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The solubility of carbon dioxide and nitrous oxide in aqueous solutions of cetyltrimethylammonium bromide, sodium dodecyl sulfate, sodium 1-heptanesulfonate, and sodium perfluorooctanoate
The Solubility of Carbon Dioxide and Nitrous Oxide in Aqueous Solutions of Cetyltrimethylammonium Bromide, Sodium Dodecyl Sulfate, Sodium 1-Heptanesulfonate, and Sodium Perfluorooctanoate D. W . O W N B Y ,
W. PRAPAITRAKUL,
AND A. D. K I N G , JR.
Department of Chemistry, School of Chemical Sciences, University of Georgia, Athens, Georgia 30602 Received July 20, 1987; accepted November 24, 1987 Measurements have been made to determine the solubilities of carbon dioxide and nitrous oxide in aqueous solutions of cetyltrimethylammonium bromide (CTAB), sodium dodecyl sulfate (SDS), sodium 1-heptanesulfonate (SHSo), and sodium perfluorooctanoate (SPFO) at elevated pressures. The solubilities of each gas, measured at 26°C for solutions of CTAB and at 25°C for the other surfactants, follow Henry's law at all surfactant concentrations. Two distinctly different types of behavior are exhibited by SHSo. At concentrations below the critical micelle concentration (CMC), the addition of SHSo acts to suppress the solubility of CO2, while at higher concentrations, the opposite behavior is observed in that addition of SHSo results in an increase in CO2 solubility. The CMC values for the other surfactants are all much lower than that for SHSo. As a result, the solubility data for N20 and CO2 in these surfactants are all taken at concentrations in excess of the CMC and in every case, the addition of surfactant is found to enhance gas solubility. The increase in gas solubility observed above the CMC for each of the four surfactants is found to be a linear function of surfactant concentration as expected for micellar solubilization. The intramicellar solubilities of COz and N20 calculated from these latter data do not correlate well with solubilitiesof other nonpolar gasesbut rather are displaced toward higher concentrations, thus suggesting that the hydrophilic character of these two gases leads to enhanced solubilization. © 1988 Academic Press, Inc.
INTRODUCTION T h e gases c a r b o n dioxide (CO2) a n d nitrous oxide ( N 2 0 ) have very s i m i l a r physical p r o p erties. This arises f r o m the fact t h a t each gas is c o m p o s e d o f l i n e a r t r i a t o m i c m o l e c u l e s which are isoelectronic a n d have virtually identical d i m e n s i o n s a n d m o l e c u l a r weights. Like m o s t c o m m o n gases, CO2 a n d N 2 0 are basically h y d r o p h o b i c in n a t u r e in t h a t t h e y are c o n s i d e r a b l y m o r e soluble in oil t h a n in water. However, a c o m p a r i s o n o f the a q u e o u s solubilities listed in T a b l e II shows t h a t CO2 a n d N 2 0 nevertheless exhibit a n o t i c e a b l e affinity for water as e v i d e n c e d b y the fact t h a t a q u e o u s solubilities o f these two gases are an o r d e r o f m a g n i t u d e greater t h a n those o f the o t h e r n o n p o l a r gases listed. This o b s e r v a t i o n is o f course n o t new b u t r a t h e r was cited b y H i l d e b r a n d in 1950 as evidence that, in ad-
0021-9797/88 $3.00 Copyright © 1988 by Academic Press, Inc. All rights of reproduction in any form rcscrve,d.
d i t i o n to the usual physical forces, m o l e c u l e s o f dissolved C O 2 a n d N 2 0 i n t e r a c t with w a t e r t h r o u g h highly specific c h e m i c a l interactions, p r e s u m a b l y o f a Lewis a c i d - b a s e n a t u r e (1). T h e r e is a b u n d a n t evidence f r o m o t h e r sources to s u p p o r t this claim. F o r e x a m p l e , n o t o n l y d o solubility m e a s u r e m e n t s o f a reciprocal n a t u r e show t h a t w a t e r is a b n o r m a l l y soluble in c o m p r e s s e d gaseous CO2 a n d N 2 0 ( 2 ) b u t surface t e n s i o n d e t e r m i n a t i o n s at high pressures also provide clear evidence that these two gases are selectively a d s o r b e d at the g a s water interface (3, 4). M o r e recently, r a d i o frequency a n d m i c r o w a v e spectra o f m o l e c u l a r b e a m s c o n t a i n i n g CO2 a n d water v a p o r have p r o v i d e d direct evidence t h a t CO2 a n d w a t e r f o r m relatively rigid p l a n a r 1:1 c o m p l e x e s having C2 v s y m m e t r y with the hydrogen a t o m s d i r e c t e d a w a y f r o m the CO2 m o l e c u l e ( 5 ) .
526 Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
S U R F A C T A N T SOLUTIONS
This suggests that, as originally proposed by Hildebrand and Scott (1), the primary mode of bonding is Lewis acid-base in nature with the lone pair electrons of water interacting with unoccupied molecular orbitals of CO2. With this in mind, a series of measurements has been made to determine the solubilities of these same two gases, CO2 and N20, in micellar solutions of a wide variety of ionic surfactants. The surfactants chosen for this study consist of a cationic surfactant, CTAB, and three anionic surfactants, SDS, SHSo, and SPFO, that have tail groups which differ markedly with respect to chemical nature and size. A wide selection such as this is of obvious value in that it allows one to compare the micellar solubilities of these two gases with similar data obtained previously with other nonpolar gases and thus ascertain the extent to which this affinity toward water exhibited by CO2 and N20 affects the degree to which these gases are solubilized within micelles having markedly different surface and internal properties. 1 EXPERIMENTAL
The apparatus and procedures used to determine gas solubilities have been described in detail previously (6i, 6n). The sources and quoted purities of the surfactants used in these experiments are cetyltflmethylammonium bromide (98%), purchased from Lancaster Synthesis Ltd.; sodium dodecyl sulfate (99.0%) BDH Prod. No. 44244 (Lot No. 908811C), purchased from Gallard-Schlesinger Chemical Mfg. Corp.; and sodium 1-heptanesulfonate (99%), purchased from Lancaster Synthesis Ltd. The sodium periluorooctanoate was prepared in situ by titrating perfluorooctanoic acid (99% n-), SCM Specialty Chemicals (Lot No. 10439), with Fisher Certified ACS grade sodium hydroxide to pH 9-12. The CO2 and N20 used in these studies were commercial grade (quoted purity of 99.5% min) and CP 1 Reference (6) constitutes a reasonably complete bibliography on the subject of gas solubilization.
527
grade (quoted purity of 99.0%), respectively, purchased from Matheson Gas Products. RESULTS A N D DISCUSSION
The gas solubilities determined for CO2 and N20 in aqueous solutions of the various surfactants are listed in Table I. These data are shown plotted as a function of surfactant concentration in Fig. 1. Here it is seen that the solubility data fall into two classes. The gas solubilities obtained with SDS, CTAB, and SPFO, all of which have low CMC values, 2 are seen to increase linearly with surfactant concentration, thus generating a series of straight lines having positive slopes whose origins coincide with the corresponding gas solubilities measured in pure water. This is the typical pattern of behavior one encounters with micellar solutions of surfactants in which the micelles responsible for solubilizing the gas molecules maintain a constant sorptive capacity which is independent of surfactant concentration. The slopes of the straight lines defined by these solubility data correspond to the mole ratios ofsolubilized gas molecules to surfactant ions in micellar form. Therefore, in view of the high dilutions involved, the numerical magnitudes of the slopes of the gas solubility data for these surfactants can be taken to equal the 1-atm mole fraction solubilities of each gas within the individual micelles, X TM, i = CO2, N20. The micellar gas solubilities derived from straight lines fitted to the experimental solubility data for SDS, CTAB, and SPFO using the least-squares criteflon are shown listed in columns 4-6 of Table II. The CMC for SHSo, on the other hand, is quite large in comparison with the other surfactants, having a value of 0.28 M(0.3 m) at 25°C (6t). As a consequence, the CO2 solubilities obtained in these experiments for aqueous solutions of this surfactant span a wide range of concentrations both above and 2 Literature values and associated references for these surfactants at 25°C are SDS, 8.1 × 10 -3 M ( 7 ) ; CTAB, 9.2 × 1 0 - 4 M ( 7 ) ; SPFO, 3.2 × 10-2 M ( 8 ) . Journal of Colloid and Interface Science, Vol. 125,No. 2, October 1988
528
OWNBY, PRAPAITRAKUL, AND KING TABLE I Solubility of CO2 and N20 in Surfactant Solutions Expressed as Moles Gas per Atmosphere in 1000 g H20 Gas solubility (m) × 102a Surfactant
Temperature (°C)
Sodium heptylsulfonate
25
Sodium dodecylsulfate
25
Cetyltrimethylammoniumbromide
26
Sodium perfluorooctanoate
25
Surfactant concentration (m)
CO:
N20
0 0.2 0.3 0.4 0.5 0.6 0.8 1.0 1.2 0.2 0.4 0.6 0
3.32, 3.392b 3.17 3.12 3.05 3.05 3.07 3.16 3.26 3.33 3.54 3.69 3.91 3.28
2.433b
0.1
3.46
0.2 0.3 0.4 0.2 0.4 0.6
3.60 3.73 3.88 3.61 3.90 4.16
2.67 2.94 3.15
a Averageerror: +0.02 × 10-2m. b Recommended values from Wilhelm, E., Battino, R., and Wilcock, R. J., Chem. Rev. 77, 219 (1977). below the CMC. Figure 1 shows that C O 2 solubility in SHSo solutions exhibits a distinctly different behavior in each of these regions. Below the CMC, it is seen that SHSo acts to suppress the solubility of CO2. One concludes that SHSo ions in monomeric form "salt out" dissolved CO2. Salting out effects are c o m m o n l y reported as Setchenow coefficients, k, defined as k = ( 1 / m ) l o g ( S ° / S ) , where S Oand S represent the solubilities of the substance of interest in pure water ( S °) and in a solution of dissolved electrolyte (S), with m denoting the molal concentration of the electrolyte. The CO2 solubility data for pure water and 0.2 and 0.3 m solutions of SHSo, when fitted to the expression above, yield a value of k = 0.089. It is interesting to compare the salting out effect of SHSo m o n o m e r s toward CO2 to that of other c o m m o n sodium salts. M a r k h a m and Kobe (9) have measured the solubility of CO2 in a number of solutions of strong electrolytes. The data they report at low salt concentrations Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
(0 < m < 0.2) yield values o f k = 0.108 and k = 0.087 for NaCI and NaNO3, respectively, at 25 °C. Thus one concludes that m o n o m e r i c SHSo and sodium nitrate are about equally effective in salting out CO2. At concentrations in excess of the CMC, this negative trend reverses and the CO2 solubility increases with surfactant concentration as expected when micelles contribute to solubilization. The slope of the linear portion of the SHSo solubility data ranging between 0.6 and 1.2 m is m u c h smaller than the corresponding slopes of the other systems, indicating that SHSo micelles have a smaller sorptive capacity toward CO2 in comparison with the other surfactants. The mole fraction solubility of CO2 in micellar SHSo corresponding to this slope is listed in column 3 of Table II along with those for other gases. Previous investigators (10) have shown that 1-atm gas solubilities measured for a given nonpolar solvent, when plotted logarithmi-
SURFACTANT SOLUTIONS 4.5 CO 2 0/ 4.0
E
0 SPFO A CTAB [] SDS
/
/ 3.5
/
@ SPFO
0 SHSo
~
"7
0 ~"
-~o.<~
> o~J~-
N20
/~j
o~
3.02.5I / / ' / ~ / / / ° / @
2.0 f l 0
I I I I I [ I I I [ I I I I I 0,2
0.4
0.6
0-8
1.0
1-2
1-4
1-6
SURFACTANT CONCENTRATION (moles kg-1)
FIG. 1. Moles of gas absorbed per atmosphere in 1000 g of H20 with added surfactant shown as a function of surfactant concentration at 25°C (data for CTAB taken at 26°C). CsHIs, n-octane; CTAB, hexadecyltrimethylammonium bromide; SDS, sodium dodecylsulfate; SHSo, sodium heptylsulfonate; C7F16, perfluoroheptane; and SPFO, sodium perfluorooctanoate.
529
the cases of the nonpolar gases, T c and AEbv are proportional to one another. 3 There is an added advantage to using T c as a scaling parameter in situations involving gases having large quadrupole m o m e n t s such as CO2 and N20. This stems from the fact that a statistical treatment of the thermodynamic properties of such systems by Rowlinson (13, 14) has shown that the effects of angle-dependent interactions can be factored out from those arising from spherically symmetric dispersion forces in a m a n n e r consistent with the principle of corresponding states, thus allowing one to define an effective critical temperature, T0c, which characterizes the contributions that dispersion forces alone make to the overall intermolecular interactions between molecules of such gases. This is important because in cases where quadrupolar gases such as CO2 or N 2 0 interact with nonpolar substances, dispersion forces constitute the major m o d e of interaction so that T c rather than T c is the appropriate scaling factor to use as a measure of the extent of interaction. According to the treatment by Rowlinson, T0c is related to the experimental critical temperature, T c, by a single parameter, 6c, according to the equation T c = T0C(1 + 2c3c).
cally, correlate well with the internal energy of vaporization of the individual gases at their respective boiling points, AE v, thus suggesting that the Gibbs free energy of solution is a linear function of AE~. Similar trends are found for chemically inert gases dissolved in moderately polar solvents. Unfortunately, the triple point of CO2 lies above 1 atm, thus precluding one from using any such scheme based on AE v to compare solubility data for CO2 with that of other gases. There are, however, a variety of other properties (boiling points, Lennard-Jones force constants, critical temperatures, etc.) that m a y be used as scaling parameters. O f these, gas critical temperature, T c, is undoubtedly the most suitable for our purposes here because not only are T c values known to a high degree of precision for the gases of interest here, including CO2, but in
[1]
The average values for 6c derived from a wide variety of thermodynamic data are found by Rowlinson (14) to be 6c = 0.10 for CO2 and 6c = 0.06 for N20. 4 Values of T~o calculated using Eq. [1] and these parameters are listed 3 AEv values taken from Ref. (10) for 11 light gases (He, Ne, Ar, Kr, Xe, H2, N2, 02, CO, C H 4 , and C2H6) when plotted against Tc (Refs. (11, 12)) yield a well-defined straight line (correlation coefficient= 0.99663 ) having essentially zero intercept (-0.1 kcal mole-l) and a slopewhich, when dividedby the gas constant, corresponds to A E V / R T c = 5.1. 4 The theoretical treatment by Rowlinson presupposes that angular-dependent interactions between axiallysymmetric moleculesact as a perturbation on the centrosymmetric dispersion force term of the pair potential and thus varieswith intermolecularseparation,r, as r -6. This closely approximates the functional form of the potential arising from quadrupole-quadrupole interactions which variesas r -3 before averaging,and it has been shown (15 ) that the Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
530
OWNBY, PRAPAITRAKUL, AND KING TABLE II Gas Solubilities at 25°C and 1 arm in Surfactants and Bulk Solvents (Mole Fraction × 104)
Gas
Tc (K)
SHSo
SDS
CTAB
Ar O2 CH4 CF4 CO2 N20 C2H6 Call8
150.75 154.58 190.55 227.6 304.15 (253) 309.57 (276) 305.43 369.82
41 41 101 -42 a -451 1021
112 135 102 105 192 295 --93 a 147 a --1 1 0 3 1555 3304 4805
SPFO
CHaOH
171 4.156 181 4.117 191 8.676 381 . 139 a 56.3 s 120 a 65 b 611 40.56 1481 130 b
CH3(CH2)2OH CH3(CH2)7OH
7.826 6.807 16.16 . . 68.07 84 b 90.06 318 b
12.739 11.329 26.879
n-CaHjs
23.369 21.779 50-269
. 93.019 134 b 1736 607 b
119.89 175 b 3241° 1200 I1
C7FI6
H20
5612 6012 7312 13612 208.813 278 b 21712 61212
0.25214 0.23014 0.251 TM -6.11114 4.383 TM 0.334 TM 0.270 TM
This work. Estimated error for micellar solubilities of CO2 and N20:_+4 × 10-4. b This work. Estimated errors for solubilities in alcohols, n-octane and perfluoroheptane: N20, _+2 × 10-4; C3H8, +_4 × 10-4. l Ref. (6t). 2 Ref. (6i). 3 Ref. (6n). 4 Ref. (6p). 5 Ref. (6q); T = 26°C. 6 Ref. (17). 7 From data of Tokunga, J., J. Chem. Eng. Data 20, 41 (1975). 8 From data of Kunerth, W., Phys. Rev. 19, 512 (1922). 9 Wilcock, R. J., Battino, R., Danforth, W. F., and Wilhelm, E., J. Chem. Thermodyn. 10, 817 (1978). 1° Hayduk, W., and Cheng, S. C., Can. J. Chem. Eng. 48, 93 (1970). 11Thomsen, E. S., and Gjaldbaek, J. Chr., Acta Chem. Scand. 17, 127, 134 (1963). 12 Ref. (6v). 13 Kobatake, Y., and Hildebrand, J. H., J. Phys. Chem. 65, 331 (1961). 14Wilhelm, E., Battino, R., and Wilcock, R. J., Chem. Rev. 77, 219 (1977). a
in parentheses in column 2 of Table II for these two gases. One-atmosphere solubilities of CO2, N20, and other light gases dissolved in the bulk solvents n-octane, perfluoroheptane, water, and three representative alcohols are listed in Table II. These data, when plotted logarithmically as a function of gas critical temperature as seen in Fig. 2, reveal several features of interest. First, it is seen that when plotted in this manner, the solubility data for nonpolar gases in a given solvent correlate well with gas critical temperature as expected. Second, since elec-
values found by Rowlinson (14) of 6c = 0.01 and 0.06 for CO2 and 1'420, respectively, can be accounted for almost exactly by assuming that quadrupole--quadrupole interactions are solely responsible for the angle dependent forces between these molecules. Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
trostatic interactions of quadrupolar origin are not expected to affect the solubilities of CO2 or N20 in n-octane, T~0rather than T c should be the appropriate scaling parameter for these two gases. This is borne out by the fact that, when plotted as a function of T~0, the solubility data for these two gases in n-octane indeed coincide well with the line defined by the other nonpolar gases. Third, one sees that the solubilities of nonpolar gases in the various alcohols describe a series of nearly parallel lines which exhibit a progressive displacement toward decreased solubility as the carbon number of each alcohol decreases. The displacement of these lines toward lower solubilities simply reflects the behavior expected of chemically unreactive nonpolar gases in that the solubilities of such gases invariably decrease as the internal pressure of the solvent in-
SURFACTANT SOLUTIONS
-1.0 - 2.O
-
o/
-4.0
%
CIOH
-5-0
/
•
-9.0 -lO.O --%--0 -11.0 --12.0 IO0
)1
-
I
-
-
-
o
I I
Ar 0 2 C H 4
I
CO2 N20 C2H6
o-
H20
I C3H 8
I
I
I
200
300
400
TC(°K] FIG. 2. One-atmosphere gas solubilities of a variety of gases in n-octane (C8), 1-0ctanol (CsOH), l-propanol (C3OH), methanol (CIOH), and water plotted logarithmically as a function of gas critical temperature using T0c for CO2 and N20.
creases.5 However, in contrast to their behavior in n-octane, the solubilities of CO2 and N20 in the various alcohols and water are seen, in every case, to lie above the line defined by the solubilities of the other gases. As noted earlier, the enhanced solubilities of CO2 and N20 relative to the other gases in water have been cited by Hildebrand as evidence for chemical association between these gases and water (1). 5 Hildebrand solubility parameters in conjunction with regular solution theory provide a convenient measure of the role played by solvent internal pressure in determining gas solubility (Ref. (1, Chap. 15)). Regular solution theory predicts that gas solubilities, expressed logarithmically, will decrease in a regular fashion as the solubility parameter of the solvent increases. This trend has been shown to hold for alcohols as well as nonhydrogen bonding solvents (16, 17). The solubility parameters of the solvents used in this study, expressed as (cal em-3)1/2are n-CsH1a, 5.7; CH3(CH2)7OH, 9.3; CH3(CH2)2OH, 11.9; CH3OH, 14.4; H20, 23.4.
531
Therefore, although effects of dipole-quadrupole interactions cannot be ruled out, it is reasonable to assume that the excess solubility exhibited by CO2 and N20 in the various alcohols is likewise a manifestation of a chemical association resulting from highly specific interactions between molecules of CO2 and N20 and the hydroxyl group common to these alcohols. 6 Figures 3a and 3b show the micellar gas solubilities contained in columns 3-6 of Table II plotted logarithmically as a function of gas critical temperature. The micellar solubilities of the nonpolar gases in the various surfactants are seen to describe a series of slightly divergent straight lines displaced toward lower values relative to the gas solubilities in the hydrocarbon, n-octane (Fig. 33), or perfluorocarbon, perfluoroheptane (Fig. 3b), having characteristics resembling the core region of the respective micelles. Here one sees that the micellar solubility of a given nonpolar gas decreases as the carbon number of the surfactant ions decreases. This decrease in micellar gas concentration has been interpreted in terms of Laplace pressure effects which cause the activity coefficient of a solubilized gas molecule to increase progressively as micelle size decreases (6t, 6v). Turning to the gases of specific interest here, one sees in Fig. 3a that in every instance the solubility of CO2 in hydrocarbon-type micelles is greater than one would predict using the solubility data of the other gases as a basis for comparison. This suggests that CO2 differs from the other nonpolar gases in that molecules of this gas experience additional modes of attraction with hydrocarbon-type micelles which augment the usual dispersion interactions and lead to enhanced solubilization. A somewhat less clear picture emerges when one examines the solubility data obtained with the 6 AS is the case with water, complementary gas-phase solubility measurements, in which saturated vapor concentrations of alcohols in compressed gases are determined at elevated pressures, provide unequivocal evidence for the existence of a specific chemical aifinity between alcohols and the two gases CO2 and N:O (l 8, 19). Journal of Colloid and Interface Science, Vol. 125,No. 2, October1988
532
OWNBY, PRAPAITRAKUL, AND KING O
0
- 1.0 --
-1.O /
-2.0 -
--
--
O
_
3.0
J,,¢
-
2.O
--
-
3-0
--
Oi
-4.0
- 4 . 0 --
u ,/.~.
o
_ 0 " ~ []
O
-
5.0
-5-0 --
-
6.0
-6.0
x t:c
~ ~o l-
-
7.0
-
-8.0 -
-
/
.
[] SDS O SHSo
10.0
- 1 1 . 0
!
-100
C7F16 SPFO
9.0
-10-O
)1
I
I
AP 0 2 CH4 - 1 2 . 0
-
i
--
0 0
-8-0
z~ C T A B
go I
7.0
CO 2
I
I
C2H 6
C3H 8
-11.0
I
I
I
200
3OO
4OO
_
-12:0 1OO
.Jl
i
I
iLL
i
A r 0 2 CH4 C ~ C 0 2 N 2 O C 2 H 6
C3H8
I
I
I
200
300
400
TC(°K)
TC(°K)
(Q)
(b)
FiG. 3. One-atmosphere micellar gas solubilities plotted logarithmically as a function of gas critical temperature using T c for CO2 and N20: (a) Gas solubilities in micelles composed of surfactant ions having ordinary alkyl groups as tails Compared with solubilities in n-octane; (b) gas solubilities in micelles of sodium perfluorooctanoate compared with corresponding solubilities in perfluoroheptane. (See Fig. 1 for definitions of abbreviations.)
perfluorinated systems shown in Fig. 3b. Here one sees that the correlation obtained using the modified critical constants T~0 for the quadrupolar gases CO2 and N20 is considerably poorer than that obtained with n-octane. In fact, an equally good correlation is obtained using the experimental values of T c for these two gases. While this may be taken to indicate that polar interactions exist between these two gases and perfluoroheptane, the possibility exists that the poor correlation reflects the fact that isomeric mixtures of perfluoroheptane may have been employed in certain of the solubility measurements used in Fig. 3b. In particular, a sample of C7F16 containing mixed isomers was used for the solubility measurements of N 2 0 in perfluoroheptane carried out in this laboratory and recorded in Table II. Turning to the gas solubilities in micellar Journal of Colloid and Interface Science, Vol. 125, No. 2. O c t o b e r 1988
SPFO shown in the lower line of Fig. 3b, one sees that the micellar solubilities of CO2 and N20 lie well above the line defined by the other gases, suggesting that, as in the case of hydrocarbon type micelles, molecules of these two gases are subject to highly specific interactions within SPFO micelles which operate in addition to the usual dispersion forces causing enhanced solubilization of these two gases. One can readily construct arguments based on the special properties of COz and N 2 0 to explain the excess sorptive capacity of micelles toward these two gases. For example, the additional forces of attraction leading to excess solubilization can be attributed to electrostatic interactions between the ions and the counterions at the micelle surface and the molecular quadrupole moments of these two gases. One can equally well argue that chemical associa-
SURFACTANT SOLUTIONS
tion between these gases and water molecules concentrated in the head group region of the micelle is responsible for the enhanced solubilization observed for CO2 and N20. However, the fact that the solubilities of CO2 and N20 in nonpolar liquids correlate well with the solubilities of other gases makes it difficult to imagine that any property associated with the interior portion of a micelle can lead to a special affinity toward these gases. This leads one to conclude that, regardless of the specific nature of the interactions involved, the excess solubilization observed with CO2 and N20 is likely to be an adsorption phenomenon resuiting in the concentration of these gases at the surface regions of micelles. ACKNOWLEDGMENT The authors express appreciation for support provided by the National Science Foundation (NSF) under Grant CHE-8218288.
7.
REFERENCES 1. Hildebrand, J. H., and Scott, R. L., "The Solubility of Nonelectrolytes." Reinhold, New York, 1950; reprinted by Dover, New York, 1964, p. 248. 2. Coan, C. R., and King, A. D., Jr., J. Amer. Chem. Soc. 93, 1857 (1971). 3. Massoudi, R., and King, A. D., Jr., J. Phys. Chem. 78, 2262 (1974). 4. Massoudi, R., and King, A. D., Jr., J. Phys. Chem. 79, 1670 (1975). 5. Peterson, K. I., and Klemperer, W., J. Chem. Phys. 80, 2439 (1984). 6. (a) McBain, J. W., and O'Connor, J. J., J. Amer. Chem. Soc. 62, 2855 (1940); (b) McBain, J. W., and O'Connor, J. J., J. Amer. Chem. Soc. 63, 875 (194t); (c) McBain, J. W., and Soldate, A. M., J. Amer. Chem. Soc. 64, 1556 (1942); (d) Ross, S., and Hudson, J. B., J. ColloidSci. 12, 523 (1957); (e) Wishnia, A., 2". Phys. Chem. 67, 2079 (1963); (f) Winters, L. J., and Grunwald, E., J. Amer. Chem. Soc. 87, 4608 (1965); (g) Somasundaran, P., and Moudgil, B. M., J. Colloid Interface Sci. 47, 290 (1974); (h) Miller, K. W., Hammond, L.,
8. 9. 10.
11. 12. 13. 14. 15. 16. 17. 18. 19.
533
and Porter, E. G., Chem. Phys. Lipids 20, 229 (1977); (i) Matheson, I. B. C., and King, A. D., Jr., J. Colloid Interface Sci. 66, 464 (1978); (j) Hoskins, J. C., and King, A. D., Jr., 9". Colloid Interface Sci. 82, 260 (1981); (k) Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 82, 264 (1981); (1) Christian, S. D., Tucker, E. E., and Lane, E. H., J. Colloid Interface Sci. 84, 423 ( 1981); (m) DellaGuardia, L., and King, A. D., Jr., J. Colloidlnterface Sci. 88, 8 (1982); (n)Bolden, P. L., Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 91, 454 (1983); (o) Ben-Naim, A., and Wilf, J., J. Solution Chem. 12, 671 (1983); (p) Ownby, D. W., and King, A. D., Jr., J. Colloid Interface Sci. 101, 271 (1984); (q) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 106, 186 (1985); (r) Ben-Naim, A., and Batfino, R., J. Solution Chem. 14, 245 (1985); (s)Flanagan, H. L., Bolden, P. L., and King, A. D., Jr., J. Colloid Interface Sci. 109, 243 (1986); (t) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 112, 387 (1986); (u) Prapaitrakul, W., Shwikhat, A., and King, A. D., Jr. J. Colloid Interface Sci. 115, 443 (1987); (v) Prapaitrakul, W., and King, A. D., Jr., J. ColloM Interface Sci. 118, 224 (1987 ). Mukerjee, P., and Mysels, K. J., "Critical Micelle Concentrations of Aqueous Surfactant Systems," Nat. Stand. Ref. Data Series, Nat. Bur. Stand. (U.S.) NSRD S-NBS 36, Feb. 1961. Treiner, C., and Chattopadhyay, A. K., J. Colloid Interface Sci. 98, 447 (1984). Markham, A. E., and Kobe, K. A., J. Amer. Chem. Soc. 63, 449 (1941). Hildebrand, J. H., Prausnitz, J. M., and Scott, R. L., "Regular and Related Solutions." Van NostrandReinhold, New York, 1970. Kudchadker, A. P., Alani, G. H., and Zwolinski, B. J., Chem. Rev. 68, 659 (1968). Mathews, J. F., Chem. Rev. 72, 71 (1972). Cook, D., and Rowlinson, J. S., Proc. R. Soc. London Ser. A 219, 405 (1953). Rowlinson, J. S., Trans. Faraday Soc. 50, 647 (1954). King, A. D., Jr., J. Chem. Phys. 51, 1262 (1969). Gjaldbaek, J. Chr., and Niemann, H., Acta Chem. Scand. 12, 15 (1958). Boyer, F. L., and Bircher, L. J., J. Phys. Chem. 64, 1330 (1960). Gupta, S. K., Lesslie, R. D., and King, A. D., Jr., J. Phys. Chem. 77, 2011 (1973). Massoudi, R., and King, A. D., Jr., J. Phys. Chem. 77, 2016 (1977).
Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
W. PRAPAITRAKUL,
AND A. D. K I N G , JR.
Department of Chemistry, School of Chemical Sciences, University of Georgia, Athens, Georgia 30602 Received July 20, 1987; accepted November 24, 1987 Measurements have been made to determine the solubilities of carbon dioxide and nitrous oxide in aqueous solutions of cetyltrimethylammonium bromide (CTAB), sodium dodecyl sulfate (SDS), sodium 1-heptanesulfonate (SHSo), and sodium perfluorooctanoate (SPFO) at elevated pressures. The solubilities of each gas, measured at 26°C for solutions of CTAB and at 25°C for the other surfactants, follow Henry's law at all surfactant concentrations. Two distinctly different types of behavior are exhibited by SHSo. At concentrations below the critical micelle concentration (CMC), the addition of SHSo acts to suppress the solubility of CO2, while at higher concentrations, the opposite behavior is observed in that addition of SHSo results in an increase in CO2 solubility. The CMC values for the other surfactants are all much lower than that for SHSo. As a result, the solubility data for N20 and CO2 in these surfactants are all taken at concentrations in excess of the CMC and in every case, the addition of surfactant is found to enhance gas solubility. The increase in gas solubility observed above the CMC for each of the four surfactants is found to be a linear function of surfactant concentration as expected for micellar solubilization. The intramicellar solubilities of COz and N20 calculated from these latter data do not correlate well with solubilitiesof other nonpolar gasesbut rather are displaced toward higher concentrations, thus suggesting that the hydrophilic character of these two gases leads to enhanced solubilization. © 1988 Academic Press, Inc.
INTRODUCTION T h e gases c a r b o n dioxide (CO2) a n d nitrous oxide ( N 2 0 ) have very s i m i l a r physical p r o p erties. This arises f r o m the fact t h a t each gas is c o m p o s e d o f l i n e a r t r i a t o m i c m o l e c u l e s which are isoelectronic a n d have virtually identical d i m e n s i o n s a n d m o l e c u l a r weights. Like m o s t c o m m o n gases, CO2 a n d N 2 0 are basically h y d r o p h o b i c in n a t u r e in t h a t t h e y are c o n s i d e r a b l y m o r e soluble in oil t h a n in water. However, a c o m p a r i s o n o f the a q u e o u s solubilities listed in T a b l e II shows t h a t CO2 a n d N 2 0 nevertheless exhibit a n o t i c e a b l e affinity for water as e v i d e n c e d b y the fact t h a t a q u e o u s solubilities o f these two gases are an o r d e r o f m a g n i t u d e greater t h a n those o f the o t h e r n o n p o l a r gases listed. This o b s e r v a t i o n is o f course n o t new b u t r a t h e r was cited b y H i l d e b r a n d in 1950 as evidence that, in ad-
0021-9797/88 $3.00 Copyright © 1988 by Academic Press, Inc. All rights of reproduction in any form rcscrve,d.
d i t i o n to the usual physical forces, m o l e c u l e s o f dissolved C O 2 a n d N 2 0 i n t e r a c t with w a t e r t h r o u g h highly specific c h e m i c a l interactions, p r e s u m a b l y o f a Lewis a c i d - b a s e n a t u r e (1). T h e r e is a b u n d a n t evidence f r o m o t h e r sources to s u p p o r t this claim. F o r e x a m p l e , n o t o n l y d o solubility m e a s u r e m e n t s o f a reciprocal n a t u r e show t h a t w a t e r is a b n o r m a l l y soluble in c o m p r e s s e d gaseous CO2 a n d N 2 0 ( 2 ) b u t surface t e n s i o n d e t e r m i n a t i o n s at high pressures also provide clear evidence that these two gases are selectively a d s o r b e d at the g a s water interface (3, 4). M o r e recently, r a d i o frequency a n d m i c r o w a v e spectra o f m o l e c u l a r b e a m s c o n t a i n i n g CO2 a n d water v a p o r have p r o v i d e d direct evidence t h a t CO2 a n d w a t e r f o r m relatively rigid p l a n a r 1:1 c o m p l e x e s having C2 v s y m m e t r y with the hydrogen a t o m s d i r e c t e d a w a y f r o m the CO2 m o l e c u l e ( 5 ) .
526 Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
S U R F A C T A N T SOLUTIONS
This suggests that, as originally proposed by Hildebrand and Scott (1), the primary mode of bonding is Lewis acid-base in nature with the lone pair electrons of water interacting with unoccupied molecular orbitals of CO2. With this in mind, a series of measurements has been made to determine the solubilities of these same two gases, CO2 and N20, in micellar solutions of a wide variety of ionic surfactants. The surfactants chosen for this study consist of a cationic surfactant, CTAB, and three anionic surfactants, SDS, SHSo, and SPFO, that have tail groups which differ markedly with respect to chemical nature and size. A wide selection such as this is of obvious value in that it allows one to compare the micellar solubilities of these two gases with similar data obtained previously with other nonpolar gases and thus ascertain the extent to which this affinity toward water exhibited by CO2 and N20 affects the degree to which these gases are solubilized within micelles having markedly different surface and internal properties. 1 EXPERIMENTAL
The apparatus and procedures used to determine gas solubilities have been described in detail previously (6i, 6n). The sources and quoted purities of the surfactants used in these experiments are cetyltflmethylammonium bromide (98%), purchased from Lancaster Synthesis Ltd.; sodium dodecyl sulfate (99.0%) BDH Prod. No. 44244 (Lot No. 908811C), purchased from Gallard-Schlesinger Chemical Mfg. Corp.; and sodium 1-heptanesulfonate (99%), purchased from Lancaster Synthesis Ltd. The sodium periluorooctanoate was prepared in situ by titrating perfluorooctanoic acid (99% n-), SCM Specialty Chemicals (Lot No. 10439), with Fisher Certified ACS grade sodium hydroxide to pH 9-12. The CO2 and N20 used in these studies were commercial grade (quoted purity of 99.5% min) and CP 1 Reference (6) constitutes a reasonably complete bibliography on the subject of gas solubilization.
527
grade (quoted purity of 99.0%), respectively, purchased from Matheson Gas Products. RESULTS A N D DISCUSSION
The gas solubilities determined for CO2 and N20 in aqueous solutions of the various surfactants are listed in Table I. These data are shown plotted as a function of surfactant concentration in Fig. 1. Here it is seen that the solubility data fall into two classes. The gas solubilities obtained with SDS, CTAB, and SPFO, all of which have low CMC values, 2 are seen to increase linearly with surfactant concentration, thus generating a series of straight lines having positive slopes whose origins coincide with the corresponding gas solubilities measured in pure water. This is the typical pattern of behavior one encounters with micellar solutions of surfactants in which the micelles responsible for solubilizing the gas molecules maintain a constant sorptive capacity which is independent of surfactant concentration. The slopes of the straight lines defined by these solubility data correspond to the mole ratios ofsolubilized gas molecules to surfactant ions in micellar form. Therefore, in view of the high dilutions involved, the numerical magnitudes of the slopes of the gas solubility data for these surfactants can be taken to equal the 1-atm mole fraction solubilities of each gas within the individual micelles, X TM, i = CO2, N20. The micellar gas solubilities derived from straight lines fitted to the experimental solubility data for SDS, CTAB, and SPFO using the least-squares criteflon are shown listed in columns 4-6 of Table II. The CMC for SHSo, on the other hand, is quite large in comparison with the other surfactants, having a value of 0.28 M(0.3 m) at 25°C (6t). As a consequence, the CO2 solubilities obtained in these experiments for aqueous solutions of this surfactant span a wide range of concentrations both above and 2 Literature values and associated references for these surfactants at 25°C are SDS, 8.1 × 10 -3 M ( 7 ) ; CTAB, 9.2 × 1 0 - 4 M ( 7 ) ; SPFO, 3.2 × 10-2 M ( 8 ) . Journal of Colloid and Interface Science, Vol. 125,No. 2, October 1988
528
OWNBY, PRAPAITRAKUL, AND KING TABLE I Solubility of CO2 and N20 in Surfactant Solutions Expressed as Moles Gas per Atmosphere in 1000 g H20 Gas solubility (m) × 102a Surfactant
Temperature (°C)
Sodium heptylsulfonate
25
Sodium dodecylsulfate
25
Cetyltrimethylammoniumbromide
26
Sodium perfluorooctanoate
25
Surfactant concentration (m)
CO:
N20
0 0.2 0.3 0.4 0.5 0.6 0.8 1.0 1.2 0.2 0.4 0.6 0
3.32, 3.392b 3.17 3.12 3.05 3.05 3.07 3.16 3.26 3.33 3.54 3.69 3.91 3.28
2.433b
0.1
3.46
0.2 0.3 0.4 0.2 0.4 0.6
3.60 3.73 3.88 3.61 3.90 4.16
2.67 2.94 3.15
a Averageerror: +0.02 × 10-2m. b Recommended values from Wilhelm, E., Battino, R., and Wilcock, R. J., Chem. Rev. 77, 219 (1977). below the CMC. Figure 1 shows that C O 2 solubility in SHSo solutions exhibits a distinctly different behavior in each of these regions. Below the CMC, it is seen that SHSo acts to suppress the solubility of CO2. One concludes that SHSo ions in monomeric form "salt out" dissolved CO2. Salting out effects are c o m m o n l y reported as Setchenow coefficients, k, defined as k = ( 1 / m ) l o g ( S ° / S ) , where S Oand S represent the solubilities of the substance of interest in pure water ( S °) and in a solution of dissolved electrolyte (S), with m denoting the molal concentration of the electrolyte. The CO2 solubility data for pure water and 0.2 and 0.3 m solutions of SHSo, when fitted to the expression above, yield a value of k = 0.089. It is interesting to compare the salting out effect of SHSo m o n o m e r s toward CO2 to that of other c o m m o n sodium salts. M a r k h a m and Kobe (9) have measured the solubility of CO2 in a number of solutions of strong electrolytes. The data they report at low salt concentrations Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
(0 < m < 0.2) yield values o f k = 0.108 and k = 0.087 for NaCI and NaNO3, respectively, at 25 °C. Thus one concludes that m o n o m e r i c SHSo and sodium nitrate are about equally effective in salting out CO2. At concentrations in excess of the CMC, this negative trend reverses and the CO2 solubility increases with surfactant concentration as expected when micelles contribute to solubilization. The slope of the linear portion of the SHSo solubility data ranging between 0.6 and 1.2 m is m u c h smaller than the corresponding slopes of the other systems, indicating that SHSo micelles have a smaller sorptive capacity toward CO2 in comparison with the other surfactants. The mole fraction solubility of CO2 in micellar SHSo corresponding to this slope is listed in column 3 of Table II along with those for other gases. Previous investigators (10) have shown that 1-atm gas solubilities measured for a given nonpolar solvent, when plotted logarithmi-
SURFACTANT SOLUTIONS 4.5 CO 2 0/ 4.0
E
0 SPFO A CTAB [] SDS
/
/ 3.5
/
@ SPFO
0 SHSo
~
"7
0 ~"
-~o.<~
> o~J~-
N20
/~j
o~
3.02.5I / / ' / ~ / / / ° / @
2.0 f l 0
I I I I I [ I I I [ I I I I I 0,2
0.4
0.6
0-8
1.0
1-2
1-4
1-6
SURFACTANT CONCENTRATION (moles kg-1)
FIG. 1. Moles of gas absorbed per atmosphere in 1000 g of H20 with added surfactant shown as a function of surfactant concentration at 25°C (data for CTAB taken at 26°C). CsHIs, n-octane; CTAB, hexadecyltrimethylammonium bromide; SDS, sodium dodecylsulfate; SHSo, sodium heptylsulfonate; C7F16, perfluoroheptane; and SPFO, sodium perfluorooctanoate.
529
the cases of the nonpolar gases, T c and AEbv are proportional to one another. 3 There is an added advantage to using T c as a scaling parameter in situations involving gases having large quadrupole m o m e n t s such as CO2 and N20. This stems from the fact that a statistical treatment of the thermodynamic properties of such systems by Rowlinson (13, 14) has shown that the effects of angle-dependent interactions can be factored out from those arising from spherically symmetric dispersion forces in a m a n n e r consistent with the principle of corresponding states, thus allowing one to define an effective critical temperature, T0c, which characterizes the contributions that dispersion forces alone make to the overall intermolecular interactions between molecules of such gases. This is important because in cases where quadrupolar gases such as CO2 or N 2 0 interact with nonpolar substances, dispersion forces constitute the major m o d e of interaction so that T c rather than T c is the appropriate scaling factor to use as a measure of the extent of interaction. According to the treatment by Rowlinson, T0c is related to the experimental critical temperature, T c, by a single parameter, 6c, according to the equation T c = T0C(1 + 2c3c).
cally, correlate well with the internal energy of vaporization of the individual gases at their respective boiling points, AE v, thus suggesting that the Gibbs free energy of solution is a linear function of AE~. Similar trends are found for chemically inert gases dissolved in moderately polar solvents. Unfortunately, the triple point of CO2 lies above 1 atm, thus precluding one from using any such scheme based on AE v to compare solubility data for CO2 with that of other gases. There are, however, a variety of other properties (boiling points, Lennard-Jones force constants, critical temperatures, etc.) that m a y be used as scaling parameters. O f these, gas critical temperature, T c, is undoubtedly the most suitable for our purposes here because not only are T c values known to a high degree of precision for the gases of interest here, including CO2, but in
[1]
The average values for 6c derived from a wide variety of thermodynamic data are found by Rowlinson (14) to be 6c = 0.10 for CO2 and 6c = 0.06 for N20. 4 Values of T~o calculated using Eq. [1] and these parameters are listed 3 AEv values taken from Ref. (10) for 11 light gases (He, Ne, Ar, Kr, Xe, H2, N2, 02, CO, C H 4 , and C2H6) when plotted against Tc (Refs. (11, 12)) yield a well-defined straight line (correlation coefficient= 0.99663 ) having essentially zero intercept (-0.1 kcal mole-l) and a slopewhich, when dividedby the gas constant, corresponds to A E V / R T c = 5.1. 4 The theoretical treatment by Rowlinson presupposes that angular-dependent interactions between axiallysymmetric moleculesact as a perturbation on the centrosymmetric dispersion force term of the pair potential and thus varieswith intermolecularseparation,r, as r -6. This closely approximates the functional form of the potential arising from quadrupole-quadrupole interactions which variesas r -3 before averaging,and it has been shown (15 ) that the Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
530
OWNBY, PRAPAITRAKUL, AND KING TABLE II Gas Solubilities at 25°C and 1 arm in Surfactants and Bulk Solvents (Mole Fraction × 104)
Gas
Tc (K)
SHSo
SDS
CTAB
Ar O2 CH4 CF4 CO2 N20 C2H6 Call8
150.75 154.58 190.55 227.6 304.15 (253) 309.57 (276) 305.43 369.82
41 41 101 -42 a -451 1021
112 135 102 105 192 295 --93 a 147 a --1 1 0 3 1555 3304 4805
SPFO
CHaOH
171 4.156 181 4.117 191 8.676 381 . 139 a 56.3 s 120 a 65 b 611 40.56 1481 130 b
CH3(CH2)2OH CH3(CH2)7OH
7.826 6.807 16.16 . . 68.07 84 b 90.06 318 b
12.739 11.329 26.879
n-CaHjs
23.369 21.779 50-269
. 93.019 134 b 1736 607 b
119.89 175 b 3241° 1200 I1
C7FI6
H20
5612 6012 7312 13612 208.813 278 b 21712 61212
0.25214 0.23014 0.251 TM -6.11114 4.383 TM 0.334 TM 0.270 TM
This work. Estimated error for micellar solubilities of CO2 and N20:_+4 × 10-4. b This work. Estimated errors for solubilities in alcohols, n-octane and perfluoroheptane: N20, _+2 × 10-4; C3H8, +_4 × 10-4. l Ref. (6t). 2 Ref. (6i). 3 Ref. (6n). 4 Ref. (6p). 5 Ref. (6q); T = 26°C. 6 Ref. (17). 7 From data of Tokunga, J., J. Chem. Eng. Data 20, 41 (1975). 8 From data of Kunerth, W., Phys. Rev. 19, 512 (1922). 9 Wilcock, R. J., Battino, R., Danforth, W. F., and Wilhelm, E., J. Chem. Thermodyn. 10, 817 (1978). 1° Hayduk, W., and Cheng, S. C., Can. J. Chem. Eng. 48, 93 (1970). 11Thomsen, E. S., and Gjaldbaek, J. Chr., Acta Chem. Scand. 17, 127, 134 (1963). 12 Ref. (6v). 13 Kobatake, Y., and Hildebrand, J. H., J. Phys. Chem. 65, 331 (1961). 14Wilhelm, E., Battino, R., and Wilcock, R. J., Chem. Rev. 77, 219 (1977). a
in parentheses in column 2 of Table II for these two gases. One-atmosphere solubilities of CO2, N20, and other light gases dissolved in the bulk solvents n-octane, perfluoroheptane, water, and three representative alcohols are listed in Table II. These data, when plotted logarithmically as a function of gas critical temperature as seen in Fig. 2, reveal several features of interest. First, it is seen that when plotted in this manner, the solubility data for nonpolar gases in a given solvent correlate well with gas critical temperature as expected. Second, since elec-
values found by Rowlinson (14) of 6c = 0.01 and 0.06 for CO2 and 1'420, respectively, can be accounted for almost exactly by assuming that quadrupole--quadrupole interactions are solely responsible for the angle dependent forces between these molecules. Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988
trostatic interactions of quadrupolar origin are not expected to affect the solubilities of CO2 or N20 in n-octane, T~0rather than T c should be the appropriate scaling parameter for these two gases. This is borne out by the fact that, when plotted as a function of T~0, the solubility data for these two gases in n-octane indeed coincide well with the line defined by the other nonpolar gases. Third, one sees that the solubilities of nonpolar gases in the various alcohols describe a series of nearly parallel lines which exhibit a progressive displacement toward decreased solubility as the carbon number of each alcohol decreases. The displacement of these lines toward lower solubilities simply reflects the behavior expected of chemically unreactive nonpolar gases in that the solubilities of such gases invariably decrease as the internal pressure of the solvent in-
SURFACTANT SOLUTIONS
-1.0 - 2.O
-
o/
-4.0
%
CIOH
-5-0
/
•
-9.0 -lO.O --%--0 -11.0 --12.0 IO0
)1
-
I
-
-
-
o
I I
Ar 0 2 C H 4
I
CO2 N20 C2H6
o-
H20
I C3H 8
I
I
I
200
300
400
TC(°K] FIG. 2. One-atmosphere gas solubilities of a variety of gases in n-octane (C8), 1-0ctanol (CsOH), l-propanol (C3OH), methanol (CIOH), and water plotted logarithmically as a function of gas critical temperature using T0c for CO2 and N20.
creases.5 However, in contrast to their behavior in n-octane, the solubilities of CO2 and N20 in the various alcohols and water are seen, in every case, to lie above the line defined by the solubilities of the other gases. As noted earlier, the enhanced solubilities of CO2 and N20 relative to the other gases in water have been cited by Hildebrand as evidence for chemical association between these gases and water (1). 5 Hildebrand solubility parameters in conjunction with regular solution theory provide a convenient measure of the role played by solvent internal pressure in determining gas solubility (Ref. (1, Chap. 15)). Regular solution theory predicts that gas solubilities, expressed logarithmically, will decrease in a regular fashion as the solubility parameter of the solvent increases. This trend has been shown to hold for alcohols as well as nonhydrogen bonding solvents (16, 17). The solubility parameters of the solvents used in this study, expressed as (cal em-3)1/2are n-CsH1a, 5.7; CH3(CH2)7OH, 9.3; CH3(CH2)2OH, 11.9; CH3OH, 14.4; H20, 23.4.
531
Therefore, although effects of dipole-quadrupole interactions cannot be ruled out, it is reasonable to assume that the excess solubility exhibited by CO2 and N20 in the various alcohols is likewise a manifestation of a chemical association resulting from highly specific interactions between molecules of CO2 and N20 and the hydroxyl group common to these alcohols. 6 Figures 3a and 3b show the micellar gas solubilities contained in columns 3-6 of Table II plotted logarithmically as a function of gas critical temperature. The micellar solubilities of the nonpolar gases in the various surfactants are seen to describe a series of slightly divergent straight lines displaced toward lower values relative to the gas solubilities in the hydrocarbon, n-octane (Fig. 33), or perfluorocarbon, perfluoroheptane (Fig. 3b), having characteristics resembling the core region of the respective micelles. Here one sees that the micellar solubility of a given nonpolar gas decreases as the carbon number of the surfactant ions decreases. This decrease in micellar gas concentration has been interpreted in terms of Laplace pressure effects which cause the activity coefficient of a solubilized gas molecule to increase progressively as micelle size decreases (6t, 6v). Turning to the gases of specific interest here, one sees in Fig. 3a that in every instance the solubility of CO2 in hydrocarbon-type micelles is greater than one would predict using the solubility data of the other gases as a basis for comparison. This suggests that CO2 differs from the other nonpolar gases in that molecules of this gas experience additional modes of attraction with hydrocarbon-type micelles which augment the usual dispersion interactions and lead to enhanced solubilization. A somewhat less clear picture emerges when one examines the solubility data obtained with the 6 AS is the case with water, complementary gas-phase solubility measurements, in which saturated vapor concentrations of alcohols in compressed gases are determined at elevated pressures, provide unequivocal evidence for the existence of a specific chemical aifinity between alcohols and the two gases CO2 and N:O (l 8, 19). Journal of Colloid and Interface Science, Vol. 125,No. 2, October1988
532
OWNBY, PRAPAITRAKUL, AND KING O
0
- 1.0 --
-1.O /
-2.0 -
--
--
O
_
3.0
J,,¢
-
2.O
--
-
3-0
--
Oi
-4.0
- 4 . 0 --
u ,/.~.
o
_ 0 " ~ []
O
-
5.0
-5-0 --
-
6.0
-6.0
x t:c
~ ~o l-
-
7.0
-
-8.0 -
-
/
.
[] SDS O SHSo
10.0
- 1 1 . 0
!
-100
C7F16 SPFO
9.0
-10-O
)1
I
I
AP 0 2 CH4 - 1 2 . 0
-
i
--
0 0
-8-0
z~ C T A B
go I
7.0
CO 2
I
I
C2H 6
C3H 8
-11.0
I
I
I
200
3OO
4OO
_
-12:0 1OO
.Jl
i
I
iLL
i
A r 0 2 CH4 C ~ C 0 2 N 2 O C 2 H 6
C3H8
I
I
I
200
300
400
TC(°K)
TC(°K)
(Q)
(b)
FiG. 3. One-atmosphere micellar gas solubilities plotted logarithmically as a function of gas critical temperature using T c for CO2 and N20: (a) Gas solubilities in micelles composed of surfactant ions having ordinary alkyl groups as tails Compared with solubilities in n-octane; (b) gas solubilities in micelles of sodium perfluorooctanoate compared with corresponding solubilities in perfluoroheptane. (See Fig. 1 for definitions of abbreviations.)
perfluorinated systems shown in Fig. 3b. Here one sees that the correlation obtained using the modified critical constants T~0 for the quadrupolar gases CO2 and N20 is considerably poorer than that obtained with n-octane. In fact, an equally good correlation is obtained using the experimental values of T c for these two gases. While this may be taken to indicate that polar interactions exist between these two gases and perfluoroheptane, the possibility exists that the poor correlation reflects the fact that isomeric mixtures of perfluoroheptane may have been employed in certain of the solubility measurements used in Fig. 3b. In particular, a sample of C7F16 containing mixed isomers was used for the solubility measurements of N 2 0 in perfluoroheptane carried out in this laboratory and recorded in Table II. Turning to the gas solubilities in micellar Journal of Colloid and Interface Science, Vol. 125, No. 2. O c t o b e r 1988
SPFO shown in the lower line of Fig. 3b, one sees that the micellar solubilities of CO2 and N20 lie well above the line defined by the other gases, suggesting that, as in the case of hydrocarbon type micelles, molecules of these two gases are subject to highly specific interactions within SPFO micelles which operate in addition to the usual dispersion forces causing enhanced solubilization of these two gases. One can readily construct arguments based on the special properties of COz and N 2 0 to explain the excess sorptive capacity of micelles toward these two gases. For example, the additional forces of attraction leading to excess solubilization can be attributed to electrostatic interactions between the ions and the counterions at the micelle surface and the molecular quadrupole moments of these two gases. One can equally well argue that chemical associa-
SURFACTANT SOLUTIONS
tion between these gases and water molecules concentrated in the head group region of the micelle is responsible for the enhanced solubilization observed for CO2 and N20. However, the fact that the solubilities of CO2 and N20 in nonpolar liquids correlate well with the solubilities of other gases makes it difficult to imagine that any property associated with the interior portion of a micelle can lead to a special affinity toward these gases. This leads one to conclude that, regardless of the specific nature of the interactions involved, the excess solubilization observed with CO2 and N20 is likely to be an adsorption phenomenon resuiting in the concentration of these gases at the surface regions of micelles. ACKNOWLEDGMENT The authors express appreciation for support provided by the National Science Foundation (NSF) under Grant CHE-8218288.
7.
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Journal of Colloid and Interface Science, Vol. 125, No. 2, October 1988