J. Phys. Chem. 1980, 84, 2413-2418
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Nature of Large Aggregates in Supercooled Aqueous Solutions of Sodium Dodecyl Sulfate Elias I. Franses,* School of Chemical Engineering, Purdue University, West Lafayette, Indiana 47907
H. Ped Davls, Wilmer G. Miller, and L. E. Scriven
J. Phys. Chem. 1980.84:2413-2418. Downloaded from pubs.acs.org by UNIV OF NEBRASKA-LINCOLN on 08/25/15. For personal use only.
Depnrtments of Chemical Engineering and Materials Science and of Chemistty, Universlty of Minnesota, Minneapolis, Minnesota 55455 (Received: November 19, 1979; In Final Form: May 8, 1980)
Preparations of 2.0 and 5.5 wt % sodium dodecyl sulfate (SDS) in 3.5 w t ’70 (0.6 M) aqueous NaCl are equilibrium micellar solutions above 28 “C, the Krafft point of the surfactant at this salinity. These systems can be supercooled and remain transparent for hours and days. At 25 “C at equilibrium they are biphasic, a hydrated crystal phase and an aqueous salt solution phase containing only 0.Ol2 w t ’70 SDS. Conductimetry and 13CNMR show that these transparent supercooled systems are indeed supersaturated solutions and not microdispersions of the hydrated crystal. The time lag for the onset of nucleation of the crystals depends strongly on stirring details and probably on presence of gas-liquid interface. The big nonequilibrium aggregates present in the supersaturated systlems resemble micelles in conductivity and molecular motion, and are likely to be metastable micelles as is presumed by Mazer, Benedek, and Carey.
Introduction In surfactant systems a great variety of equilibrium and nonequilibrium microstructures have been reported. Micelles are a special class of equilibrium aggregates. Because there is not, yet a universally accepted definition for micelles, it is important to define the term precisely. The definition adopted here agrees with that of the IUPAC,l Hartley,2 and T a n f ~ r d Further .~ justification of the definition and a discussion of other uses of the term micelle are given in ref 4,section 1.1.2. The term micelle Eitands for stable, disjoint, cooperative, closed, equilibrium colloidal aggregates. By “stable” is meant not permanence of the individual aggregates, but constancy of the properties of the entire population of aggregates. By “disjoint” is meant that the aggregates are of limited extent in all three dimensions and remain clearly identifiable in principle even when they are closely packed with one another. By “closed” is meant that a closed connected surface can be constructed such that all the hydrophilic surfactant moieties lie on the same side of the surface and all the lipophilic moieties lie on the other side. Molecular imperfections can of course occur. Disjointedness and closedness are ensured if each aggregate has an inside and an outside, Le., a topological order. The simplest examples are closed globular particles and tubules of genus zero.‘ Cooperativity refers to the association pattern of the aggregates, i.e., the dependence of the aggregation free energy on the aggregate size.6 And by “equilibrium” is meant that the aggregates form spontaneously and reversibly and are limited in size.2 This definition of micelles excludes nonequilibrium microstructures. Moreover, dispersed phase particles and nuclei are excluded because they do not form reversibly, are not limited in size, and have sizes and shapes depending on their thermal and mixing history, thus violating classical criteria for thermodynamic equilibrium. As in the case of phase equilibria of molecular solutions, micelles in metastable equilibrium car1 be defined if nucleation and growth of a precipitating phase can be suppressed. The size and the shape of micelles are central issues of micellization theory. An important question is: how big can micelles be? Available data have been summarized by T a n f ~ r dwho , ~ qualitatively classifies micelles as small if they consist of more than 100 molecules or so and large 0022-36541ao12oa4-24 1 3 ~ I0.OOIO
if they consist of more than 1000 molecules. Large aggregates tend to be more polydisperse in size than small micelles, and this aspect of aggregation equilibria has been modeled by Mukerjeea6Although it has been established that the small micelles generally conform to our defiiition,3 it has not been established that the large aggregates do so. Nevertheless many authors call them micelles and generally use models analogous to those of the small micelles. The small micelles are almost certainly globular on the average, although their size and shape can fluctuate. It is difficult, however, to discriminate experimentally oblate from prolate spheroidal shapes, which seem to be the logical choices, unless the aspect ratio is substantial enough to produce measurable dissymmetry of scattered visible lights7Small micelles of ionic surfactants can grow to giant “micelles” as surfactant concentration is raised, salt concentration is increased, or temperature is lowered. Small micelles of nonionic surfactants can grow as temperature is raised. The mode of growth of small micelles to giant aggregates is still debated.8 There are two leading possibilities: growth of small, globular micelles to long tubular aggregates; and secondary reversible aggregation of small micelles to clusters of micelles, which are in contact but remain identifiable. Available evidence3v9indicates that the so-called large micelles tend to occur closer to the solubility limit of the surfactant than small micelles do: by “solubility limit” is meant the largest amount of surfactant that can be dissolved spontaneously and reversibly. Indeed, very large micelles have been reported in supersaturated solution,9 where they are a t best metastable. For this reason, the thermodynamic stability or metastability has to be carefully established before the detected aggregates can be identified as micelles. In order to discriminate micelles from particles of dispersed second phase, it is essential t o determine the solubility of the surfactant precisely. Moreover, micelles can scatter light similarly to concentration fluctuations near consolute points, and, to avoid confusing the two, the phase diagram must be known. By both quasi-elastic light scattering and conventional light scattering measurements, Mazer et aL9Jodeduce the presence of tubular aggregates of sodium dodecyl sulfate (SDS) in aqueous NaC1. The aggregates are larger the closer the surfactant concentration is to the solubility limit 0 1980 American Chemical Society
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The Journal of Physical Chemistry, Vol. 84, No. 19, 1980
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TABLE I: Solubilities of Sodium Dodecyl Sulfate in Water a t 2 4 "C
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L,
MOLECULAR SOLUTION
L,
MICELLAR SOLUTION
S
HYDRATED SURFACTANT CRYSTALS
Franses et al,
-
( w t % NaCl/ [(wt % NaCl) + ( w t % H,0)] 19.8 9.35 3.60
0.93 0.0
maximum solubility (wt % SDS in total) 0.004a 0.004a
0.020b 0.Ollb~C 0.015a 0.012a'c 7.8d 2 5d
Solubilities were determined by equilibrating SDS crystals with water or brine, allowing excess crystals to settle, and analyzing SDS in supernatant by hyamine titration. The values given are upper limits of solubilities; see text. Samples were prepared by first dissolving surfactant and then adding brine. e Supernatant was centrifuged to remove suspended crystals. Because no crystals were detected, this concentration is a lower limit of solubility. a
1
0 0
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W T % SURFACTANT
Flgure 1. Sketch of the phase diagram of sodium dodecyl sulfate with water or salt water (pseudobinary). The Krafft point depends on salinity; see text. This sketch is consistent with all our observations ( * marks regions of insufficient information).
(Figure l),and they are largest aboue the solubility limit. In order to get transparent systems above the solubility, it is necessary to supercool the solutions or to raise their salt concentration. The aggregates observed above the solubility (0.Ol2 wt %, Table I) are definitely not equilibrium micelles, because the equilibrium state is a hydrated crystal of SDS in equilibrium with a molecular solution of SDS, as we establish below (see Figure 1). Mazer et ala9suggest that the aggregates they observe in supercooled solutions are metastable as long as visible crystallites are absent. By "metastable" they mean that the detected aggregates exist in equilibrium with the monomers and, moreover, that the equilibria between aggregates and monomers follow similar patterns as those of small equilibrium micelles. For the system studied by Mazer et al., SDS in 0.6 M NaCl at 25 "C, Corti and Degiorgio"J2 obtain sizes which follow the same trends with temperature and salinity but are about 30% lower than those of Mazer et ale9The disparity is probably due to the different sample of SDS used.1° The existence of metastable micelles implies intersection of the cmc and solubility curves as indicated in Figure 1. Below the solubility curve micellar solution, in analogy with molecular solution, may be metastable over a range of temperatures and unstable below this temperature range. Little is known about the thermodynamics of metastable micellar solution. The measurements of both Corti and Degiorgio"J2 and Mazer et al. were reproducible to 5% or better. Neither group reports detailed tests of metastable equilibrium, such as thermal cycles and dependence of observed sizes on thermal history and time after supercooling, although Mazer et al. mention that the determined sizes do not depend on direction of temperature change, of course as long as the solution has not phase separated. Important questions remain, however. Are all the aggregates of similar shape and size, or are some of them small micelles and others small microcrystallites of the hydrated crystal phase (Figure l ) ? How are the observed aggregates related to the nuclei of the hydrated crystal? Are the differences between the results of Mazer et al. and Corti and Degiorgio due only to the established differences in the surfactant sample used or to slight variations in preparing the samples as well? To resolve these issues, it was decided to determine the aqueous solubilities of SDS at certain salinities examined by Mazer et al.9 and to investigate the mechanism of the crystallization of the hydrated SDS crystal. The crys-
tallization process in SDS solutions of similar composition to those used in ref 9-12 was followed by direct visual observations, spectroturbidimetry, conductimety, and 13C NMR spectroscopy. The latter two methods were used to complement the light scattering techniques employed previously"'2 by providing molecular mobility probes of the surfactant in the supersaturated solutions. Our results show that the observed giant aggregates resemble well equilibrium micelles in their conductivity and 13C NMR behavior, even though they definitely lack thermodynamic stability.
Materials and Methods Sodium dodecyl sulfate (SDS), specially pure from BDH Chemicals Ltd., Poole, England, had, according to the manufacturer, no more than 1% Clo and C14sulfates. It was used without further purification. The conductivities of aqueous solutions a few days after mixing agreed to better than 5% with literature values,13as did the critical ~ water was drawn micelle concentration ( c ~ c ) . 'Distilled through a Millipore four-stage cartridge system. Its conductivity did not exceed 1.5 X S m-l. The sodium chloride was Certified ACS from Fisher Scientific. Details of the spectrophotometer, low-frequency conductivity bridge, and I3C NMR spectrometer are given e l ~ e w h e r e . ~Temperatures J~ were controlled to f0.5 "C in spectroturbidimetry, f0.1 "C in conductimetry, and f l "C in 13C NMR. Results and Discussion Direct Visual Observations and Spectroturbidimetry. The solubilities of SDS in aqueous sodium chloride at 24 f 1 "C are shown in Table I. Because floating crystallites may inadvertedly have been entrained during sampling the solutions, the reported values at high salinities are actually upper limits of the solubility. At 0.6 M NaCl(w3.6 wt %), the temperatures at which turbid aqueous dispersions of 0.5, 1, 2, and 5.5 w t 9'0 SDS became clear without stirring were determined by increasing the temperature 1 "C at a time and holding it constant for 30 min or more. All clearing points were 28 f 1"C. This narrow temperature range above which the solubility increases more than 500-fold-from 0.01, to more than 5.5%-is the Krafft temperature of SDS at 0.6 M NaC1. This Krafft temperature, or Krafft point, is higher than the value of 25 "C reported by Mazer et al.,9J0who call it critical micelle temperature, or cmt. The difference could be due to the
The Journal of Physical Chemistry, Vol. 84, No. 19, 1980
Aggregates in Supercooled Aqueous Solutions of SDS
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Sample I Sample I (Exp. Repeated)
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J. Phys. Chem. 1980.84:2413-2418. Downloaded from pubs.acs.org by UNIV OF NEBRASKA-LINCOLN on 08/25/15. For personal use only.
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Flgure 3. Time dependence of resistance of 1.99 wt % SDS in 3.50 wt % aqueous NaCI. For R = 148.0 a, the conductivity is 5.569 S m-'.
TIME, min
Flgure 2. Time course of specific absorbance at 400 nm of supersaturated solutions of 1.99 wi % SDS in 3.50 wi % aqueous NaCl at 25 "C. Samples 1 and 2 were heated to clarity and recooled; for sample 2 (experiment repeated), only one point is shown; sample 3 was producedby mlxhg aqueous SDS,2.93 wt %, with a q w NaCI, 10.81 wt % , SDS had been contacting water for 60 days in sample 1, and for 1 day in samples 2 and 3. Measurements were taken until the onset of visible crystallization (indicated by arrows).
difference in purity of the surfactant samples used; we both used SDS from BDH, but they further purified the material they used.10b This difference, however, is not important in comparing our results with those of Mazer et al., because we both consider supercooled systems. Corti and Degiorgio12have estimated by extrapolation from lower salinities the cmc of SDS at 0.6 M NaCl at 25 OC as 0.016 wt 7%. Because the solubility determined here, 0.012 wt 3' % or less, is lower than the calculated cmc, it can be argued that few if any micelles are present in the equilibrium saturated solution. Thus it seems plausible that the solubility below the Krafft temperature is small because no micelles form up to the solubility limit, as suggested by Murray and Hartley16first and recently by Mazer et al.9 and Franses et a1.16 Therefore, the system of 2 wt 90 SDS in 3.5 wt 9O aqueous NaCl at 25 "C is biphasic, a hydrated crystal in stable equilibrium with a solution which contains 0.012 f 0.002 wt 9O SDS. And the solution may be free of micelles a t equilibrium. Supersaturated solutions were prepared either by dispersing 2 wt 9O SDS in brine, heating above 30 OC to dissolution, and then cooling back to 25 "C or by dissolving SDS in water and then mixing with brine as needed. Initially supersaturated solutions were transparent over path lengths of 10 cm or less, and then crystallites appeared. The f i t detectable crystallites were small, like specks of dust. Crystallites then grew fast to 100 pm or larger and settled. Nucleation occurred sooner and growth of crystals was faster when samples were stirred or shaken, whereas unstirred supercooled samples could remain clear for hours. In vigorously stirred samples crystallization was completed within 30 min. Opalescence, which would indicate a substantial number of particles of size -0.1 pm, was not detected in samples, over path lengths of 10 cm or less, at any time after supercooling. All observations suggested that crystals grew within minutes or seconds once they were nucleated and that nucleation was the slow step in the phase separation process. The apparent absorbance over 1-cm path length of supersaturated solutions was followed (Figure 2) until crystallites became visible. What was measured as ab-
sorbance was in fact due to scattering by the clear supersaturated solutions and not to any impurities, because the specific absorbance--0.005 cm-l- was substantially higher than the specific absorbance of a 10 wt 9O aqueous solution of NaCl with no SDS-hem., In press. L. M. Kushner and W. D. Hubbard, J. Col/o/d Sci., 10, 428 (1955). H. F. Huisman, Proc. K. Ned. Acad. Ser. Wet., Ser. B, 67, 367, 375, 388, 407 (1964). K. W. Wagner, Arch. flektrotech. (Berlln), 2, 371 (1914), clted by T. Hanai in “Emulslon Science”, P. Sherman, Ed., Academic Press, New York, 1968, p 379. R. T. Roberts and C. Chachaty, Chem. Phys. Lett., 22, 348 (1973). Y. Talmon, H. T. Davis, L. E. Scrlven, and E. L. Thomas, Rev. Sci. Instrum., 50, 698 (1979).
Interactions and Aggregation in Mlcroemulsions. A Small-Angle Neutron Scattering Study R. Ober and C. Taupin” Physique de la Mati&e CondensOe, E R A . 542 du Centre National de la Recherche Scientlfique, Coll6ge de France, 75231 Paris Cedex 05,France (Received: January 22, 1980)
We performed a study of various systems of well-defined water-in-oil microemulsions by means of small-angle neutron scattering. The equation of state, which is deduced from the scattered intensity at zero angle, shows that the interactions in these systems are essentially of the hard-sphere type. Some systems behave as pure hard-sphere liquids, but a small attractive term has to be added in most cases. The order of magnitude of this attraction is not compatible with van der Waals forces. The analysis of the scattered intensity as a function of the momentum transfer shows that the hard-sphere type systems behave as isolated spheres. On the contrary, the attractive systems present several characteristic features of doublets of spheres. The physical origin of the attraction is discussed.
Introduction In the past few years, much interest has been raised by microemulsions. These transparent fluid systems, which were identified as colloidal systems by Schulman in 1943,l have the capability of solubilizing water and oil in nearly all relative proportions. Various theoretical and experimental articles2 have been published in an attempt to understand their structure, stability, and occurrence as a function of chemical composition. In a previous paper,3 we determined the structure of typical microemulsion systems (water, cyclohexane, sodium dodecyl sulfate, and 1-pentanol) in the oil-rich region. The main conclusions were the following: The microemulsion is well described by spherical water droplets dispersed in an oily medium, the polydispersity in size being remarkably low. At constant ionic strength, the size of the water core of the droplets is determined by the area per polar head of the soap molecule which remains almost constant when the ratio of soap to water is varied. These conclusions are in good agreement with the experimental findings of other authors on similar ~ y s t e m s . ~ In our previous studies, we explored several systems the characteristics of which are reported in Table I. This table shows that besides decreasing the area per polar head of the soap molecule, which was expected, the increase of the ionic strength of the aqueous phase induces a decrease of the hydrodynamic volume. In fact, the hydrodynamic thickness rh - rw decreases from >23 to 16.5 A, revealing that the continuous phase has been expelled from the interfacial film. 0022-3654/80/208~-2418$01.OO/O
TABLE I: Structural Data for Four Types of Microemulsionsa A
B
C
1.25 0.18 34.5 52 46.5 12 18
2.5 0.087 56.5 68 65.5 9 23
3.75 0.045 83 66 92 9 27
NaCl mol/L
w/sb’
P/Ci rw, A A,g At rwe ‘4 r C - r W ,A rh - r w , f A
B1M 1 2.5 0.026 66.5 52 74.5 8
16.5
a Systems A, B, C correspond to increasing amounts of solubilized water.3 System B1M (1M NaCl) corresponds to unpublished results of J. P. L e Pesant. W/S is the weight ratio of water to soap. cP/C is the volume ratio of pentanol to cyclohexane in the continuous phase.j rw is the radius of the water core. e rc is the radius which is not penetrated by the continuous phase (see ref 3). f % is the hydrodynamical radius. g A is the area per polar head of the SDS molecules.
One of the puzzling questions which is raised by the wide domain of existence of the microemulsionsis the possibility of structural changes associated with an increase of the amount of the dispersed water phase. It is known that a good approximation to obtain systems of variable droplet concentration is to maintain the water-to-soap ratio, WIS, constant and to add to the water-soap-cyclohexane mixture the minimum quantity of alcohol which is necessary to obtain a clear phase. The “titration curve” (pentanol volume vs. cyclohexane volume) shows that the pentanol 0 1980 American Chemical Society
