J . Phys. Chem. 1988, 92, 3531-3537
the Mo 3d5,, level at 228.8 eV to the Mo(I1) state. The presence of the new line at 516.9 eV after relatively high Mo coverage of the a-Alz03surface is assigned to an interatomic Auger transition in terms of Mo-oxygen interaction similar (but at a different energy) to that observed in the case of Mo on silica. The presence of an oxide oxygen at 530.8 eV was not immediately apparent, and this may be due to the relatively lower 0 1s binding energy in the a-A1203at 531.8 eV. However, its presence can be deduced from the relative broadening of the 0 Is spectral line at higher Mo coverage. The stability of this interatomic Auger transition even after the oxidation of molybdenum may indicate that the electrostatic environment of the interacting oxygen did not change. This may indicate that in fact the interacting Mo atoms are already in an oxidized state. Further evidence for this stable Mo(I1) oxide formation on the a-A1203surface can be deduced from the fact that the kinetic energy of the interatomic Auger transition line observed upon the oxidation of Mo 011silica is the same as in the case of unoxidized and oxidized Mo on a-A1203and differs by 1.5 eV from that of the unoxidized Mo on silica. The reduction of a freshly vapor-deposited Mo for 26 min on a-AI2O3by argon ion bombardment from Mo 3d5,? at 228.8 to 227.8 eV demonstrates clearly that molybdenum exists in fact in an Mo(I1) oxidized state, where the argon ions are expected to break these Mo-0 bonds. Codc1usion From the Results and Discussion the following can be concluded: Vapor-deposited molybdenum on silica shows that molybdenum atoms in the first few layers are interacting with the support. A metallic Mo state is obtained after higher Mo coverage. Exposure of the surface to air for 30 min results in the oxidation of Mo to Mo(1V) and Mo(V1) states. A metal-support interaction is observed in the first few layers of vapor-deposited Mo on an a-A1203single crystal. Vapor-de-
3531
posited Mo for 26 min results in the presence of an Mo(I1) state, which is oxidized to Mo(1V) and Mo(V1) states upon exposure of the surface to air for 30 min. Higher coverage of the surface by Mo results in the presence of molybdenum in the metallic state. The presence of Mo on silica or a-A1203surfaces at different concentrations and oxidation states can be monitored by the relative changes in the intensity and energy of the first and the second (0 2p) transitions in the valence-band energy region. Depending on the Concentration of Mo on either silica or aAl,03, the interaction between the metal and the support may vary. At higher concentrations this results in an interatomic Auger transition involving M o - 0 atoms in the form of a chemical bond. Such an interatomic Auger transition results in the presence of a spectral line at a kinetic energy of 5 18.4 eV in Mo on silica and at 516.9 eV in Mo on a-AlzO3. The 0 KVV Auger transition energies and line shapes are also modified. The oxidation of Mo on silica by exposure to air results in the decrease of the interatomic Auger transition kinetic energy from 5 18.4 to 5 16.9 eV, while it remains unaffected in the case of oxidized Mo on a-AlZO3. This suggests that Mo at a certain concentration on a-A1203is already in an oxidized state and that further oxidation did not significantly alter the electronic environment of the active and interacting Mo-0 atoms responsible for this transition. On the contrary, molybdenum on silica is mainly in the metallic state, and its oxidation results in a considerable change in the electronic environment of the interacting Mo-0 atoms. The observation that the Si 2p and AI 2p line shapes remain symmetrical through all these experiments clearly indicates that interatomic Auger transitions occur mainly between Mo and oxygen atoms. One may expect that the differences in the interaction of molybdenum with AI2O3and SiO, will have implications for the catalytic behavior of such systems. This point will be dealt with in a separate article. Registry No. Mo, 7439-98-7; A1,0,, 1344-28-1; SiO,, 7631-86-9.
Double-Layer and Solvation Forces Measured in a Molten Salt and I t s Mixtures with Water R. G . Horn,*,+D. F. Evans,$and B. W. Ninhamt Department of Applied Mathematics, Research School of Physical Sciences, Australian National University, Canberra 2601, Australia and Department of Chemical Engineering and Materials Science, University of Minnesota, Minneapolis, Minnesota 55455 (Received: August 4, 1987; In Final Form: December 28, 1987)
Measurements have been made of the force between molecularly smooth mica surfaces immersed in ethylammonium nitrate, which is a molten salt at room temperature, and in mixtures of this salt with water across the concentration range from loe4 M to that of the pure salt, which is 11.2 M. At low concentrations the salt behaves as a typical 1:l electrolyte, and we measure an electrical double-layer force whose range decreases with increasing salt concentration. At high concentrations, above about 1 M, the double-layer force becomes so weak and short-ranged that it is completely dominated by a solvation force extending up to 5 nm. In the pure molten salt the solvation force is an oscillatory function of surface separation comparable to that measured in simple nonpolar liquids. No monotonic component of solvation force is found.
Introduction Ethylammonium ,-,itrate (EAN) is a salt which is molten at room temperature.’ It is miscible with water in all proportions and so provides an to study electrolytes at uncommonly high concentrations. addition, it is an extensively hydrogen-
bonded liquid that mimics certain of the unusual properties of water.2 We are interested in it for both reasons. One of the most vital and least understood properties of water is its interaction with different molecules, molecular groups, and surfaces in ways that enable them to be classified as hydrophilic
*To whom correspondence should be addressed at Materials Science Building 223, National Bureau of Standards, Gaithersburg, MD 20899. ‘Australian National University. University of Minnesota.
( 1 ) Evans, D. F.; Yamauchi, A,; Roman, R.; Casassa, E. 2. J . Colloid Interface Sci. 1982, 88, 89. (2) Evans, D. F.; Chen, S.-H.;Schriver, G W.; Arnett, E. M. J . Am. Chem. SOC.1981, 103, 481.
f
0022-3654/88/2092-3531$01.50/0
0 1988 American Chemical Society
Horn et al.
3532 The Journal of Physical Chemistry, Vol. 92, No. 12, 1988 or hydrophobic. Once so classified, we have a surprisingly accurate expectation of how these entities will interact with each other in water, yet the details remain dim. Presumably, molecular dipoles and the particular hydrogen-bonding characteristics of water are important ingredients. A characteristic of small hydrophobic molecules is their negative entropy of transfer from a nonpolar solvent into water, which is explained by saying that the water molecules are somehow more ordered around the molecule by forming a hydrogen-bonded “cage”. It has been found that certain nonpolar gases show the same behavior in ethylammonium nitrate, suggesting a “solvophobic effect” in that liquid too, again attributed to its ability to form extensive hydrogen-bonded networks.2 Some years ago Rand, Parsegian, and co-workers demonstrated clearly the existence of a hydration repulsion between hydrophilic, uncharged surfaces:phospholipid bilayer^.^" This repulsion accounts for the swelling of these and other surfactant lamellar phases in water, an effect that cannot be explained by any classical force such as the van der Waals or electrical double-layer interactions. There have been various attempts to account theoretically for the hydration repulsion, but so far there is no general agreement about the physical origins of the effect. Suggestions have included perturbation of the hydrogen-bonded structure of water,7 inhomogeneous polarization of ~ a t e r ,and ~ . image ~ charge interactions,’O although recent work” suggests that proper calculation of electrostatic effects cannot produce the requisite repulsion. A similar swelling of a lamellar phospholipid phase has been observed in ethylene glycol’2 and in ethylammonium nitrate,I3 with a repulsion between the lipid head group surfaces extending to 2.5 nm. Since the latter liquid is effectively an 11 M electrolyte, it would screen any electrostatic interaction very effectively. The observation therefore supports the idea that the repulsion stems from the hydrogen-bonding properties of the solvent. Further light has been shed on the problem by the series of force measurements between mica surfaces immersed in various liquids, obtained using the surface force apparatus developed by Israel a ~ h v i l i . ’In ~ aqueous electrolyte solutions at low concentrations the experimental results are consistent with the predictions of the Derjaguin-Landau-Verwey-Overbeek (DLVO) theory, namely, that the interaction is the sum of an electrical double-layer repulsion and a van der Waals a t t r a c t i ~ n . ’ ~ . ’The ~ former is calculated from the Poisson-Boltzmann equation for the distribution of ions in the vicinity of a charged surface, as described in any textbook on colloid science. This leads to a quasiexponential repulsion with a decay length (the Debye length) that decreases as the ionic strength increases. At moderate to high concentrations of simple electrolytes PashleyIs has shown that there are deviations from DLVO theory. There appears to be an additional short-range repulsion between the mica surfaces once the number of cations adsorbed from solution becomes significant, these cations apparently remaining
(3) LeNeveu, D. M.; Rand, R. P.; Parsegian, V . A. Nature (London) 1976, 259, 601.
(4) LeNeveu, D. M.; Rand, R. P.; Parsegian, V. A,: Gingell, D. Blophys. J . 1977, 18, 209. ( 5 ) Parsegian, V . A,; Fuller, N.; Rand, R. P. Proc. Nut/. Acad. Sci. ( U . S. A . ) 1979, 76, 2750. (6) Lis, L. J.; McAlister, M.; Fuller, N.; Rand, R. P.: Parsegian, V. A. Biophys. J . 1982, 37, 657. (7) Marcelja, S.; Radic, N. Chem. Phys. Lett. 1976, 42, 129. (8) Gruen, D. W. R.; Marcelja, S. J . Chem. SOC.,Faraday Trans. 2 1983, 79, 225. (9) Belaya, M. L.; Feigel’man, M. V.; Levadny, V . G. Chem. Phys. Lett. 1986, 126, 361. (10) Jonsson, B.; Wennerstrom, H. J . Chem. Soc., Faraday Trans. 2 1983, 79, 19. (11) Attard, P.; Mitchell, D. J. Chem. Phys. Lett. 1987, 133, 347. (12) Persson, P. K. T.; Bergenstahl, B. A. Biophys. J . 1985, 47, 743. (13) Evans, D. F.: Kaler, E. W.; Benton, W. J. J . Phys. G e m . 1983, 87, 533. (14) Israelachvili, J. N.; Adams, G . E. J . Chem. SOC.,Faraday Trans. I 1978, 74, 975. (15) Pashley, R. M. J . Colloid Interface Sci. 1981, 83. 5 3 1
TABLE I: Compositions of Solutions Studied in These Experiments, Giving the Concentrations of Ethylammonium Nitrate (EAN) Expressed in Various Units EAN molality molarity vol wt mol (mol/kg (mol/L fraction fraction fraction of water) of soh) m 1 .oo 1.oo 1.oo 11.2 0.90 0.92 0.65 100.9 10.1 0.83 0.86 0.50 54.7 9.4 0.74 0.70 7.9 0.32 26.1 0.50 0.55 0.17 11.2 5.7 0.10 1.13 0.12 0.022 1.25 8.9 x 10-3 1.1 x 10-2 1.8 x 10-3 0.10 0.10 8.9 x 10-4 1.1 x 10-3 1.8 x 10-4 1.0 x 10-2 1.0 x 10-2 8.9 X 10-5 1.1 x 10-4 1.8 x 10-5 1.0 x 10-3 1.0 x 10-3 8.9 X 1.1 x 10-5 1.8 x 10-6 1.0 x 10-4 1.0 x 10-4
hydrated and so in turn hydrating the surface. Once again, a detailed understanding of this effect, also dubbed a hydration repulsion, is lacking. Closer examination of the short-range force in this system has shown that in fact it is (at least under some circumstances) an oscillatory function of surface separation.I6 Similar oscillatory forces had been observed previously in simple nonpolar and polar and ascribed to the arrangement of molecules of the liquid in layers adjacent to the smooth solid surface. The spacing of the oscillations is commensurate with the diameter of the solvent molecule. The short-range (or solvation) force measured in the aqueous electrolytes does, however,’differ from that so far measured in other liquids. The oscillations appear to be superimposed on a monotonic background that becomes progressively more repulsive as the surface coverage of cations on mica increases.I6 It is this monotonically repulsive component that is thought to be related to the repulsion observed between lipid bilayers. In an endeavor to understand its origins, a series of liquids of varying properties has been investigated in the surface force apparatus. What property of water bestows this ability to support a repulsive force over a few nanometers? Can that property be found by systematically investigating other liquids with and without significant dipole moments and hydrogen bonding? In a series of simple nonpolar liquids the force is found to be oscillatory with no underlying repulsive comp~nent.’~*’~ In a polar liquid with no hydrogen bonding the same is true,” suggesting that dipolar interactions are not the answer. In liquids that have some hydrogen bonding it is still true, although the only cases studied so far (methanolz0 and ethylene glycol2’) have rather limited hydrogen bonding, which would not allow networks as extensive as those found in water to exist. Part of the motivation for the present work is to extend this series to a liquid that is known to have extensive hydrogen bonding. Ethylammonium nitrate is a suitable candidate that is already known to share some of water’s unusual properties. In addition to those mentioned above, it has been found that surfactants aggregate into in this liquid, again suggesting a solvophilic/solvophobic dichotomy comparable to that found in water. Another reason for interest in this liquid is that it provides an opportunity, in its mixtures with water, to study an electrolyte at extremely high concentrations, so high that the calculated Debye length would be less than the molecular size. In this paper we present measurements of the force between mica surfaces immersed in water-ethylammonium nitrate mixtures across the entire concentration range from very dilute electrolyte (16) Pashley, R. M.; Israelachvili, J. N. J . Colloid Interface Sci. 1984, 101, 511. (17) Horn, R. G.; Israelachvili, J. N. J . Chem. Phys. 1981, 75, 1400. (18) Christenson, H. K . J . Chem. Phys. 1983, 78, 6906. (19) Christenson, H. K.; Horn, R. G. Chem. Phys. Lett. 1983, 98, 4 5 . (20) Christenson, H. K. J . Chem. Soc., Faraday Trans. 1 1984,80, 1933. (21) Christenson, H. K.; Horn, R. G. J . Co/loid Interface Sci. 1985, 103, 50. (22) Evans, D. F.; Yamauchi, A.; Wei, G. J.; Bloomfield, V . A. J . Phys. G e m . 1983, 87, 3 5 3 7 .
The Journal of Physical Chemistry, Vol. 92, No. 12, 1988 3533
Double-Layer and Solvation Forces
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crossed cylindrical mica surfaces in dilute aqueous solutions of ethylammonium nitrate, plotted as a function of the surface separation D. Filled circles show experimental points, and solid lines show numerical solutions of the nonlinear Poisson-Boltzmann equation with a charge regulation model described in the Discussion. The concentrations are (a) 0.80 X lo4 M and (b) 0.93 X lo-' M, which are within experimental error of the nominal values of lo4 and lo-' M respectively.
to the pure molten salt. The results will be presented in that order and then discussed in terms of the questions raised above.
Experimental Section The results presented were obtained by using the surface force apparatus developed by I~raelachvili.'~ Briefly, the force between crossed cylinders of molecularly smooth mica is determined with a resolution of lo-' N as a function of their separation, which is measured by optical interference to an accuracy of 0.1-0.2 nm. According to the Derjaguin a p p r ~ x i m a t i o n the , ~ ~force divided by 2?r times the radius of the cylinders, about 1 cm, is equal to the interaction energy between two flat plates at the same separation. Water was purified by following the procedure described by P a ~ h l e y .Ethylammonium ~~ nitrate was prepared by interacting equimolar amounts of ethylamine and nitric acid according to the procedure described by Evans et al.' If not used within a few days of preparation, the molten salt turns slightly yellow due to the accumulation of nitrous oxide impurities, but the discoloration can be removed by adding acetonitrile, allowing the mixture to stand with activated charcoal, filtering, and then pumping off the acetonitrile under vacuum.22 In Table I there is a list of the compositions used in these experiments, expressed in various units. The conversions were effected by using density and partial molar volume dataZ5interpolated to 22 OC, the temperature at which all the present measurements were made. Results Forces measured between mica surfaces in low concentrations of ethylammonium nitrate (approximate concentrations of 104, and lo-' M) are shown in Figures 1 and 2. A t long range there is a double-layer repulsion whose extent decreases as the electrolyte concentration increases, as predicted by DLVO theory, based on the Poisson-Boltzmann equation. The solid lines in these figures show numerical solutions of the nonlinear Pois(23) Derjaguin, B. V. Kolloid Zh. 1934, 69, 155. (24) Pashley, R. M. J. Colloid Interface Sei. 1981, 80,153. (25) Allen, M.; Evans, D. F.; Lumry,R.J. Solution Chem. 1985, 14,549.
-
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0 (nm) Figure 2. Forces measured in ethylammonium nitrate solutions of nominal concentrations and lo-' M. The first two are the same results as shown in Figure 1, plotted on an expanded scale. In c and d the crosses show repeat measurements; the lines are theoretical curves for 0.91 X and 0.93 X lo-' M, respectively. (a) At M, the surfaces jump in from D = 2 nm (arrowhead) to adhesive contact at D = 0. (b) At M there is a very shallow minimum at D = 0.8 nm, manifested by a small outward jumpI4 indicated by the arrowhead, but primary contact is prevented by a repulsive barrier at around D = 0.5 nm. (c) As the concentration is increased, the barrier is found slightly further out, at D = 0.8-0.9 nm, and no minimum is observed at M. (d) At lo-' M there is no minimum near contact, but there is a secondary minimum at D = 6-7 nm.
son-Boltzmann equation from the algorithm of Chan et a1.26and a surface charge regulation model discussed in the next section. At separations beyond 5 nm there is excellent agreement with experiment, obtained using concentrations that are within experimental error of the nominal values. The forces at short range are a little more complicated. In pure water and in lo4 M ethylammonium nitrate there is a van der (26) Chan, D. Y .C.; Pashley, R. M.; White, L. R. J. Colloid Interface Sei. 1980, 77, 283.
3534 The Journal of Physical Chemistry, Vol. 92, No. 12, 1988 Waals attraction that dominates at short range and pulls the surfaces into adhesive contact at D = 0, as predicted by DLVO theory. However, at higher concentrations this so-called primary minimum is not observed. The van der Waals force is masked M and by a steep short-range repulsion that is present at above, as shown in Figure 2. At M there is a very shallow minimum in the force at D = 0.8 nm, though the minimum is still in the repulsive region ( F > 0). As the surfaces approach, the force reaches a maximum at D slightly greater than 0.8 nm, then begins to decrease as the van der Waals attraction becomes significant, but then almost immediately increases sharply as a steep repulsion is encountered at around 0.5-0.7 nm. The resultant shallow minimum is detected as a very weak “adhesion”, giving a small jump outwards from D = 0.8 nm when the surfaces are separated, indicated by the arrowhead in Figure 2 . At IO-*M ethylammonium nitrate no minimum is found, and there is a steep repulsion at D = 0.8-0.9 nm. At lo-’ M the steep repulsion is still there, and the only minimum is a secondary minimum at around D = 6-7 nm. At very high concentrations, as we approach the pure ethylammonium nitrate end of the spectrum, we find a force that is an oscillatory function of surface separation. This is shown in Figure 3. The spacing of the oscillations is 0.5-0.6 nm, and this does not vary measurably as the concentration changes. As the fraction of water in the mixture decreases, more and more oscillations are found, and by the time we reach the pure molten salt they extend to almost 5 nm. Here there are probably eight or nine oscillations, of which only the weakest four or five are measurable. At separations less than about 2 nm the repulsive peaks become so strong that they prevent closer approach of the surfaces, a situation that is exacerbated by progressive flattening of the surfaces.17 Even at 10% ethylammonium nitrate v/v (approximately 1 M), where the oscillations have the shortest range, there is no measurable double-layer force. The model discussed below would predict a surface potential of less than 10 mV at this concentration, with a Debye length of 0.3 nm, giving a force so weak that it is exceeded by the van der Waals attraction at all separations. At higher ethylammonium nitrate concentrations it would be weaker still. Beyond the range of the oscillations an attraction is observed. The whole attractive region of the force cannot be measured with this apparatus, but it is possible to determine the gradient at one point.14 The result shows that the attraction in 10% ethylammonium nitrate is significantly weaker than that in water. We will elaborate on this in the Discussion. There is no clear indication in Figure 3 that the oscillatory force is superimposed on any monotonic background force, in contrast to that observed in concentrated aqueous solutions of potassium ch1oride.l6
Discussion At low concentrations in water, ethylammonium nitrate behaves as a simple electrolyte. There is an electrical double-layer force between mica surfaces, with a Debye length that decreases with concentration as expected. By comparison of the measured force with that predicted from the Poisson-Boltzmann equation, it is possible to determine the surface potential of isolated surfaces, as a function of concentration. When this is done, it is found that the potential first decreases, then increases, and then decreases again as the electrolyte concentration is increased from IO4 to lo-’ M. To explain similar behavior in solutions of the alkali-metal chlorides, Pashley developed a model of competitive cation adsorption to micaI5 based on the earlier ion-exchange model of Ninham and Parsegian.*’ The adsorption of hydrogen and other cationic species to the negative mica surface is described by mass-action relations with suitable dissociation constants. If a cation occupies a surface area greater than the lattice area, then (27) Ninham, B. W.; Parsegian, V . A. J . Theor. Biol. 1971, 31, 405
Horn et al.
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Figure 4. Top: variation of surface potential with concentration for isolated mica surfaces in aqueous solutions of ethylammonium nitrate, as predicted by the ion-exchange model of Pashley,Is and Miklavic and Ninham2*described in the Discussion (the horizontal scale is pConc = -log concentration). Three cases are shown, corresponding to a surface area per adsorbed ethylammonium ion of 0.54 nmz (dotted line), 0.55 nm2 (solid line), and 0.56 nm2 (dashed line). The theoretical curves shown in Figures 1 and 2 were generated by using 0.55 nm2 and the concentrations indicated by the plus signs. Bottom: fractional surface coverage of ethylammonium ions (solid line) and hydrogen ions (dashed line) from the same model with 0.55 nm*/ion. Because this is greater than the lattice area of 0.48 nm2, the ethylammonium ion coverage cannot exceed 87%.
Pashley’s model and found an inconsistency in using the simple mass-action equations when the area occupied by adsorbed cations can restrict further adsorption. To correct this, they derived slightly more complex mass-action equations; otherwise, the essence of the model remains unchanged. With the more recent model, the variation of surface potential with concentration of ethylammonium nitrate can be explained satisfactorily if the ethylammonium cation is ascribed a surface dissociation constant of 10-3,s5M and a surface area of 0.55 nm2, combined with a hydrogen ion dissociation constant of 10-6.0M,Is an area per adsorbed hydrogen equal to the mica lattice area of 0.48 nm2, and the experimental pH of 5.7. Figure 4 shows how the surface potential and surface coverage of the two cationic species vary with concentration under these conditions. Potential versus concentration curves are also shown for two slightly different cation areas to illustrate the sensitivity to this parameter. The theoretical double-layer repulsion is computed by using the algorithm of Chan et aLZ6to generate a numerical solution of the nonlinear Poisson-Boltzmann equation. To this is added a van der Waals attraction with a Hamaker constant of 2.2 X 1OT20 J, appropriate for water,I4 to generate the curves shown in Figures 1 and 2 . Once the above parameters are established, the infinite-separation surface potential is determined by the electrolyte concentration and can no longer be varied independently in fitting the forces. Plus signs in Figure 4 mark the concentrations used to fit the curves; they are within error of the nominal experimental values. Furthermore, the same cation adsorption model predicts how the surface charge varies as a function of surface separation, regulating between the limits of constant charge and constant potential interactions. This regulation is taken into account in obtaining the theoretical force curves. The agreement with experiment over a range of concentrations is very satisfactory. At concentrations of M and above, the force at short range deviates from the predictions of DLVO theory. There is a repulsion that dominates any van der Waals attraction. This be-, havior is reminiscent of that found in the alkali-metal chlorides
and other electrolytes,1s~z9-3’in which a short-range repulsion appears above a certain concentration, dependent on the particular cation but generally in the range 10-4-10-2 M. For the alkali metals PashleyI5 has shown that the “hydration repulsion” occurs when the cation, still hydrated, is the predominant species adsorbed to the surface, and this gives rise to the extra force. In the present case the appearance of the repulsive force can also be correlated with the surface coverage of the ethylammonium cation, In Figure 4 (bottom) we see that the ion-exchange model discussed above predicts that as concentration increases, the ethylammonium displaces hydrogen ions from the surface, the M. exchange taking place substantially between loT4and Because it occupies an area greater than the site area, the ethylammonium coverage can never exceed 87%, and it is already M. At this concentration and above, there approaching this at appears to be a “hard wall” in the force at D = 0.9 nm (Figure 2 ) , suggesting that a fully packed layer of ethylammonium cations M the model is about 0.45 nm thick on each surface. At predicts a coverage of 51%, and the hard wall appears at D = 0.5 nm, consistent with slightly more than half a monolayer per surface. In the present case the extra repulsive force appears to be very and lo-’ M, although at loT2M there does steep, at !east at seem to be a “softer” repulsion extending to 3 nm. A steep repulsion would suggest a steric barrier at a fixed position, consistent with the idea of a Stern layer of ethylammonium ions, rather than the softer repulsion found with metal cations15~29J0 and ascribed to the hydration of the adsorbed ions. It is reasonable to expect ethylammonium to be less hydrated than the metal ions. In previous measurements on a series of tetraalkylammonium salts3’ there was also evidence for Stern layers. In that case the extra repulsion was softer than a “hard wall” but could be accounted for surprisingly well by assuming that the plane of charge was shifted out from the surface to the edge of the Stern layer (while the van der Waals force still acted from the mica surface). There was no need to invoke an additional “hydration” force. It has been observed that both pure ethylammonium nitrate and its aqueous solutions are autophobic on mica, Le., a droplet first spreads on the surface and then retracts until the contract angle is about 40’. This suggests that the cation adsorbs with its charged ammonium end next to the mica surface, presenting ethyl groups to the liquid phase, reminiscent of the way that cationic surfactants adsorb to mica and render it hydrophobic. Despite that, no hydrophobic attraction or strong adhesion of the kind observed between adsorbed surfactant monolayers32was seen at any concentration in the present experiments. An estimate given below shows that some water is still present in the Stern layer, which may account for this difference. Nevertheless, the ethylammonium ion is not sufficiently hydrated to give rise to a secondary hydration repulsion of the kind found with metal cations. In a discussion of the results at concentrations of 1 M and above, it is perhaps easiest to start at the pure ethylammonium nitrate end of the spectrum (Figure 3f). Here the force is an oscillatory function of separation, with a spacing between force minima of 0.5-0.6 nm. Not all of the oscillations are measurable, the repulsive barriers at D < 2 nm being so strong that they prevent closer approach of the surfaces. The weakest minimum is probably about the ninth from mica-mica contact. This force is very much like those found in a variety of nonpolar and polar where typically eight to ten oscillations are measured, with minima below the F = 0 axis and maxima above it and a spacing commensurate with the molecular diameter. In the present case the liquid consists of ions rather than molecules, and choosing the appropriate “diameter” is not so straightforward. As a crude estimate, we can use a method that Hugo Christenson ~~
(29) Pashley, R. M.; Israelachvili,J. N. J . Colloid Interface Sci. 1984, 97,
446. (30) Pashley, R. M. J . Colloid Interface Sci. 1984, 102,23. (31) Claesson, P. M.; Horn, R. G.; Pashley, R. M. J . Colloid Interface Sci. 1984, 100,250. (32) Pashley, R. M.; McGuiggan, P. M.; Ninham, B. W.; Evans, D. F. Science (Wushingfon,D.C., 1979-) 1985, 229, 1088.
3536 The Journal of Physical Chemistry, Voi. 92, No. 12, 1988 (private communication) has found effective for other liquids. From the known density we calculate the volume per ethylammonium nitrate ion pair and take the cube root of that to get 0.53 nm, a value that agrees with the measured spacing. It is reasonable to suggest, therefore, that the fused salt forms layers in the confined space between mica sheets in close proximity in the same way that molecular liquids do and that both the cation and the anion are present in the layers in roughly equal numbers. The force measurements cannot tell us any more about the arrangement of ions within the layers. For example, we cannot rule out the possibility that there are alternating "sublayers" of cations and anions that together form a single layer as we detect it, but such an arrangement seems unlikely on the grounds that it would involve a large electrostatic energy. As water is added progressively to ethylammonium nitrate, the range of the oscillatory force decreases. Clearly the number of oscillations in the force decreases markedly, from perhaps eight in the pure fused salt to two at 50% water. Very similar behavior and explained has been found in mixtures of nonpolar in the following way. The extensive oscillatory force in a pure liquid results from packing restrictions on molecules of a uniform size in a confined space: at certain separations of the solid surfaces it will not be possible to fill the space efficiently, and a high free energy result^.'^ If two molecular species of different size are present, there will be more possible ways to fill a given space, and therefore fewer unfavorable separations will be encountered. In the present case the ability of both ethylammonium nitrate and water to hydrogen bond will complicate the issue of how they pack between mica surfaces, but from the similarity of these results to those in nonpolar liquid mixtures it would seem that the simple size effects are the most important consideration. We have argued above that there should be roughly equal numbers of cations and anions in each layer of the fused salt, but the very first "layer" adjacent to each mica sheet is probably different. Since the mica is negatively charged we would expect a layer of cations to form next to it. This is just the Stern layer discussed earlier, and it would approximately neutralize the charge, so that there would not be a strong tendency for the next layer to consist solely of anions. If there are only cations in the first layer, it will be thinner than subsequent layers. The measurements in dilute solutions give a value of 0.45 nm. The surface area occupied by each ion in the Stern layer was deduced to be 0.55 nm2; multiplying this by the layer thickness gives a volume of 0.25 nm3, which is more than twice the volume one would estimate for the ethylammonium cation. Since the presence of a sizeable fraction of anions in this layer would not be consistent with the low surface charge measured at lo-* and lo-' M, this suggests that the cations in the Stern layer are significantly hydrated. The gradient of force measured at a single point in the attractive region can be used to calculate the Hamaker constant ( A ) if the van der Waals force is assumed to have the nonretarded form FIR = A / 6 D Z . This gives a value in 10% ethylammonium nitrate of A = 4 X lo-*' J, compared with 2.2 X J for water. It was not measured accurately in more concentrated mixtures, but the attraction certainly remained much smaller than in water. There is not sufficient optical data on ethylammonium nitrate available to allow a theoretical calculation of the Hamaker constant for comparison. The experimental value is very low; possible explanations include the absence of an infrared absorption, screening of the zero-frequency term, and retardation effects, which would also alter the functional form of the force and make the above calculation inappropriate. There is a qualitative difference between the forces measured between mica surfaces in pure water and pure ethylammonium nitrate. In the former there is an electrical double-layer force dominant at long range and an apparently monotonic van der Waals attraction strong enough to dominate at short range. Oscillations are not seen, although that does not prove that they are absent. Because the water molecule has a small diameter and the Hamaker constant in this system is large, any oscillations would (33) Christenson, H. K. Chem. Phys. Lett. 1985, 118, 455.
Horn et al.
occur only at very short range and be entirely within the attractive regime,I6 in which case they would not be detected by the present technique. In ethylammonium nitrate the double-layer force is completely screened, the van der Waals force is much weaker, and the "molecules" are larger, so that the oscillatory force extends further, dominating the interaction at short range. Thus we can rationalize the differences between the two liquids without resorting to a discussion of any differences in their hydrogen bonding. The oscillatory force measured in ethylammonium nitrate does not appear to be superimposed on any monotonic repulsive component. This is in contrast to the force measured in concentrated aqueous solutions of KCl, where a monotonic background to the hydration force has been compared to the hydration force measured between lecithin bilayers in water.I6 However, a clear distinction between the two systems must be made. The polar headgroup of lecithin is strongly hydrated, so the bilayer surface is intrinsically hydrated, and a force arises from having to dehydrate it when two such surfaces are pushed together. The mica surface itself is not hydrated in pure water. Only when a significant number of hydrated metal cations are adsorbed to it does a hydration force appear; there is a "secondary" hydration of the mica via the cation^.'^ The situation in ethylammonium nitrate is similar. There is a strong repulsive force between lecithin bilayers, indicating that lecithin is solvated by this liquid, but mica is not. In this sense our experiments do not distinguish between the properties of ethylammonium nitrate and water but serve to remind us that the mica surface is different from that of lecithin. We have not explored a situation in which there could be secondary solvation of mica in ethylammonium nitrate via some other species being solvated and adsorbing to it. Cations would not adsorb as readily as they do from aqueous solution because the liquid is itself a salt and would screen any electrostatic attraction to the surface.
Conclusions At low concentrations in water, ethylammonium nitrate behaves as a simple electrolyte and the force between mica surfaces follows the predictions of DLVO theory. As the concentration increases, the cations adsorb to mica in a way that is described satisfactorily by the ion-exchange model proposed by P a ~ h l e y and ' ~ revised by Miklavic and Ninham.** When cations are adsorbed to 50% or more of the mica lattice sites, they form a Stern layer that causes a steric repulsive barrier preventing the surfaces from coming into primary adhesive contact. The adsorption also reduces the effective surface charge, so that at high concentrations (above 10% v/v, or about 1 M) the double-layer repulsion is so weak and so short ranged that it is dominated by the van der Waals attraction. However, the finite sizes of the ethylammonium and nitrate ions, the former already evident from the measurable thickness of the Stern layer, now play a role. At 50% v/v another layer, containing both cations and anions, appears between the surfaces. This is manifested in a second, weaker repulsive barrier at a larger separation. At still higher concentrations more and more layers are detected, until in the pure fused salt there are eight or nine, giving rise to a solvation force that is an oscillatory function of surface separation. The oscillatory force is very similar to that found in nonpolar l i q ~ i d s . ' ~ ,Neither '~ the ionic nature of etylammonium nitrate nor its extensive hydrogen bonding have had much effect: the packing of the finite-sized ions is the predominant feature determining the force. No monotonic solvation force comparable to the hydration force observed between mica surfaces in other aqueous s ~ l u t i o n s ' ~ ~ ~ ~ ~ ~ ~ is found here. That can be considered as a secondary hydration force occurring only when a hydrated solute ion is adsorbed to mica, but it does not happen with ethylammonium nitrate considered as the solute. Taking it as the solvent, we have not explored the situation of other ions dissolved in it but have argued that there would be little driving force to adsorb cations, whether solvated or not, from a medium of such high ionic strength. Solvation of the head groups coupled with extensive hydrogen bonding remains the most likely explanation for the repulsion
3537
J . Phys. Chem. 1988, 92, 3537-3542
between lipid bilayers observed in both water and ethylammonium nitrate, but it is clear from the present measurements that neither of the pure liquids solvates the mica surface in the same way.
Acknowledgment. These experiments commenced while R.G.H. was on leave at the University of Minnesota, and he acknowledges
the hospitality there of Matt Tirrell. We thank Hal Beesley and Patty McGuiggan for preparing the ethylammonium nitrate samples and Stan Miklavic and Hugo Christenson for helpful discussions. Registry No. EAN, 22113-86-6.
Complexation of Redox-Active Surfactants by Cyclodextrins Abigail Diaz, Pablo A. Quintela, Jodi M. Schuette, and Angel E. Kaifer* Department of Chemistry, University of Miami, Coral Gables, Florida 331 24 (Received: August 7 , 1987; In Final Form: October 27, 1987)
The effects of cyclodextrins on the aggregation behavior and electrochemicalproperties of N-ethyl-N’-hexadecyl-4,4’-bipyridinium bromide (CI6VBr2)and N-ethyl-N’-octadecyl-4,4’-bipyridinium bromide (ClsVBr2)were assessed by using electrochemical, optical, and surface tension techniques. The association constants of these two amphiphilic viologens with a- and @-cyclodextrin (ACD and BCD) were determined with a conductance method. Values in the range 103-104 M-I were obtained for all four complexes. In the presence of ACD, the first reduction couple of both viologens exhibits reversible, diffusion-controlled voltammetric behavior. ACD also inhibits completely the formation of dimers from the corresponding cation radicals. In striking contrast, the presence of BCD does not eliminate the precipitation on the electrode surface of the hydrophobic, reduced viologen species. Furthermore, BCD is unable to prevent the extensive dimerization of the viologen cation radicals.
Introduction Cy~lodextrinsl-~ are cyclic glucopyranose oligomers having a characteristic toroidal shape. These compounds are soluble in aqueous media because of the hydrophilic nature of the outer surface of the torus. The inner surface is more hydrophobic, and thus hydrated cyclodextrins represent a high-energy state that can readily accept guest molecules in place of the inner water molecules.’ Indeed, the resulting host-guest complexes are more stable as the hydrophobicity of the guest molecule increases. As expected for inclusion complexes, the better the fit of the guest molecule in the inner cavity of the cyclodextrin, the more stable the host-guest complex will @-Cyclodextrin (BCD, seven glucopyranose units) forms complexes with many organic molecules because its cavity is well suited to bind phenyl moieties and other functional groups. a-Cyclodextrin (ACD, six glucopyranose units) does not form so many complexes because of its smaller inner cavity. However, simple size considerations indicate that ACD is well suited to bind compounds having long alkyl chains such as surfactants. There have been some scattered reports on the interactions of cyclodextrins and surfactants. For instance, Ise and co-workers have reported on the association of cyclodextrins and common surfactants like sodium dodecyl sulfate and cetyltrimethylammonium bromide.4a They found that the addition of cyclodextrins increases the apparent critical micelle concentrations (cmc) of both surfactants. Using conductometric techniques, they determined equilibrium constants between ACD, BCD, and the two surfactants. Similar results have been reported for ACD by Satake et al.4b Thomas and co-workers have studied the formation of CD-surfactant-aromatic fluorophore ternary c o m p l e x e ~ . ~ Quite recently, Kusumoto et al. have also reported on the interaction of pyrene with BCD in aqueous surfactant solutions.6 Yasuda et al. have studied the interactions of several rather hy(1) Szejtli, J. Cyclodextrins and their Inclusion Complexes; Akademiai Kiado: Budapest, 1982. (2) Bender, M. L.; Komiyama, M. Cyclodextrin Chemistry; SpringerVerlag: Berlin, 1978. (3) Saenger, W. Angew. Chem., In?. Ed. Engl. 1980, 19, 344. (4) (a) Okubo, T.; Kitano, H.; Ise, N. J . Phys. Chem. 1976,80, 2661. (b) Satake, I.; Ikenoue, T.; Takeshita, T.; Hayakawa, K.; Maeda, T. Bull. Chem. SOC.Jpn. 1985, 58, 2746. ( 5 ) Hashimoto, S.; Thomas, J. K. J. Am. Chem. SOC.1985, 107, 4655. (6) Kusumoto, Y.; Shizuka, M.; Satake, I. Chem. Phys. Lett. 1986, 125, 64.
0022-3654/88/2092-3537$01 S O / O
drophobic viologens’ with cyclodextrins, aiming at the development of a practical electrochromic display system. Our group is interested in the redox properties of viologens solubilized in hydrophobic, membrane mimetic environments.8-’0 Therefore, we decided to investigate the effects of cyclodextrins on the electrochemical behavior and aggregation properties of the two viologen compounds CI6VBr2and C,,VBr2. These two
‘1fjVBr2
6-CH2-t
“3-c \
\
/N+-CH2-CH3
2 Br-
C18VBr2
compounds (CI6VBr2and C18VBr2)were selected because of their amphiphilic nature that provides two possible binding sites for the cyclodextrin hosts: the aromatic viologen group or the alkyl chain. We report here the results of this study.
Experimental Section Materials. The surfactant viologen bromides, CI6VBr2and C,,VBr2, were prepared according to published procedures for the asymmetric quaternization of 4,4’-bipyridine.” 1-Ethyl4-(4’-pyridyl)pyridinium bromide was first synthetized by mixing 4.0 g of 4,4’-bipyridine (Aldrich) with 20 mL of ethyl bromide. The reaction mixture was kept at room temperature for 4 days and then at 40-50 OC for 3 more days. The resulting mixture was stirred in. 100 mL of toluene to remove the unreacted bipyridine and filtered to separate the solid product. Recrystallization of this solid from hot acetonitrile yielded a material with an N M R spectrum (DMsO-d,) consistent with the product structure. The second quaternization was performed in acetonitrile by reacting the ethylated intermediate with either hexadecyl or (7) Yasuda, A,; Kondo, H.; Itabashi, M.; Seto, J. J. Electroanal. Chem. 1986, 210, 265.
(8) Kaifer, A. E.; Bard, A. J. J. Phys. Chem. 1985, 89, 4876 (9) Kaifer, A. E. J. A m . Chem. SOC.1986. 108. 6837. (10) Quintela, P. A,; Kaifer, A. E. Langmuir 1987, 3, 769. (1 1) Pileni, M. P.; Braun, A. M.; Gratzel, M. Photochem. Photobiol. 1980, 31, 423.
0 1988 American Chemical Society
