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The influence of pH on gas solubilities in aqueous solutions of sodium octanoate at 25°C
The Influence of pH on Gas Solubilities in Aqueous Solutions of Sodium Octanoate at 25°C W. P R A P A I T R A K U L ,
A. S H W I K H A T ,
AND A. D. K I N G , JR.
Department of Chemistry, School of Chemical Sciences, University of Georgia, Athens, Georgia 30602 Received February 4, 1986; accepted May 6, 1986 Measurements have been made to determine the solubilities of oxygen, methane, ethane, and propane in aqueous solutions of sodium octanoate at 25°C and at pH values ranging from 7 to 13. The solubility of each gas obeys Henry's law at all surfactant concentrations and pH values. The solubilities measured for each gas are effectivelyindependent of sodium octanoate concentration as well as pH at concentrations well below the CMC, and, within experimental error, are the same as the respective gas solubilities in pure water at 25°C. Above the CMC, the solubility of each gas increases with surfactant concentration indicating micellar solubilization. The degree of solubilization is greatest for propane and decreases in the order: propane > ethane > methane > oxygen. However, unlike the situation at low surfactant concentrations, the individual gas solubilities vary with pH at higher surfactant concentrations. The degree of solubilization is found to be quite constant for each gas at pH values greater than 9; and intramicellarsolubilities derived from gas solubility data taken at pH 12 are identical to those determined for a similar surfactant having a seven carbon alkyl chain, sodium 1-heptane sulfonate. However, the extent of solubilization for each gas increases rapidly as the solution pH decreases from 9 to 7, suggesting that mixed micelles composed of octanoic acid and sodium octanoate are more effective solubilizing agents than those of the pure soap. © 1987AcademicPress,Inc. INTRODUCTION Surface active salts o f long-chain c a r b o x y l i c acids are generally designated as soaps. As a class, soaps differ f r o m m o s t o t h e r a n i o n i c surfactants in t h a t the m o d e r a t e l y basic carb o x y l a t e h e a d g r o u p s are subject to hydrolysis equilibria (1-3). As a result, the h y d r o p h i l i c c h a r a c t e r o f these surfactants is sensitive to the acidity o f t h e a q u e o u s e n v i r o n m e n t a n d p r o p e r t i e s such as surface t e n s i o n (4), interfacial t e n s i o n (5), d e t e r g e n c y (6), a n d e m u l sifying characteristics (7) o f soap s o l u t i o n s all exhibit a strong d e p e n d e n c e on pH. This p a p e r e x a m i n e s the role that p H plays in d e t e r m i n i n g the extent o f solubilization o f gases.1 It reports the results o f a series o f solubility m e a s u r e m e n t s at 25 °C involving four gases with widely differing p r o p e r t i e s ( 0 2 , CH4, C2H6, a n d l Reference (8) constitutes a reasonably complete bibliography on the subject of gas solubilization in micellar solutions.
C3H8) dissolved in a q u e o u s solutions o f a typical soap, s o d i u m octanoate, at p H values ranging b e t w e e n 7 a n d 13. EXPERIMENTAL T h e m e t h o d used to d e t e r m i n e gas solubilities has b e e n previously described in detail (8m). It differs from other m e t h o d s c o m m o n l y used to d e t e r m i n e gas solubility in t h a t it involves m e a s u r i n g the v o l u m e o f gas released f r o m solutions s a t u r a t e d at elevated pressures. T h e p r o c e d u r e consists o f three steps: (i) the surfactant s o l u t i o n is allowed to equilibrate with the gas o f interest at a n elevated pressure in a t h e r m o s t a t t e d b o m b e q u i p p e d with a m a g n e t i c a l l y d r i v e n stirrer; (ii) the stirrer is s t o p p e d a n d the solution is allowed to b e c o m e still, w h e r e u p o n the pressure is released; a n d (iii) after a short delay ( ' ~ 30 s) to allow t h e r m a l e q u i l i b r i u m to b e reestablished following the release o f gas, the m a g n e t i c stirrer is again ac-
443 0021-9797/87 $3.00 Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
Copyright © 1987 by Academic Press, Inc. All rights of reproduction in any form reserved.
444
PRAPAITRAKUL, SHWIKHAT, AND KING
tivated causing the dissolved gas to effervesce from the formerly supersaturated solution. The volume of gas released is measured manometrically at ambient pressure. In this manner one determines the number of moles of gas released, An, as the solution proceeds isothermally from what was previously an equilibrium state at the upper pressure to a new equilibrium state at ambient pressure. The difference in pressure between these two states, AP, is numerically equal to the gauge pressure recorded during the initial equilibration step (i). The pressures employed in these experiments were such that the amounts of dissolved gas were 30 vol% or less, measured at ambient pressure. The series of steps, (i)(iii), are repeated one or more times at different pressures in order to establish that the ratios 2xn/AP are independent of pressure. This was found to be the case in every instance so that the averaged values of An/&P listed in Table I can be taken to represent gas solubilities in the respective solutions at a partial pressure of 1 atm with Henry's law being satisfied for each system.
The sodium octanoate used in all experiments was produced in situ by titrating aqueous suspensions containing known amounts of octanoic acid (Eastman Cat. No. 665, Lot Nos. 691 and 692) with a concentrated solution of sodium hydroxide (Baker analyzed reagent grade) to the desired pH. The pH measurements were performed with a Fisher Acumet Model 825 MP pH meter using standard buffers and two point calibration to standardize the instrument. The gases used were CP grade or the equivalent, having quoted purities of 99.0% or better for the hydrocarbon gases which were purchased from Matheson and 99.6% for the oxygen which was obtained from Selox Corp. Doubly distilled water was used to prepare all solutions. RESULTS AND DISCUSSION
Gas solubilities determined as a function of sodium octanoate concentration at pH 7.5 and 12 are listed in Table I. These data are shown plotted as a function of soap concentration in Figs. 1-4. Inspection of the data shown in Figs.
TABLE I Gas Solubilities in Aqueous Solutions of Sodium Octanoate at pH 12 and pH 7.5 Expressed as Moles Gas per Atmosphere in 1000 g H20 at 25°C Gas solubility (m) × 103 ,o Sodium octanoate concentration (m)
02
C1J~
C2I%
C3Hs
0 0.2 0.3 0.4 0.5 0.6 0.8 1.0 1.2
1.41 c 1.35 (1.37) - - (1.33) 1.34 (1.38) 1.28 (--) 1.33 (1.49) 1.39 (1.47) 1.44 (1.58) 1.57 (1.64)
1.55 c 1.46 (1.61) - - (1.52) 1.50 (1.53) 1.47 (--) 1.56 (1.88) 1.82 (2.07) 1.95 (2.31)
1.76 c 1.83 (1.87) - - (1.87) 2.01 (2.00) 2.08 (--) 2.48 (3.32) 3.49 (4.71) 4.48 (6.17)
1.42 d 1.4 (1.4) - - (1.5) 1.6 (2.4) 2.1 ( - - ) 3.0 (5.1) 5.0 (8.6) 7.2 (11.5)
a First value listed for solubility is for solution having pH 12. Second value in parenthesis is for solution having pH 7.5. bAverage errors of solubilities measured in these experiments for 02, CH4, C2H6, and C3Hs are +1 × 10-5, +2 × 10-5, +4 × 10-5, and _1 × 10-4, respectively. c Data taken from Ref. (8m). a Datum taken from Ref. (8p).
Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
GAS SOLUBILITY 2.0
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FIG. 1. Moles of oxygen absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12. 1-4 reveals that within experimental error, the solubility o f each gas remains the same as that in pure water until the concentration o f sod i u m octanoate reaches approximately 0.4 m at which point the solubility o f each gas begins to increase with surfactant concentration. The C M C of sodium octanoate has been determ i n e d to be 0.405 m at 2 5 ° C (9) suggesting that, as found previously with other surfactants (8), the e n h a n c e d solubility observed for each gas above 0.4 rn is the result o f solubilization by micelles present at soap concentrations in excess o f the CMC. A c o m p a r i s o n o f the data shows that above the CMC, the degree to which the solubility is e n h a n c e d by the soap
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FIG. 3. Moles of ethane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
micelles is greatest for propane and decreases in the order o f decreasing boiling point for the individual gases: C3H8 (231 K) > C2H6 (184 K) > CH4 (109 K) > 02 (90 K). This trend with boiling point is quite general and is c o m m o n l y observed with gases solubilized in a wide variety o f surfactants (8r). It is the behavior expected for systems in which dispersion forces constitute the d o m i n a n t m o d e o f interaction between solute and solvent molecules. A n o t h e r feature o f interest in the data o f Figs. 1-4 is the fact that the degree to which each gas is solubilized in sodium octanoate micelles is a function of the p H o f the solution, with the micelles absorbing m o r e gas at the lower pH. This unusual feature is not c o m m o n to anionic surfactants having strong acid head groups 2 and is u n d o u b t e d l y related to the p H dependent hydrolysis o f the mildly basic carboxylate head group. In general, the paucity o f publications on the topic o f gas solubilization (Ref. (8) contains
I 1,0
SODIUMOCTANOATECONC. (mol kg-1]
FIG. 2. Moles of methane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
2 The insensitivity toward pH exhibited by anionic surfactants is illustrated by the fact that the solubility of propane in 0.5 rn sodium octylsulfate at 25°C remains constant, having a value of(7.5 + 0.1) X 10 -3 mole C3H8kg-1 atm -~, at pH values ranging between 2.5 and 12. Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
446
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SODIUMOCTANOATE CONC, (m01 kg -I) FIO. 4. Moles of propane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
only 18 citations) precludes frequent comparisons between data obtained in different laboratories. It is therefore very fortunate that the solubility data obtained here with the hydrocarbon gases, methane, ethane, and propane, at pH 12 can be compared directly with results obtained previously by Ben-Naim and Wilf
(8n) for these same gases under identical conditions using a different experimental method. Within experimental error, perfect agreement is found between solubilities measured here for ethane and propane and those measured by Wilf and Ben-Naim. However, in the case of methane, a significant discrepancy exists between the results obtained in the two different laboratories. Below the CMC, the solabilities determined here for CH4 are as much as 13% greater than those published by BenNaim and Wilf. As a consequence, no evidence is found in Fig. 2 for the puzzling minimum in solubility observed by Ben-Naim and Wilf for methane above the CMC of sodium octanoate. The rising portions of the data shown in Figs. 1-4 extend in a linear fashion to quite high soap concentrations, suggesting that the sorptive capacity of the individual micelles remains constant, even at concentrations as great as 1.2 m. If it is assumed that soap monomer concentration remains effectively constant at concentrations greater than the CMC, then the slopes of the rising portions of Figs. 1-4 can be equated to the mole ratio ofsolubilized gas to micellized soap. Since the solutions are all quite dilute with respect to dissolved gas, these mole ratios may in turn be taken to represent mole fraction solubilities of the various gases solubilized within the sodium octanoate micelles at 1 atm, X~. The micellar gas solubilities obtained in this manner are listed in Table II along with values obtained previously for
TABLE II Micellar Solubilities of Gases in Sodium Octanoate at 25°C and 1 atm (Mole Fraction × l04)
Gas
0 2
CH4 C2H6 C3Hs
Sodium octanoate a o H 7.5
Sodium octanoat¢ a p H 12.0
Sodium l-heptane sulfonate b
Octanoic acid c
3 13 70 150
4 10 49 100
4 10 45 102
--190 700
"Estimated error: 02, +1 × 10-4; CH4, +1 × 10-4; C2H6, _+1 × 10-4; C3H8,_+4 × 10-4. b Data from Ref. (80. c Estimated error: C2H6, _+4 × 10-4; C3H8, _+10 X 10-4. Journal of Colloid and Interface Science, Vol. I 15, No. 2, February 1987
447
GAS SOLUBILITY
the closely related surfactant, sodium 1-heptane sulfonate. Previous studies in this laboratory have shown that alkyl-chain length is the dominant factor in determining the sorptive capacity of miceUes derived from ionic surfactants (8m, 8r). Therefore, since hydrolysis effects are expected to be negligible at pH 12, it is reasonable to expect the micellar solubilities of 02, C H 4 , C 2 H 6 , and C3H8 to have the same values in sodium octanoate at pH 12 as determined previously with sodium 1heptane sulfonate (8r). A comparison of the micellar solubilities listed in the middle columns of Table II shows this to be the case. The micellar solubilities measured at pH 7.5 equal or exceed those measured at pH 12 and, with the exception ofO2 for which any change in sorptive capacity is masked by experimental error, it appears that reducing the pH from 12 to 7.5 results in approximately a 40% increase in micellar gas solubility. In terms of percentage, this observed enhancement is considerably greater than the small decrease in CMC caused by the reduction in pH. Clearly, the enhanced gas solubility accompanying the reduction in pH cannot be explained in terms of a simple reapportionment of monomer anions into micelles as might be expected if a simple electrolyte were added to the solution. Rather, some other process producing a much greater increase in the sorptive capacity of the micellized octanoate ions must be involved. In order to more carefully examine the effect o f p H on gas solubility, a series of experiments was performed in which the solubilities of ethane and propane were determined in 0.8 m solutions of sodium octanoate to which small amounts of sodium hydroxide or hydrochloric acid had been added in order to adjust the pH of the solutions to values falling within the range 7-13. The lower limit of these experiments was dictated by the fact that below pH 7 a precipitate composed of a mixture of octanoic acid and sodium octanoate forms, causing the solution to become a gel. The resuits of these experiments are listed in Table III. These data, presented graphically in Fig. 5, reveal that the solubilities of ethane and
TABLE III The Solubility of Ethane and Propane in 0.8 m Sodium Octanoate at 25°C as a Function o f p H C2I~
C3Hs
pH
Solubility (m) × 10s °
7.07 7.20 7.41 7.81 8.11 8.89 9.28 10.16 11.54 12.20 12.52
5.02 4.65 4.01 3.67 3.42 3.11 3.38 3.36 3.27 3.27 3.41
pH
Solubility (m) × 103a
7.17 7.40 7.55 7.80 8.20 9.68 11.60 13.15
11.4 11.8 9.9 6.9 5.1 4.6 4.4 4.7
a Average errors in solubility found in footnote b of Table I.
propane remain quite constant as the pH falls from 13 to 9 and then increase rapidly as the pH of the solution drops from pH 9 to 7. The phase behavior and pH characteristics of sodium octanoate/octanoic acid mixtures having a formal concentration of 0.8 M are summarized graphically in Fig. 6. 3 These data show that octanoic acid is 100% converted into the salt form, i.e., pure soap is present, over the pH range of 9-13 for which low but constant gas solubilities are found for both ethane and propane in Fig. 5. Therefore, one concludes that the gas solubilities recorded at pH 12 for 02, CH4, C2H6, and C3H8 in Table I and Figs. 1-4 reflect micellar solubilization by pure sodium octanoate. However, it is also clear from Fig. 6 that a significant conversion of octanoate anion to the acid form occurs as the pH falls from 9 to 7. The solubility of octanoic acid in water is quite low and is esti3 The data used to construct Fig. 6 were obtained using solutions/suspensions formed by adding known a m o u n t s of reagent grade sodium hydroxide to preweighed samples of octanoic acid. Doubly distilled water was used to dilute each solution/suspension to a formal concentration of 0.8 M with respect to the acid. Journal of Colloid and Interface Science, Vol. 115,No. 2, February 1987
448
PRAPAITRAKUL, SHWIKHAT, AND KING
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pH FIG. 5. Solubilities of ethane and propane in 0.8 m solutions of sodium octanoate at 25°C shown as a function of pH. Inset: Solubilities of ethane and propane shown as a function of' micelle composition. ©, C2H 6; O, C3H8.
mated to be only 5.1 × 10 -3 M a t 2 5 ° C . 4 Since the solutions remain transparent at these pH values, one must assume that octanoic acid is solubilized within micelles derived from the remaining sodium octanoate. Therefore, it is reasonable to conclude that the increase in gas solubility observed with failing pH is the result of an internal conversion of sodium octanoate to octanoic acid resulting in the formation of mixed micelles having enhanced sorptive capacities toward these gases. With this in mind, one can use the pHcomposition data of Fig. 6 to recast the solubilities of ethane and propane as a function of micellar composition as shown in the inset of Fig. 5. Although there is considerable scatter among the data, as might be expected from 4 The solubility of octanoic acid at 25 °C is interpolated from solubility data taken from Refs. (10, 11). Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
accumulated errors in the calculations, the solubilities of both gases clearly exhibit a linear dependence on octanoic acid concentration. The slopes of the individual lines which have been fitted to the data according to the leastsquares criterion are 9.3 X 10 -3 and 47.3 X 10 -3 mole kg -1 atm -1 for C2H6 and C3H8, respectively. When divided by the soap concentration of these solutions (0.8 m), these slopes take on the significance of being the change in gas solubility, ASol, accompanying the conversion of a mole of octanoate ion into octanoic acid. They are found to be ASol = 9.3 × 1 0 - 3 / 0 . 8 0 = 120 × 10 -4 atm -1 and ASol = 47 X 1 0 - 3 / 0 . 8 0 = 5 9 0 X 10 . 4 a t m -1 for ethane and propane, respectively. Since each mole of octanoic acid is produced at the expense of a mole of octanoate ion, this change in gas solubility, ASol, is the net change resulting from the loss of solubility attributed to micellized octanoate ions and a concomitant gain in solubility due to the octanoic acid formed; i.e., ASol = ASol (acid) - zXSol (anion). The molar solubility due to micellized octanoate ions, 2×Sol (anion), can be equated with the mole fraction gas solubility determined at pH 12, which is shown in Table II to be 49 X 10 -4 and 100 × 10 .4 for ethane and propane, respectively. It follows that the
14
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Mole NoOH/Mole H00C-CTHI5 F[o. 6. Phase and pH characteristics of 0.8 M solution of sodium octanoate at 25°C. 2@, two phase region; l~b, clear isotropic solution.
GAS SOLUBILITY increases in the solubilities o f ethane and propane attributed to the formation o f octanoic acid in the micelle at 1 arm are ASol (acid) = 120 X 10 -4 + 49 X 10 -4 = 169 X 10 -4 for ethane; and ASol (acid) = 590 X 10 -4 + 100 X 10 -4 = 690 X 10 -4 for propane. Although the resulting values are surprisingly close to those at 1 atm, solubilities measured for ethane and propane in bulk octanoic acid at 2 5 ° C are (as seen in Table II) 190 X 10 -4 and 700 X 10 -4, respectively. The close correspondence observed is probably little m o r e than coincidence since, barring the highly improbable situation in which the octanoic acid undergoes phase separation within the micelle, there is no reason to expect the observed increment in solubility to equal the bulk solubility. CONCLUSION Gas solubilities for four gases, 02, CH4, C2H6, and C3H8, in micellar solutions o f sodium octanoate show a p r o n o u n c e d sensitivity to p H as the p H o f the solutions approaches a value o f p H 7. The origin o f this effect is not understood, although the linear relationship observed between enhanced solubilization and degree o f p r o t o n a t i o n o f octanoate ions suggests that solubilized octanoic acid plays a major role in whatever mechanism is operative in these systems. ACKNOWLEDGMENT The authors express appreciation for support provided by the National Science Foundation (NSF) Grant CHE8218288. REFERENCES 1. John, L. M., and McBain, J. W., J. Amer. Oil Chem. Soc. 25, 141 (1948).
449
2. Stainsby, G., and Alexander, A. E., Trans. Faraday Soc. 45, 585 (1949). 3. Lucassen, J., J. Phys. Chem. 70, 1824 (1966). 4. Powney, J., Trans. Faraday Soc. 31, 1510 (1935). 5. Powney, J., and Addison, C. C., Trans. Faraday Soc. 34, 356 (1938). 6. Schwartz, A. M., and Perry, J. W., "Surface Active Agents," Vol. 1, p. 374. Interscience, New York, 1949. 7. Qutubuddin, S., Miller, C. A., and Fort, T., Jr., J. Colloid Interface Sci. 101, 46 (1984). 8. (a) McBain,J. W., and O'Conner, J. J., J. Amer. Chem. Soc. 62, 2855 (1940); (b) McBain, J. W. and O'Conner, J. J., J. Amer. Chem. Soc. 63, 875 (1941); (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., J. Phys. Chem. 67, 2079 (1963); (f) Winters, L. J., and Grunwald, E., J. Amer. Chem. Soc. 87, 4608 (1965); (g) Miller, K. W., Hammond, L., and Porter, E. G., Chem. Phys. Lipids 20, 229 (1977); (h) Matheson, I. B. C., and King, A. D., Jr., J. Colloid Interface Sci. 66, 464 (1978); (i) Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 82, 260 ( 1981); (j) Hoskins, J. C., and King, A. D., Jr., J. Colloidlnterface Sci. 82, 264 (1981); (k) Christian, S. D., Tucker, E. E., and Lane, E. H., J. Colloid Interface Sci. 84, 423 (1981); (1) DellaGuardia, L., and King, A. D., Jr., J. Colloid Interface Sci. 88, 8 (1982). (m) Bolden, P. L., Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 91, 454 (1983). (n) Ben-Naim, A., and Wilf, J., J. Solution Chem. 12, 671 (1983). (o) Ownby, D. W., and King, A. D., Jr., J. Colloid Interface Sci. 101, 271 (1984). (p) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 106, 186 (1985). (q) Flanagan, H. L., Bolden, P. L., and King, A. D., Jr., J. Colloid Interface Sci. 109, 243 (1986); (r) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 112, 387 (1986). 9. Vikingstad,E., Skauge, A., and Hoiland, H., J. Colloid Interface Sci. 66, 240 (1978). 10. Ralston, A. W., and Hoerr, C. W., J. Org. Chem. 7, 546 (1942). 11. Eggenberger,D. N., Broome, F. K., Ralston, A. W., and Harwood, H. J., J. Org. Chem. 14, 1108 (1949).
A. S H W I K H A T ,
AND A. D. K I N G , JR.
Department of Chemistry, School of Chemical Sciences, University of Georgia, Athens, Georgia 30602 Received February 4, 1986; accepted May 6, 1986 Measurements have been made to determine the solubilities of oxygen, methane, ethane, and propane in aqueous solutions of sodium octanoate at 25°C and at pH values ranging from 7 to 13. The solubility of each gas obeys Henry's law at all surfactant concentrations and pH values. The solubilities measured for each gas are effectivelyindependent of sodium octanoate concentration as well as pH at concentrations well below the CMC, and, within experimental error, are the same as the respective gas solubilities in pure water at 25°C. Above the CMC, the solubility of each gas increases with surfactant concentration indicating micellar solubilization. The degree of solubilization is greatest for propane and decreases in the order: propane > ethane > methane > oxygen. However, unlike the situation at low surfactant concentrations, the individual gas solubilities vary with pH at higher surfactant concentrations. The degree of solubilization is found to be quite constant for each gas at pH values greater than 9; and intramicellarsolubilities derived from gas solubility data taken at pH 12 are identical to those determined for a similar surfactant having a seven carbon alkyl chain, sodium 1-heptane sulfonate. However, the extent of solubilization for each gas increases rapidly as the solution pH decreases from 9 to 7, suggesting that mixed micelles composed of octanoic acid and sodium octanoate are more effective solubilizing agents than those of the pure soap. © 1987AcademicPress,Inc. INTRODUCTION Surface active salts o f long-chain c a r b o x y l i c acids are generally designated as soaps. As a class, soaps differ f r o m m o s t o t h e r a n i o n i c surfactants in t h a t the m o d e r a t e l y basic carb o x y l a t e h e a d g r o u p s are subject to hydrolysis equilibria (1-3). As a result, the h y d r o p h i l i c c h a r a c t e r o f these surfactants is sensitive to the acidity o f t h e a q u e o u s e n v i r o n m e n t a n d p r o p e r t i e s such as surface t e n s i o n (4), interfacial t e n s i o n (5), d e t e r g e n c y (6), a n d e m u l sifying characteristics (7) o f soap s o l u t i o n s all exhibit a strong d e p e n d e n c e on pH. This p a p e r e x a m i n e s the role that p H plays in d e t e r m i n i n g the extent o f solubilization o f gases.1 It reports the results o f a series o f solubility m e a s u r e m e n t s at 25 °C involving four gases with widely differing p r o p e r t i e s ( 0 2 , CH4, C2H6, a n d l Reference (8) constitutes a reasonably complete bibliography on the subject of gas solubilization in micellar solutions.
C3H8) dissolved in a q u e o u s solutions o f a typical soap, s o d i u m octanoate, at p H values ranging b e t w e e n 7 a n d 13. EXPERIMENTAL T h e m e t h o d used to d e t e r m i n e gas solubilities has b e e n previously described in detail (8m). It differs from other m e t h o d s c o m m o n l y used to d e t e r m i n e gas solubility in t h a t it involves m e a s u r i n g the v o l u m e o f gas released f r o m solutions s a t u r a t e d at elevated pressures. T h e p r o c e d u r e consists o f three steps: (i) the surfactant s o l u t i o n is allowed to equilibrate with the gas o f interest at a n elevated pressure in a t h e r m o s t a t t e d b o m b e q u i p p e d with a m a g n e t i c a l l y d r i v e n stirrer; (ii) the stirrer is s t o p p e d a n d the solution is allowed to b e c o m e still, w h e r e u p o n the pressure is released; a n d (iii) after a short delay ( ' ~ 30 s) to allow t h e r m a l e q u i l i b r i u m to b e reestablished following the release o f gas, the m a g n e t i c stirrer is again ac-
443 0021-9797/87 $3.00 Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
Copyright © 1987 by Academic Press, Inc. All rights of reproduction in any form reserved.
444
PRAPAITRAKUL, SHWIKHAT, AND KING
tivated causing the dissolved gas to effervesce from the formerly supersaturated solution. The volume of gas released is measured manometrically at ambient pressure. In this manner one determines the number of moles of gas released, An, as the solution proceeds isothermally from what was previously an equilibrium state at the upper pressure to a new equilibrium state at ambient pressure. The difference in pressure between these two states, AP, is numerically equal to the gauge pressure recorded during the initial equilibration step (i). The pressures employed in these experiments were such that the amounts of dissolved gas were 30 vol% or less, measured at ambient pressure. The series of steps, (i)(iii), are repeated one or more times at different pressures in order to establish that the ratios 2xn/AP are independent of pressure. This was found to be the case in every instance so that the averaged values of An/&P listed in Table I can be taken to represent gas solubilities in the respective solutions at a partial pressure of 1 atm with Henry's law being satisfied for each system.
The sodium octanoate used in all experiments was produced in situ by titrating aqueous suspensions containing known amounts of octanoic acid (Eastman Cat. No. 665, Lot Nos. 691 and 692) with a concentrated solution of sodium hydroxide (Baker analyzed reagent grade) to the desired pH. The pH measurements were performed with a Fisher Acumet Model 825 MP pH meter using standard buffers and two point calibration to standardize the instrument. The gases used were CP grade or the equivalent, having quoted purities of 99.0% or better for the hydrocarbon gases which were purchased from Matheson and 99.6% for the oxygen which was obtained from Selox Corp. Doubly distilled water was used to prepare all solutions. RESULTS AND DISCUSSION
Gas solubilities determined as a function of sodium octanoate concentration at pH 7.5 and 12 are listed in Table I. These data are shown plotted as a function of soap concentration in Figs. 1-4. Inspection of the data shown in Figs.
TABLE I Gas Solubilities in Aqueous Solutions of Sodium Octanoate at pH 12 and pH 7.5 Expressed as Moles Gas per Atmosphere in 1000 g H20 at 25°C Gas solubility (m) × 103 ,o Sodium octanoate concentration (m)
02
C1J~
C2I%
C3Hs
0 0.2 0.3 0.4 0.5 0.6 0.8 1.0 1.2
1.41 c 1.35 (1.37) - - (1.33) 1.34 (1.38) 1.28 (--) 1.33 (1.49) 1.39 (1.47) 1.44 (1.58) 1.57 (1.64)
1.55 c 1.46 (1.61) - - (1.52) 1.50 (1.53) 1.47 (--) 1.56 (1.88) 1.82 (2.07) 1.95 (2.31)
1.76 c 1.83 (1.87) - - (1.87) 2.01 (2.00) 2.08 (--) 2.48 (3.32) 3.49 (4.71) 4.48 (6.17)
1.42 d 1.4 (1.4) - - (1.5) 1.6 (2.4) 2.1 ( - - ) 3.0 (5.1) 5.0 (8.6) 7.2 (11.5)
a First value listed for solubility is for solution having pH 12. Second value in parenthesis is for solution having pH 7.5. bAverage errors of solubilities measured in these experiments for 02, CH4, C2H6, and C3Hs are +1 × 10-5, +2 × 10-5, +4 × 10-5, and _1 × 10-4, respectively. c Data taken from Ref. (8m). a Datum taken from Ref. (8p).
Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
GAS SOLUBILITY 2.0
445 0 x
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FIG. 1. Moles of oxygen absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12. 1-4 reveals that within experimental error, the solubility o f each gas remains the same as that in pure water until the concentration o f sod i u m octanoate reaches approximately 0.4 m at which point the solubility o f each gas begins to increase with surfactant concentration. The C M C of sodium octanoate has been determ i n e d to be 0.405 m at 2 5 ° C (9) suggesting that, as found previously with other surfactants (8), the e n h a n c e d solubility observed for each gas above 0.4 rn is the result o f solubilization by micelles present at soap concentrations in excess o f the CMC. A c o m p a r i s o n o f the data shows that above the CMC, the degree to which the solubility is e n h a n c e d by the soap
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FIG. 3. Moles of ethane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
micelles is greatest for propane and decreases in the order o f decreasing boiling point for the individual gases: C3H8 (231 K) > C2H6 (184 K) > CH4 (109 K) > 02 (90 K). This trend with boiling point is quite general and is c o m m o n l y observed with gases solubilized in a wide variety o f surfactants (8r). It is the behavior expected for systems in which dispersion forces constitute the d o m i n a n t m o d e o f interaction between solute and solvent molecules. A n o t h e r feature o f interest in the data o f Figs. 1-4 is the fact that the degree to which each gas is solubilized in sodium octanoate micelles is a function of the p H o f the solution, with the micelles absorbing m o r e gas at the lower pH. This unusual feature is not c o m m o n to anionic surfactants having strong acid head groups 2 and is u n d o u b t e d l y related to the p H dependent hydrolysis o f the mildly basic carboxylate head group. In general, the paucity o f publications on the topic o f gas solubilization (Ref. (8) contains
I 1,0
SODIUMOCTANOATECONC. (mol kg-1]
FIG. 2. Moles of methane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
2 The insensitivity toward pH exhibited by anionic surfactants is illustrated by the fact that the solubility of propane in 0.5 rn sodium octylsulfate at 25°C remains constant, having a value of(7.5 + 0.1) X 10 -3 mole C3H8kg-1 atm -~, at pH values ranging between 2.5 and 12. Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
446
PRAPAITRAKUL, SHWIKHAT, AND KING ]2.0
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SODIUMOCTANOATE CONC, (m01 kg -I) FIO. 4. Moles of propane absorbed per atmosphere in 1000 g of H20 with added sodium octanoate shown as a function of soap concentration at 25°C. O, pH 7.5; ®, pH 12.
only 18 citations) precludes frequent comparisons between data obtained in different laboratories. It is therefore very fortunate that the solubility data obtained here with the hydrocarbon gases, methane, ethane, and propane, at pH 12 can be compared directly with results obtained previously by Ben-Naim and Wilf
(8n) for these same gases under identical conditions using a different experimental method. Within experimental error, perfect agreement is found between solubilities measured here for ethane and propane and those measured by Wilf and Ben-Naim. However, in the case of methane, a significant discrepancy exists between the results obtained in the two different laboratories. Below the CMC, the solabilities determined here for CH4 are as much as 13% greater than those published by BenNaim and Wilf. As a consequence, no evidence is found in Fig. 2 for the puzzling minimum in solubility observed by Ben-Naim and Wilf for methane above the CMC of sodium octanoate. The rising portions of the data shown in Figs. 1-4 extend in a linear fashion to quite high soap concentrations, suggesting that the sorptive capacity of the individual micelles remains constant, even at concentrations as great as 1.2 m. If it is assumed that soap monomer concentration remains effectively constant at concentrations greater than the CMC, then the slopes of the rising portions of Figs. 1-4 can be equated to the mole ratio ofsolubilized gas to micellized soap. Since the solutions are all quite dilute with respect to dissolved gas, these mole ratios may in turn be taken to represent mole fraction solubilities of the various gases solubilized within the sodium octanoate micelles at 1 atm, X~. The micellar gas solubilities obtained in this manner are listed in Table II along with values obtained previously for
TABLE II Micellar Solubilities of Gases in Sodium Octanoate at 25°C and 1 atm (Mole Fraction × l04)
Gas
0 2
CH4 C2H6 C3Hs
Sodium octanoate a o H 7.5
Sodium octanoat¢ a p H 12.0
Sodium l-heptane sulfonate b
Octanoic acid c
3 13 70 150
4 10 49 100
4 10 45 102
--190 700
"Estimated error: 02, +1 × 10-4; CH4, +1 × 10-4; C2H6, _+1 × 10-4; C3H8,_+4 × 10-4. b Data from Ref. (80. c Estimated error: C2H6, _+4 × 10-4; C3H8, _+10 X 10-4. Journal of Colloid and Interface Science, Vol. I 15, No. 2, February 1987
447
GAS SOLUBILITY
the closely related surfactant, sodium 1-heptane sulfonate. Previous studies in this laboratory have shown that alkyl-chain length is the dominant factor in determining the sorptive capacity of miceUes derived from ionic surfactants (8m, 8r). Therefore, since hydrolysis effects are expected to be negligible at pH 12, it is reasonable to expect the micellar solubilities of 02, C H 4 , C 2 H 6 , and C3H8 to have the same values in sodium octanoate at pH 12 as determined previously with sodium 1heptane sulfonate (8r). A comparison of the micellar solubilities listed in the middle columns of Table II shows this to be the case. The micellar solubilities measured at pH 7.5 equal or exceed those measured at pH 12 and, with the exception ofO2 for which any change in sorptive capacity is masked by experimental error, it appears that reducing the pH from 12 to 7.5 results in approximately a 40% increase in micellar gas solubility. In terms of percentage, this observed enhancement is considerably greater than the small decrease in CMC caused by the reduction in pH. Clearly, the enhanced gas solubility accompanying the reduction in pH cannot be explained in terms of a simple reapportionment of monomer anions into micelles as might be expected if a simple electrolyte were added to the solution. Rather, some other process producing a much greater increase in the sorptive capacity of the micellized octanoate ions must be involved. In order to more carefully examine the effect o f p H on gas solubility, a series of experiments was performed in which the solubilities of ethane and propane were determined in 0.8 m solutions of sodium octanoate to which small amounts of sodium hydroxide or hydrochloric acid had been added in order to adjust the pH of the solutions to values falling within the range 7-13. The lower limit of these experiments was dictated by the fact that below pH 7 a precipitate composed of a mixture of octanoic acid and sodium octanoate forms, causing the solution to become a gel. The resuits of these experiments are listed in Table III. These data, presented graphically in Fig. 5, reveal that the solubilities of ethane and
TABLE III The Solubility of Ethane and Propane in 0.8 m Sodium Octanoate at 25°C as a Function o f p H C2I~
C3Hs
pH
Solubility (m) × 10s °
7.07 7.20 7.41 7.81 8.11 8.89 9.28 10.16 11.54 12.20 12.52
5.02 4.65 4.01 3.67 3.42 3.11 3.38 3.36 3.27 3.27 3.41
pH
Solubility (m) × 103a
7.17 7.40 7.55 7.80 8.20 9.68 11.60 13.15
11.4 11.8 9.9 6.9 5.1 4.6 4.4 4.7
a Average errors in solubility found in footnote b of Table I.
propane remain quite constant as the pH falls from 13 to 9 and then increase rapidly as the pH of the solution drops from pH 9 to 7. The phase behavior and pH characteristics of sodium octanoate/octanoic acid mixtures having a formal concentration of 0.8 M are summarized graphically in Fig. 6. 3 These data show that octanoic acid is 100% converted into the salt form, i.e., pure soap is present, over the pH range of 9-13 for which low but constant gas solubilities are found for both ethane and propane in Fig. 5. Therefore, one concludes that the gas solubilities recorded at pH 12 for 02, CH4, C2H6, and C3H8 in Table I and Figs. 1-4 reflect micellar solubilization by pure sodium octanoate. However, it is also clear from Fig. 6 that a significant conversion of octanoate anion to the acid form occurs as the pH falls from 9 to 7. The solubility of octanoic acid in water is quite low and is esti3 The data used to construct Fig. 6 were obtained using solutions/suspensions formed by adding known a m o u n t s of reagent grade sodium hydroxide to preweighed samples of octanoic acid. Doubly distilled water was used to dilute each solution/suspension to a formal concentration of 0.8 M with respect to the acid. Journal of Colloid and Interface Science, Vol. 115,No. 2, February 1987
448
PRAPAITRAKUL, SHWIKHAT, AND KING
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pH FIG. 5. Solubilities of ethane and propane in 0.8 m solutions of sodium octanoate at 25°C shown as a function of pH. Inset: Solubilities of ethane and propane shown as a function of' micelle composition. ©, C2H 6; O, C3H8.
mated to be only 5.1 × 10 -3 M a t 2 5 ° C . 4 Since the solutions remain transparent at these pH values, one must assume that octanoic acid is solubilized within micelles derived from the remaining sodium octanoate. Therefore, it is reasonable to conclude that the increase in gas solubility observed with failing pH is the result of an internal conversion of sodium octanoate to octanoic acid resulting in the formation of mixed micelles having enhanced sorptive capacities toward these gases. With this in mind, one can use the pHcomposition data of Fig. 6 to recast the solubilities of ethane and propane as a function of micellar composition as shown in the inset of Fig. 5. Although there is considerable scatter among the data, as might be expected from 4 The solubility of octanoic acid at 25 °C is interpolated from solubility data taken from Refs. (10, 11). Journal of Colloid and Interface Science, Vol. 115, No. 2, February 1987
accumulated errors in the calculations, the solubilities of both gases clearly exhibit a linear dependence on octanoic acid concentration. The slopes of the individual lines which have been fitted to the data according to the leastsquares criterion are 9.3 X 10 -3 and 47.3 X 10 -3 mole kg -1 atm -1 for C2H6 and C3H8, respectively. When divided by the soap concentration of these solutions (0.8 m), these slopes take on the significance of being the change in gas solubility, ASol, accompanying the conversion of a mole of octanoate ion into octanoic acid. They are found to be ASol = 9.3 × 1 0 - 3 / 0 . 8 0 = 120 × 10 -4 atm -1 and ASol = 47 X 1 0 - 3 / 0 . 8 0 = 5 9 0 X 10 . 4 a t m -1 for ethane and propane, respectively. Since each mole of octanoic acid is produced at the expense of a mole of octanoate ion, this change in gas solubility, ASol, is the net change resulting from the loss of solubility attributed to micellized octanoate ions and a concomitant gain in solubility due to the octanoic acid formed; i.e., ASol = ASol (acid) - zXSol (anion). The molar solubility due to micellized octanoate ions, 2×Sol (anion), can be equated with the mole fraction gas solubility determined at pH 12, which is shown in Table II to be 49 X 10 -4 and 100 × 10 .4 for ethane and propane, respectively. It follows that the
14
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8
6
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~-~-2~
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.[
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i I i i i I i i i i i i i 0.25 0.50 0.75 1.00 1,25 1.50 I,'75
Mole NoOH/Mole H00C-CTHI5 F[o. 6. Phase and pH characteristics of 0.8 M solution of sodium octanoate at 25°C. 2@, two phase region; l~b, clear isotropic solution.
GAS SOLUBILITY increases in the solubilities o f ethane and propane attributed to the formation o f octanoic acid in the micelle at 1 arm are ASol (acid) = 120 X 10 -4 + 49 X 10 -4 = 169 X 10 -4 for ethane; and ASol (acid) = 590 X 10 -4 + 100 X 10 -4 = 690 X 10 -4 for propane. Although the resulting values are surprisingly close to those at 1 atm, solubilities measured for ethane and propane in bulk octanoic acid at 2 5 ° C are (as seen in Table II) 190 X 10 -4 and 700 X 10 -4, respectively. The close correspondence observed is probably little m o r e than coincidence since, barring the highly improbable situation in which the octanoic acid undergoes phase separation within the micelle, there is no reason to expect the observed increment in solubility to equal the bulk solubility. CONCLUSION Gas solubilities for four gases, 02, CH4, C2H6, and C3H8, in micellar solutions o f sodium octanoate show a p r o n o u n c e d sensitivity to p H as the p H o f the solutions approaches a value o f p H 7. The origin o f this effect is not understood, although the linear relationship observed between enhanced solubilization and degree o f p r o t o n a t i o n o f octanoate ions suggests that solubilized octanoic acid plays a major role in whatever mechanism is operative in these systems. ACKNOWLEDGMENT The authors express appreciation for support provided by the National Science Foundation (NSF) Grant CHE8218288. REFERENCES 1. John, L. M., and McBain, J. W., J. Amer. Oil Chem. Soc. 25, 141 (1948).
449
2. Stainsby, G., and Alexander, A. E., Trans. Faraday Soc. 45, 585 (1949). 3. Lucassen, J., J. Phys. Chem. 70, 1824 (1966). 4. Powney, J., Trans. Faraday Soc. 31, 1510 (1935). 5. Powney, J., and Addison, C. C., Trans. Faraday Soc. 34, 356 (1938). 6. Schwartz, A. M., and Perry, J. W., "Surface Active Agents," Vol. 1, p. 374. Interscience, New York, 1949. 7. Qutubuddin, S., Miller, C. A., and Fort, T., Jr., J. Colloid Interface Sci. 101, 46 (1984). 8. (a) McBain,J. W., and O'Conner, J. J., J. Amer. Chem. Soc. 62, 2855 (1940); (b) McBain, J. W. and O'Conner, J. J., J. Amer. Chem. Soc. 63, 875 (1941); (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., J. Phys. Chem. 67, 2079 (1963); (f) Winters, L. J., and Grunwald, E., J. Amer. Chem. Soc. 87, 4608 (1965); (g) Miller, K. W., Hammond, L., and Porter, E. G., Chem. Phys. Lipids 20, 229 (1977); (h) Matheson, I. B. C., and King, A. D., Jr., J. Colloid Interface Sci. 66, 464 (1978); (i) Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 82, 260 ( 1981); (j) Hoskins, J. C., and King, A. D., Jr., J. Colloidlnterface Sci. 82, 264 (1981); (k) Christian, S. D., Tucker, E. E., and Lane, E. H., J. Colloid Interface Sci. 84, 423 (1981); (1) DellaGuardia, L., and King, A. D., Jr., J. Colloid Interface Sci. 88, 8 (1982). (m) Bolden, P. L., Hoskins, J. C., and King, A. D., Jr., J. Colloid Interface Sci. 91, 454 (1983). (n) Ben-Naim, A., and Wilf, J., J. Solution Chem. 12, 671 (1983). (o) Ownby, D. W., and King, A. D., Jr., J. Colloid Interface Sci. 101, 271 (1984). (p) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 106, 186 (1985). (q) Flanagan, H. L., Bolden, P. L., and King, A. D., Jr., J. Colloid Interface Sci. 109, 243 (1986); (r) Prapaitrakul, W., and King, A. D., Jr., J. Colloid Interface Sci. 112, 387 (1986). 9. Vikingstad,E., Skauge, A., and Hoiland, H., J. Colloid Interface Sci. 66, 240 (1978). 10. Ralston, A. W., and Hoerr, C. W., J. Org. Chem. 7, 546 (1942). 11. Eggenberger,D. N., Broome, F. K., Ralston, A. W., and Harwood, H. J., J. Org. Chem. 14, 1108 (1949).