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W. R. GILLAP,N. D. WEINER,AND M. GIBALDI
sively larger deviations occurred for sodium dodecyl and tetradecyl sulfates. It is quite likely that the results observed in the present study reflect the difference in dimerization potential of SDS and SDDS. Mukerjee, et ~ l . , ~have ’ also suggested that certain discrepancies in the results of their study of hydrophobic bonding in micelle formation may be due to dimerization of long-chain ions in submicellar concentrations. Since dimerization depends primarily on the hydrophobic interactions between hydrocarbon chains and since SDDS exhibits a much greater tendency to dimerize than SDS, the entropy for the bulk monomer-adsorbed ion equilibrium for SDDS would be
greater than the apparent measured value. The present determination may reflect an equilibrium of the type bulk monomer
21
bulk dimer
Jl”
adsorbed species
Hence, the difference in the driving force to adsorption between the SDS and SDDS systems must be enthalpic. However, when considering the microscopic process of bulk monomer -+ surface monomer, the difference in driving force between the two surface-active agents may still be predominantly entropic on origin.
Ideal Behavior of Sodium Alkyl Sulfates at Various Interfaces. Thermodynamics of Adsorption at the Oil-Water Interface by William R. Gillap, Norman D. Weiner, and Milo Gibaldi College of Pharmaceutical Sciences, Columbia University, New York, New York 10026 (Received December 14, 1967)
Adsorption of sodium decyl sulfate and sodium dodecyl sulfate has been studied at very high areas per mole,> ,. cule at the hydrocarbon-0.1 M NaCl interface for a homologous series of n-alkanes (CC Ideal gaseous behavior was observed in all systems studied. The chain length of the oil had no apparent effect on adsorption in the ideal region. The standard thermodynamic functions for adsorption at the dodecane-0.1 M NaCl interface were determined and evaluated. Based on comparisons of the data obtained for both air-water and hydrocarbon-water interfaces and on evaluation of the results in terms of the major forces involved in the adsorption process, it is suggested that sodium dodecyl sulfate dimers, which are present in the aqueous bulk, dissociate at the oil-water interface and that the monomer is the exclusive adsorbed species. The importance of hydrophobic bonding and entropy as a driving force to adsorption at the oil-water interface is discussed.
Introduction Few data are available regarding the energetics of adsorption at the oil-water interface. Haydon and Taylor’ calculated the free energy of adsorption per methylene group, from the “initial” slopes of interfacial tension us. concentration curves for solutions of sodium alkyl sulfates (C8,Cl0,and Clz) in 0.1 M NaC1, to be 758 cal/mol of CH2. A review of their data, however, indicates that the bulk concentrations employed exceeded those required to maintain ideality at the interface.2 Davies8has reported an energy of adsorption per methylene group for the petroleum ether (bp >120°)-water interface of 811 cal/mol of CH2. This was calculated by comparing rates of desorption (by measuring changes in interfacial tension at constant area) of monolayers of cetyltrimethylamnionium bromide injected into a spreading solution at the 0.01 N HC1-petroleum ether The Journal of Physical Chemistry
interface and rates of desorption of monolayers of octadecyltrimethylammonium chloride at the air-0.01 N HC1 surface. Since rates of desorption were measured (which could differ significantly from rates of adsorption, particularly as a function of chain length), two different interfaces were employed, and the difference in the rates of desorption of the spreading solvent (propyl alcohol) at the two interfaces was not considered: the significance of this value is doubtful. Cassie and Palmer14using the data of Powney and Addison16report (1) D. A. Haydon and F. H. Taylor, Trans. Faraday SOC.,58, 1233 (1962). (2) W. R. Gillap, N. D. Weiner, and M. Gibaldi, J . Colloid Interfac. Sci., 25, 441 (1967). (3) J. T. Davies, Trans. Faraday Soc., 48, 1052 (1952). (4) A. B. D. Cassie and R. C. Palmer, ibid., 37, 156 (1941). (6) J. Powney and C. C. Addison, ibid., 33, 1243 (1937).
THERMODYNAMICS OF ADSORPTION AT
THE
2223
OIL-WATERINTERFACE
Interfacial tension was determined by the Wilhelmy a heat of adsorption per methylene group of 0.85 IcT. plate method. lo The hydrocarbon-sodium decyl sulThe latter estimate, however, is based on heats of fate and -sodium dodecyl sulfate solution interfacialmicellization rather than on heats of adsorption. tension values were unaffected by mutual saturation of Furthermore, the purity of the surfactants used by the aqueous and hydrocarbon phases. It was, therePowney and Addison6 has been questioned because of fore, considered more important to ensure the purity of the presence of minima in the interfacial tension os. conthe hydrocarbons by passing them through the silica centration curves. Therefore, although these data are as indicative of free gel-alumina column immediately prior to use, rather quoted in texts and other than routinely using preequilibrated systems. All soluenergies or heats of adsorption at the oil-water intertions were run at least in duplicate. Interfacial tension face, they do not represent precise thermodynamic was continually noted until constant values were obvalues and may serve only as approximations. The importance of water structure and hydrophobic served. Usually 15 min was required for equilibration. Interfacial tension data were obtained for sodium bonding in the association and adsorption of surfaceactive agents has been discussed in the previous a r t i ~ l e . ~ dodecyl sulfate at the hexane-, dodecane-, tetradecane-, and hexadecane-0.1 M NaCl interfaces at 25 f 0.1”. Gillap, et al.,l0 have also considered the role of water Data for SDDS at the dodecane-0.1 M NaCl interface structure and hydrophobic bonding in the phenomenon were also obtained at 35 and 45 f 0.1”. of hydrocarbon-water interfacial tension. I t was proInterfacial tension data were obtained for sodium posed that hydrocarbon molecules have a tendency to interact and to stabilize water clusters at the interface, decyl sulfate at the hexane-, octane-, dodecane-, and as well as the fact that they have a tendency toward selfhexadecane-0.1 M NaCl interfaces at 25 0.1’ and at the dodecane-0.1 M NaCl interface at 35 and 45 f 0.1’. interaction and disruption of water clusters. The latter effect becomes more significant with increasing chain Results length. In view of these observations, hydrophobic In all systems studied, an ideal region was observed. bonding forces for both the surface-active agent and Representative curves are shown in Figures 1 and 2. hydrocarbon interactions must be considered in adsorpThe extremely low concentration of surfactant required tion processes at the oil-water interface. No reports to approach the ideal region should be noted. A comconcerning such competitive interactions may be found parison of the concentrations required for ideality at the in the available literature. It was the purpose of this investigation to examine the effect o€ a homologous series of hydrocarbons on the adsorption of sodium alkyl sulfates at the oil-water Table I : Slopes ( 0 1 Values) of Interfacial Tension vs. interface, particularly at low, ideal concentration levels Molar Concentration Plots for Sodium Decyl Sulfate and Sodium Dodecyl Sulfate at 25’ of the surfactants, and to determine and to evaluate the standard thermodynamic functions of the adsorption process. Of equal importance was the comparison of SDDS the interfacial and surface-adsorption properties of ( X 108) sodium alkyl sulfates and the further elucidation of the Hexane ... 3.29 role of hydrophobic bonding in adsorption at these interOctane 7.75 3.24 Decane ... 3.25 faces.
*
Experimental Section Fractionally distilled, demineralized water was used in all experiments. Deionized water was fractionally distilled in the presence of basic potassium permanganate. The hydrocarbons, hexane, octane, decane, dodecane, tetradecane, and hexadecane, were purchased from Eastman Organic Chemicals. The oils were purifed according to the method of Weiner, et al.,” and repeatedly passed through an alumina-silica gel column until constant interfacial tension values against water were obtained.1° The surface tension of each hydrocarbon after purification agreed with reported data.8s12 The surface-active agents, sodium decyl sulfate (SDS) and sodium dodecyl sulfate (SDDS), were synthesized and purified as described in the previous article.$ Aqueous solutions were freshly prepared before each experiment.
Dodecane Tetradecane Hexadecane a
8.00
...
7.73 7.83“
3.18 3.18 3.24 3.23“
Average value.
(6) J. N. Phillips, Trans. Faraday SOC.,51, 561 (1955), (7) D.Shinoda, “Colloidal Surfactants,” Academic Press Inc., New York, N. Y., 1963,p 39. (8) L. I. Osipow, “Surface Chemistry, Theory and Industrial Applications,” Reinhold Publishing Corp., New York, N. Y., 1962, pp 184, 238. (9) W.R.Gillap, N. D. Weiner, and M. Gibaldi, J. Phus. Chem., 7 2 , 2218 (1968). (10) W. R. Gillap, N. D. Weiner, and M. Gibaldi, J . Amer. Oil Chemists Soc., 44, 71 (1967). (11) N. D. Weiner, H. C. Parreira, and G. Zogmfi, J . Pharm. Sci., 55, 187 (1966). (12) “International Critical Tables,” Vol. 4, McGcaw-Hill Book Co., Inc., New York, N. Y., p 456. Volume 78, Number 6 June 1068
W. R. GILLAP,N. D. WEINER,AND M. GIBALDI
2224
52
f
51
50
49
I
I
5
10 WAR
I I5
I 20
d 15
CONCENTPATION x 107
Figure 1. Interfacial tension (y) vs. molar concentration of sodium decyl sulfate a t various interfaces a t 25': 0, hexadecane-0.1 M NaCI; 6,dodecane-0.1 M NaC1; 0, octane-0.1 M NaC1. Figure 3. Surface and interfacial pressure ( T ) os. molar concentration of sodium decyl sulfate a t various interfaces a t 25": 0, octane-0.1 M NaCl; 8, hexadecane-0.1 M NaCl; 0 , air-0.1 M NaCl.
52
50
08
46
44 M L A R CONCENTRATION
X
10'
Figure 2. Interfacial tension ( 7 ) vs. molar concentration of sodium dodecyl sulfate a t various interfaces a t 25': 0, tetradecane-0.1 M NaC1; e, decane-0.1 M NaC1; 0, octane-0.1 M NaCl; 0 , hexane-0.1 M NaCl.
hydrocarbon-water and air-water interfaces is shown in Figures 3 and 4. The data for each surface-active agent were fitwith the a0, where 6 is the interfacial tenequation 6 = -aC sion, C is the concentration of surface-active agent, &, is the interfacial tension of the hydrocarbon-0.1 M NaCl solution, and a is a constant. The value of (*I for each surface-active agent was independent of the chain length of the hydrocarbon oil. This is apparent from the parallel curves shown in Figures 1 and 2 and the data in Table I. The dependency of interfacial tension on the chain length of the hydrocarbon oil at concentrations of alkyl sulfates in excess of the ideal region has been considered previously. l 3 The effect of temperature on the ideal adsorption of
+
The Journal of Physical Chemistry
5.0
10.0
15.0
25.0
20,O
WLAR CONCEKfMTION
X
30.0
35.0
IO6
Figure 4. Surface and interfacial pressure ( T ) vs. molar concentration of sodium dodecyl sulfate a t various interfaces a t 25": 0, hexane-0.1 M NaC1; 8, hexadecane-0.1 M NaC1; 0 , air-0.1 M NaC1.
SDS and SDDS at the dodecane-0.1 M NaC1 interface is shown in Figure 5, These data permitted calculation of the standard thermodynamic functions of adsorption according to the method of Betts and Pethica,I4 as (13) W. R. Gillap, N. D. Weiner, and M. Gibaldi, J. CoZloid Interfac. Sei., in press.
THERMODYNAMICS OF ADSORPTION AT THE OIL-WATERINTERFACE
2225
52
I1
50
49
-
I 45
I 21
46
I . 5.0
I
I .
10.0
15.0 H O ~ A RCONCEHPUTION x 107
1
10.0
I
described in the previous reportsg The slopes of the 6-C plots ( a values) were calculated by the method of least squares and, in each case, linearity was confirmed at the 95% confidence level. Plots of AGO as a function of temperature, used to determine AS" values, are shown in Figure 6. Values of a, as well as the standard thermodynamic functions of adsorption of SDS and SDDS, are presented in Table 11. A comparison of the standard thermodynamic functions of adsorption at the air-water and oil-water interfaces, at 25" for SDS and SDDS, is shown in Table 111.
Table 11: Standard Thermodynamic Functions for the Adsorption of Sodium Decyl and Dodecyl Sulfates a t Various Temperatures a t the Dodecane-0.1 M NaCl Interface Ab8
pound
SDS SDDS
temp, deg
298.2 308.2 318.2 298.2 308.2
318.2
a
8.00 X 6.60 X 5.68 X 3.17 X 2.95 X 2.78 X
106 106 10' 10" 106 106
AGO,
AH',
koal/mol
kaal/mol
-8.00 -8.18 -8.35 -8.83 -9.09 -9.34
-2.9 -2.9 -2.9 -1.2 -1.2 -1.2
Figure 6. Standard free energy of adsorption (AGO) plotted as a function of temperature from interfacial tension data for sodium decyl sulfate (0)and sodium dodecyl sulfate ( 0 ) a t the dodecane-0.1 M NaCl interface.
25-0
Figure 5. Interfacial tension ( 7 ) us. molar concentration of sodium decyl sulfate (upper plot) and sodium dodecyl sulfate (lower plot) in the ideal region (dodecane system) a t 25 (0), 35 (e),and 45" (e).
Corn-
35 TEWKUTURB ('e)
AS0, eu
17 25
These thermodynamic data indicate that the major driving force to adsorption at the hydrocarbon-water interface is entropic rather than enthalpic. As ex-
Table I11 : Comparison of Standard Thermodynamic Functions for Adsorption at the Air-Water and Dodecane-Water Interfaces for Sodium Decyl Sulfate and Sodium Dodecyl Sulfate a t 25'
AGO (air-water), kcal/mol AGO (oil-water), kcal/mol A(AGo), kcal/mol A H o (air-water), kcal/mol A H o (oil-water), kcal/mol A( A H "), kcal/mol A S o (air-water), eu A S " (oil-water), eu A ( A S o ) , eu
SDS
SDDS
-5.92 -8.00 -2.08 -2.0 -2.9 -0.9 13 17 4
-7.31 -8.83 -1.52 -3.3 -1.2 2.1 13 25 12
pected, the negative standard free energy of adsorption of each alkyl sulfate is greater at the oil-water interface than at the air-water interface. Furthermore, the standard entropy of adsorption of each surfactant is greater at the oil-water interface than at the airwater surface. The AS" value of SDDS adsorption a t the dodecane-water interface is significantly larger than the AS" value for SDS adsorption at this interface. An unexpected finding is the large increase in enthalpy for SDDS at the oil-water interface compared to the value observed at the air-water surface. The increased enthalpy, however, is more than offset by the very large increase in entropy.
Discussion The interactions between the hydrocarbon and the alkyl chain of the surface-active agent, which exist even at high areas per adsorbed molecule at the hy(14) J. J. Betta and B. A. Pethioa, Proc. Intern. Congr. Surface Activity, 2nd (London), 1, 152 (1957).
Volume 7'2,Number 6 June 1068
2226
W. R. GILLAP,N.D. WEINER,AND M. GIBALDI
drocarbon-water interface, obviously do not preclude the existence of an ideal region. Ideality at the interface is apparently dependent on the absence of interactions between the adsorbed surfactant molecules. It is postulated that at the oil-water interface in the ideal region, the surfactant molecules extend into the oil phase and are aligned as a result of the induced dipole-induced dipole and other van der Waals interactions between the alkyl chains of the surfactant and oil molecules. Orientation is expected to be constant over the entire ideal region because of the absence of self-interaction of the surfactant molecules. One would expect greatly increased adsorption at the oil-water interface (because of the hydrocarbonsurfactant interaction) compared with the extent of adsorption at the air-water surface at comparable bulk concentrations. Table I V shows a comparison of interfacial excesses of SDS and SDDS at the air-water and oil-water interfaces. The data demonstrate the significantly greater adsorption of each surfactant at the oil-water interface compared with adsorption at the air-water surface.
Table IV : Comparison of Interfacial Excess (r)for Sodium Decyl Sulfate and Sodium Dodecyl Sulfate a t Air-Water and Dodecane-Water Interfaces a t 25' in the Ideal Regions C = 10 x 10-7 M Air
Dodecane-
0 . 1 M NaCl,=
0 . 1 M NaCl,*
Ratio
Compound
X l O l a molecules/cm(
X 1012 molecules/cm*
oil-water: air-water
SDS SDDS
0.6 5.4
21 70
35
a Ratio of interfacial excesses of SDDS : SDS = 9. interfacial excesses of SDDS:SDS = 3.2.
13
b
Ratio of
Determination of the standard thermodynamic functions for adsorption of SDS and SDDS (Tables I1 and 111) shows the expected decrease in AGO with increasing chain length of surfactant. This decrease (which is not as great as the decrease found a t the air-water surface) is the net result of an increase in the AH" of SDDS and an even larger increase in the magnitude of the entropic term (TAso). The expected increase in entropy with increased chain length is observed at the oil-water interface, whereas no difference in the ASo of the alkyl sulfates was found a t the air-water surface. The difference in the behavior of SDS and SDDS a t the two interfaces may be explained from an interpretation of the thermodynamic functions a t the oil-water and air-water interfaces. It has been postulated in the previous articleg that SDDS is present in the bulk in the form of both monomers and The Journal of Physical Chemistry
dimers and that it is adsorbed at the air-water surface in both forms. At the oil-water interface, however, owing to van der Waals adhesional forces, it is probable that the SDDS dimers dissociate at the interface and the monomer is the exclusive adsorbed species. It is likely that each adsorbed monomer is surrounded by neighboring oil molecules as suggested by Hutchinson.I5 The data which support this theory may be summarized in terms of the following major forces. 1. Hydrophobic bonding is present as the monomers of SDS and the monomers and dimers of SDDS leave the bulk, are adsorbed, and lose structured water. This is an entropic contribution which is independent of the nature of the adsorption interface. The standard entropy of adsorption for both alkyl sulfates at the air-0.1 M NaCl interface was 13 eu. It is suggested that this value represent the entropy gain associated with loss of structured water about the hydrophobic chain(s) of the surfactant molecules upon adsorption. 2. Hydrophobic bonding, owing to the loss of structured water surrounding the hydrocarbon oil chains at the interface, is present as the surfactant molecules penetrate the oil phase and interact with the hydrocarbon molecules replacing, in effect, the "icebergs" at the interface. This force is present only at the oil-water interface. The difference in entropy of adsorption of SDS at the oil-water and air-water interfaces is 4 eu. A similar difference in ASo of SDDS at the two interfaces would be expected. However, the difference in AS" of SDDS at -the oil-water and air-water interfaces is 12 eu. This finding supports the possibility of the existence of SDDS dimers which dissociate upon adsorption at, the oil-water interface and significantly increase the entropic contribution to the adsorption process. 3. An enthalpy effect is present that is associated with the separation of the SDDS dimers into monomers a t the interface. This would be a positive enthalpy effect, since the dimer represents a more stable species in the bulk. This enthalpy effect is apparently absent at the air-water interface, since the dimer is probably adsorbed as such. The effect is reflected in the increased AH" value for SDDS at the oil-water interface, as compared to the value at the air-water interface. A similar comparison for SDS reveals a decrease in AH". 4. An enthalpy effect is present that is associated with the separation of oil molecules as the surfactant molecules penetrate the oil phase. 5. An enthalpy effect is present that is associated with adhesional forces between the alkyl chains of the oil and surfactant molecules. Hutchinson16 has suggested that the forces of adhesion are greater than the forces of cohesion, a situation which would account for the penetration of the hydrophobic chain of the surfactant into the oil phase. If this were the case, the (15) E.Hutchinson, J . Colloid rSci., 3 , 235 (1948).
ULTRASONIC ABSORPTION MEASUREMENTS IN GLYCINE, DIGLYCINE, AND TRIQLYCINE negative enthalpy resulting from interaction and penetration would exceed the positive enthalpy resulting from the separation of the oil molecules. The decrease
2227
in AH" of SDS a t the oil-water interface compared with the value a t the air-water surface supports the probability of this situation.
Ultrasonic Absorption Measurements in Aqueous Solutions of
Glycine, Diglycine, and Triglycinel by Gordon G. Hammes and C. Nick Pace2 Depurtment of Chemistry, Cornell University, Ithaca, New York 14860 (Received December 14, 1967)
Ultrasonic absorption and velocity measurements have been made on aqueous solutions of glycine, diglycine, and triglycine at 4" over the frequency range 10-175 MHa. Similar measurements were also made on solutions of diglycine in water and in aqueous urea, guanidine hydrochloride, and NaCl a t 10'. In all cases, a single relaxation process was observed, and the relaxation times were virtually identical. Detailed consideration of various possible mechanisms suggests that the relaxation process is associated with solute-solvent interactions. The solvation apparently involves both the charged groups and the portion of the molecules between the charged groups. Since all of the solutes are mainly in their solvated form, the reciprocal relaxation time (-4 X lo8sec-l) is essentially equal to the rate constant for solvation, which is apparently identical for all three compounds. The differences in the amplitudes of the ultrasonic relaxation are consistent with the interpretation that glycine is more solvated than diglycine, which is more solvated than triglycine. The addition of urea, guanidine hydrochloride, or NaCl results in a decrease of the amplitude of the relaxation effect. However, the relaxation times are not significantly different from those found in water, indicating that the rate constant for solvation in these solutions is similar to that in water.
Introduction Water plays an important and unique role in determining the structure and stability of many biological macromolecules. The presence of water is known to give rise to the hydrophobic forces3 which are of prime importance in stabilizing the native, globular structure of protein^,^ the helical structure of nucleic acids,s and probably the structure of lipids in biological membranes. In addition, water plays an important role in determining the stability of hydrogen bonding in these molecules.6 Studies of ultrasonic absorption in aqueous solutions of a variety of different molecules ranging from polymers such as poly-L-glutamic acid7 and polyethylene glycolS-loto simple molecules such as amines1' and ethersl1rl2have shown that this technique is useful for investigating water-solute interactions. In this work ultrasonic absorption studies of glycine, diglycine, and triglycine in water and in aqueous solutions of urea, guanidine hydrochloride, and NaCl are reported. These compounds serve as simple models for the polypeptide chain of proteins. Aqueous urea and guanidine hydrochloride are well known protein denaturants, and solubility studies have shown that one mechanism of importance in this regard is their ability to increase the solubility of peptide groups.13~14
Experimental Section Diglycine and triglycine, Mann Assayed grade, were obtained from R'lann Research Laboratories Inc. Glycine and guanidine hydrochloride were obtained from the J. T. Baker Chemical Co. The guanidine hydrochloride was purified using the procedure of Nozaki and Tanford.16 Urea, A grade, was obtained (1) This work was supported by a grant from the National Institutes of Health (GM13292), (2) National Institutes of Health Postdoctoral Fellow, 1966-1968. (3) W. Kauzmann, Advan. Protein Chem., 14, 1 (1959). (4) C. Tanford, J . Amer. Chem. SOC., 84, 4240 (1962). (5) T. T. Herskovitz, Biochemistry, 2, 235 (1963). (6) H. A. Scheraga, The Proteins, 1, 478 (1963). (7) J. J. Burke, G. G. Hammes, and T. B. Lewis, J . Chem. Phys., 42, 3520 (1965). (8) G. G. Hammes and T. B. Lewis, J . Phys. Chem., 70, 1610 (1966). (9) G. G. Hammes and P. R. Schimmel, J . Amer. Chem. Soc., 89, 442 (1967). (10) G. G. Hammes and J. C. Swann, Biochemistry, 6, 1591 (1967). (11) J. H. Andreae, P. D. Edmonds, and J. F. McKellar, Acustica, IS, 74 (1965). (12) G. G. Hammes and W. Knoche, J . Chem. Phys., 45, 4041 (1966). (13) D. R. Robinson and W. P. Jencks, J . Amer. Chem. Sac., 87, 2462 (1965). (14) Y . Nozaki and C. Tanford, J . BWZ. Chem., 238, 4074 (1963).
Volume YO, Number 6 June 1968
