol 25.0°C
35.0°C
2-Propanol 25.0°C
0
-7.93 + 0.009 a -7.99 b -7.93 b -4.76 + 0.14 -4.82 b
-13.16 + 0.06 -13.10 c -13.09 a -12.98 e -12.88 f
0.090
0.184
0.377
-7.53 + 0.15
-7.14 ± 0.15
- 6 . 1 4 + 0.15
-4.46 + 0.10
-4.23 + 0.10
-3.33 ___0.10
-12.68 + 0.10
-12.44 + 0.10
-11.67 + 0.10
a Error limits are expressed as twice the calculated standard deviation of the mean. b Ref. (12). *Ref. (13). dRef. (14). *Ref. (15). fRef. (16). Journal of Colloid and Interface Science, Vol. 115,No. 1, January 1987
60
JOHNSON AND OLOFSSON
isfactory description of the partitioning of various solutes in SDS solutions from N M R self-diffusion measurements (4). Four different solutes including pentanol were studied in solutions containing between 0.04 and 0.25 mole liter -1 of SDS and the concentration of two solutes, benzene and benzyl alcohol, was varied between 0.03 and 0.24 mole liter -1, while the SDS concentration was kept at 0.24 mole liter-1. Stilbs has also studied the solubilization of n-alcohols at higher solubilizate and SDS concentrations (5). For n-butanol there is a good agreement between the value of Kc derived at low concentrations and values derived for solutions containing up to 0.46 mole liter-~ of SDS and about 1 mole liter -1 of butanol. The self-diffusion behavior of the different species indicates that up to these concentrations the system is best described as a micellar solution (5). The value ofK~ = 49 for pentanol derived for a solution containing 0.34 mole liter -~ ofpentanol and 0.25 mole liter -1 of SDS is in satisfactory agreement with the value of 55 _+ 3 found at low alcohol concentration (4). At higher alcohol and SDS concentrations the self-diffusion results for SDS indicate a breakdown of the micelles. However, the self-diffusion of the solubilizate remains constantly slow irrespective of solution composition, which can be seen as indicating relatively constant alcohol partition equilibria between polar and apolar domains (5). Thus it appears that the simple two-state equilibrium model can describe the solubilization of butanol and pentanol in solutions containing up to at least 0.46 and 0.25 mole liter-~ of SDS, respectively, as long as the molar ratio of alcohol:SDS in the mixed micelles does not much exceed 1:1.1 At higher concentrations the alcohol is still solubilized in apolar domains with a distribution resembling that at lower concentrations. Assuming the two-state partition equilibrium model to be applicable, the observed en-
thalpy change A/l(obs) in our experiments can be expressed as
A/~(obs) = aA[nic/-I-t- (1 - a)A~q/~.
[31
The partial molar enthalpy of solubilization of A into the micellar phase from the liquid state is denoted ~{nic/~ and the partial molar enthalpy of solution of liquid A in the aqueous phase is denoted A~q/1. The degree of solubilization a was calculated using values for the partition coefficient Ko determined by Stilbs (4) from N M R self-diffusion measurements. The reported value for pentanol is 5 5 + 3 and for 2-propanol J.J-l.stZ t:+2.5. Values of a were calculated according to a/(1 - a) = Kc. Vmio/Vaq using the value of 250 cm 3 (mole SDS) -l for Vmi~ (4, 20) and 18.0 cm 3 (mole H20) -1 for Vaq. This means that Vmio and V,q are proportional to the amount of SDS and water in the solutions and may differ from the actual volumes, particularly in the more concentrated solutions. The concentration of free SDS is assumed to be equal to the cmc, 8 × 10-3 mole kg -1. This is an overestimate of the monomer concentration, but since it is not known how it varies with alcohol and SDS concentration, this value has been used throughout. The error introduced in Vmi~ is small and will not significantly influence the results. Changes in monomer concentration with alcohol content could give an enthalpy contribution to the measured A/t(obs). This effect is considered to be less than - 0 . 5 ld mole -1. The estimate was made assuming the monomer concentration to decrease to the same extent as has been observed close to the cmc (7) and using an enthalpy of micelle formation o f - 5 . 0 kJ/mole. 2 The self-diffusion measurements were made in D20 but a study of the solubilization of nalcohols in sodium decanoate micelles in both D20 and H 2 0 indicated that partition coefficients are the same in the two solvents (23). The effect of temperature on K~ was found to
I The abrupt changes in solubility and K~ observed 2 The enthalpy of micelle formation of SDS in pure for hexanol and heptanol at certain SDS concentrations (17-19) has not been observedfor butanol and pentanol water is closeto zero at 25°C and -5.1 kJ mole-1 at 35°C (21, 22). (5, 18). Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
61
SOLUBILIZATION OF ALCOHOLS
be small (4) and for pentanol we have used the same value at 25 and 35°C. The solution enthalpy in water, A~q/~, will vary with the composition of the aqueous phase. In evaluating A~i~/laccording to relation [3], we have assumed that the ionic medium effect on A~q/~ can be estimated from the enthalpies of solution of the alcohol in NaC104 solutions having the same molality as the SDS solutions. Further, it is assumed that the variation of A~q/-I with alcohol concentration in the aqueous phase is the same as in pure water. Estimates of the partial molar enthalpies of solution in water were based on results from dilution experiments on C1-C4 alcohols (12). Values of the pairwise and triplet interaction coefficients for 2-propanol were assumed to be the same as for 1-propanol, 557 J mole -~ (mole kg-1) and 180 J mole -1 (mole kg-l) -2, respectively. Extrapolation of the values for the coefficients for the C1-C4 alcohols gave for pentanol 1930 J mole -1 (mole kg-1)-1 for the pairwise and 740 J mole -1 (mole kg-1)-2 for the triplet interaction coefficient. Values of the solubilization enthalpy A~i~/lfor pentanol derived in this manner are summarized in Fig. 3, where A~i¢/7 is plotted against the molar ratio pentanol/SDS in the micellar phase, np~oH/nSDS. As can be seen, A~nic/tis closely the same in the three systems at the same composition of the micellar phase. However, A~i¢/7 is not constant but becomes increasingly more endothermic as the alcohol content of the micellar phase increases. The effect of the 10°C change of temperature in A~i~/t is negligible. Values of A~i¢/1 for the solubilization of 2-propanol are summarized in Fig. 4. The degree of solubilization of this alcohol is low and the value of K~ has large error limits. Therefore, the derived values of A~'CH are uncertain. The difference between A~°/t for the different SDS solutions seen in the figure is not significant, as a small change of K¢ will shift the relative position of the curves and can also make them coincide. We have defined the solubilization enthalpy as the partial molar enthalpy of solution of liquid alcohol into the micellar phase A~nic/~.
b t=35.5°C 0.18: ~
P
~-ii/~ a t=25.0°C "7,
0.182 1
.o
~
J
~
<1
3 . 89
J J__ 4 5 rl PeOH/FISDS
-2
FIG. 3. Partial molar enthalpy of solubilization of pentanol into the micellar phase from the liquid state, A[~ic/1, as a function of calculated molar ratio pentanol/ SDS in the micellar phase (a) at 25.0°C, (b) at 35.5°C. The molalities of the different SDS solutions are indicated in the figures.
It is the difference between the partial molar enthalpy content of the alcohol in the micellar phase and the enthalpy content of liquid alcohol: A~ic/~ =/tmic -- HI. Only/~mic will vary with the composition of the system. The solubilization enthalpy can also be defined as the enthalpy of transfer of alcohol from aqueous to micellar phase: Aaq micH- =/tmi¢ --/l~a. Values of A~ ~/lcan be calculated by subtracting from A~i°/1 the appropriate value for the solution enthalpy in aqueous solution, A~q~. The partial molar content of alcohol will vary both in the micellar and aqueous phase with composition. Values of enthalpy changes for the transfer ofpentanol from aqueous solution to the SDS micellar phase, A~¢H have been reported previously (8, 10, 11). These values have been derived from calorimetric measurements at only one alcohol concentration. The value reported by Lisi et aL (10) of 3.44 + 0.06 kJ Journal of Colloid and Interface Science, Vol. 115, No. 1, January 1987
62
JOHNSON AND OLOFSSON o
0.381 ~
E<~
~..~177JJ.~. 0.250.
-8 t
\0.150
0'5
1'o n iPrOH / rlSDS
FIG. 4. Partial molar enthalpy of solubilization of 2-propanol into the micellarphase fromthe liquid state, A[nie/~, as a functionofcalculatedmolarratio2-propanol/ SDSin the micellarphaseat 25.0°C. The molalitiesofthe differentSDS solutionsare indicatedin the figure.
mole -1 is in good agreement with the value we calculate by subtracting the enthalpy of solution ofpentanol in water (Table II) from the limiting value of A ~ n i c / t , s e e Fig. 3. Other reported values of A~qi~/4do not agree so well, the discrepancy stemming mainly from the use of a different and, as we believe, incorrect value for the enthalpy of solution of pentanol in water. Measurements have been made at 25°C of densities and heat capacities of the ternary systems water-SDS-butanol (24) and waterSDS-2-propanol (25) over the complete alcohol mole fraction (solubility) range. Apparent and partial molar volumes and heat capacities of the alcohols were derived. Assuming that the partition equilibria at low alcohol content can be described by the partition coefficients determined by Stilbs (4), partial molar heat capacities of the solubilized alcohols can be estimated. Values derived from the reported Cp,~ at low alcohol and SDS concentrations are close to the values for the heat capacities of the liquid alcohols, indicating that the heat capacity change for the transfer of liquid alcohol to micellar phase is close to zero for both 2-propanol and butanol. DISCUSSION The observed enthalpies of solution of alcohols in SDS solutions are less exothermic than in pure water or NaC104 solutions due to partial solubilization of the alcohols in the micellar phase. Pentanol and 2-propanol differ Journal of Colloid and Interface Science, Vo]. 115, No. 1, January 1987
significantly in their solubilization behavior, with pentanol showing a strong preference for the micellar phase as indicated by the value 55 for the partition coefficient, while 2-propanol is solubilized to a much less extent, Kc being about 5.5. This means that in the calorimetric measurements on pentanol in the 0.18 and 0.37 mole liter-1 SDS solutions, a major part ofpentanol added in each experiment enters the micellar phase and the observed enthalpy change is dominated by the enthalpy contribution from the solubilization of liquid pentanol into the micelles, A~ic/1. As can be seen from Fig. 1, there is a smooth change in A/4(obs) as the pentanol concentration increases, indicating a smooth change in A~i¢/-I with composition up to high alcohol content, where micellar breakdown occurs (5). The incremental solution enthalpies become steadily more endothermic, indicating that the alcohol continues to dissolve in a nonaqueous environment and that the milieu for the solubilizate changes smoothly from discrete micelles with low alcohol content to alcohol-rich apolar domains. The change of A/4(obs) with temperature is small, indicating that the heat capacity change for the transfer of liquid pentanol to the micellar phase is small. In the experiments on 2-propanol, a larger part of each added portion will remain in the aqueous phase, but also for 2-propanol, A/l(obs) is significantly less exothermic than the solution enthalpy in pure water. To get a clearer picture of the solubilization process, a separation of the enthalpy contributions from solubilization and hydration has been attempted by the derivation of differential solubilization enthalpies A{nic/1as described previously. The calculated values are tentative due to the approximations made, but at least at low alcohol content, where the two-state partition equilibrium model is applicable, the derived values are considered to give a fair representation of the partial molar enthalpies of transfer of liquid alcohol to micellar phase. This assumption is supported by the simplified features of the plots of A~iCH against composition of the micellar phase, Figs. 3 and 4.
SOLUBILIZATION OF ALCOHOLS
63
They show, for instance, that the derived val- for the incorporation of liquid pentanol into ues of A~nic/~ do not vary with SDS concen- the miceUar phase indicates that the solubitration for molar ratios alcohol/SDS of up to lized alcohol is not exposed to water, but that the alkyl chain is in a nonaqueous environat least 1:1 in the micellar phase. As seen from Fig. 3, the partial molar en- ment. This conclusion may also apply to buthalpy of transfer, A~n'cH, of liquid pentanol tanol and 2-propanol, since the heat capacity to micellar phase varies with the composition results from the measurements of the ternary of the micellar phase. At low alcohol content, solutions (24, 25) are fully compatible with A~ic/1 is exothermic, being -4.4 ___ 0.2 kJ very small or zero heat capacity changes for mole-1 in the limit of np~oH/nSDSequal to zero. the transfer of liquid alcohol to miceUar phase The corresponding value for 2-propanol is at low alcohol content. -6.5 + 0.5 kJ mole -1. At higher concentraIt is generally assumed that an additive like tions the values of A~iCH for 2-propanol are pentanol is solubilized in the micelles with the too uncertain to give any information about polar group enclosed at the surface among the their variation with the composition of the surfactant head groups and with the alkyl micellar phase. We refrain from calling the chain immersed in the hydrocarbon case. The limiting values "standard state" values. In our observed zero heat capacity change for the opinion, the concept of standard state is not transfer of liquid pentanol into SDS micelles particularly useful in this situation, as it is not supports this model. The incorporation ofsolpossible to derive standard state values from ubilizate in the micelles will decrease the surexperimental results without using a particular face charge density, the decrease for SDS mimodel for the solubilization process. The re- celles being at least to a first approximation suiting values will therefore depend on the linear with the mole fraction of additive in the model used. Besides, even within the frame- micellar phase (28). As can be seen from Fig. work of a particular model, in the present case 3, A[nic/t increases almost linearly with penthe partition equilibrium model, it is not ev- tanol content in the micellar phase from -4.4 ident how to define meaningful standard kJ mole -1 at very low molar ratio pentanol/ state(s). SDS to zero at a molar ratio of about 2. A striking feature of the results for pentanol In the partition process A(aq) ~ A(mic), is that Arai~/ldoes not vary with temperature, the transfer ofpentanol and 2-propanol from at least not for npeoH/nsosless than 1. If Figs. the aqueous to the micellar phase is endo3a and b are superimposed, the curves coincide thermic at all concentrations studied. The to within +0.2 kJ mole -1 at low molar ratios, positive transfer enthalpies indicate that the which is remarkable considering the simpli- transfer entropy is positive and the driving fying assumptions made. At higher molar ra- factor for the solubilization process. Positive tios, the curves for the two more dilute SDS entropy contributions could stem from the solutions diverge somewhat, but they still lie dehydration of the alkyl chain of the additive within +0.5 kJ mole -1. Thus the heat capacity and a decrease in ion binding to the micelle change for the transfer of liquid pentanol to as the surface charge density decreases. For micellar phase is zero or close to it. This can the solubilization of a series of additives like be compared with the heat capacity change of 1-alkanols, the dehydration entropy can be 300 J K -l mole -1 for the transfer of liquid expected to be proportional to the chain length pentanol to aqueous solution, see Table II. The (29) while the change in surface charge density large positive value of the heat capacity of so- will not vary. The increase in partition coeflution in water stems from the strong temper- ficient with chain length observed for the solature dependence of the interaction between ubilization of 1-alkanols in SDS solution (4, the alkyl groups and the surrounding water 7) is probably governed by the increase of the layer (26, 27). The zero heat capacity change dehydration entropy with chain length. Journal of Colloid and Interface Science, Vol. 115, No. l, January 1987
64
JOHNSON AND OLOFSSON REFERENCES
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