Langmuir 1997, 13, 5433-5439
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Role of Tail-Tail Interactions versus Head-Group/Subphase Interactions in the Pressure-Area Isotherms of Fatty Amines at the Air-Water Interface. 1. Influence of Subphase Acid Counterions P. Ganguly* and D. V. Paranjape Physical and Materials Chemistry Division, National Chemical Laboratory, Pashan Road, Pune 411008, India
F. Rondelez Institut de Pierre et Marie Curie, 11 rue Pierre et Marie Curie, 75231 Paris Cedex 05, Paris, France Received October 8, 1996. In Final Form: July 11, 1997X The pressure-area isotherms obtained in the initial stages after the spreading of fatty primary amines on the aqueous subphase has been investigated as a function of pH, length of hydrocarbon chain, basicity of the amine, nature of the acid counterion, etc. The main emphasis is on the high-area liquidlike features observed prominently at low pH. The dependence on the specific nature of the acid counteranion or the nature of the amine (aliphatic or aromatic) highlights the importance of the interactions at the interface in deciding the nature of the pressure-area isotherm in addition to the well-known tail-tail interactions between the hydrocarbon chains. The marked influence of small amounts (∼2 × 10-4 M) of alcohol such as n-butanol in influencing the stability as well as the nature of the pressure-area isotherms of fatty acids and amines is demonstrated. The relation between the lift-off areas and length scales, such as the size of the solvated subphase ionic species, repulsive limits between charged -NH3+ head groups, and ion-pair interaction distances, as well as the valence of the subphase acid species is examined. The role of the area-expanding head-group/subphase interactions vis-a-vis the area-condensing tail-tail interactions is discussed in terms of recent two-state models for liquid monolayers.
I. Introduction There has been surprisingly little attention paid to the systematic study of the pressure-area isotherms of fatty amines compared to that of the fatty acids,1-3 although the current developments in several areas of materials science and biology makes such studies on fatty amines necessary. Studies4 on biologically important molecules include linear alkyl chain phosphatidylcholines or ethanolamines. Condensed Langmuir monolayers of primary amines have been used for depositing large complex ammonium salts of organic or inorganic anions at low bulk pH from an aqueous subphase,5 which include the titanyl oxalate anion,6 [TiO(C2O4)2]2-, or the chloroplatinate anion,7 [PtCl6]2-. The films of salts of fatty amines are useful, after suitable treatment, for forming ultrathin films of oxides8 such as TiO2 or for forming nanoclusters of Pt metal or other sulfides, for example. Most recent studies on the π-A isotherms on the existence of a universal or generic π-T phase diagram9,10 X Abstract published in Advance ACS Abstracts, September 15, 1997.
(1) Adam, N. K. The Physics and Chemistry Surfaces; Oxford Press: London, 1941. Harkins, W. D. The Physical Chemistry Of Surfac Films; Reinhold: New York, 1952. (2) Bell, G. M.; Combs, L. L.; Dunne, L. J. Chem. Rev. 1981, 81, 15. (3) Goddard, E. D.; Ackilli, J. A. J. Colloid Sci. 1963, 18, 585. Goddard, E. D.; Smith, S. R.; Kao, O. J. Colloid Sci. 1966, 21, 320. (4) See, for example: Baltes, H.; Schwendler, M.; Helm, C. A.; Mohwald, H. J. Colloid Interface Sci. 1996, 178, 135 and references therein. (5) Gains, G. L., Jr. Nature 1992, 298, 544. (6) Ganguly, P.; Paranjape, D. V.; Sastry, M. Langmuir 1993, 9, 571. (7) Ganguly, P.; Paranjape, D. V.; Sastry, M. J. Am. Chem. Soc. 1993, 117, 793. (8) Paranjape, D. V.; Sastry, M.; Ganguly, P. Appl. Phys. Lett. 1993, 63, 18. Sastry, M.; Paranjpe, D. V.; Ganguly, P. J. Electron Spectrosc. 1993, 67, 163.
S0743-7463(96)00981-X CCC: $14.00
at areas close to the cross-sectional area Aperp, perpendicular to the axis of the all-trans hydrocarbon chains, ignore the role of head-group/subphase interactions, except in describing11 the fractionation of subphase metal ions and protons between the bulk and the surface of the liquid for a floating Langmuir monolayer of fatty acids. The head-group/subphase interaction is expected to be important in the liquid-expanded (L1; see ref 10)/liquidcondensed (LE/LC) phase transitions at higher areas.12 It is here that one expects the relative roles of the tail-tail and the head-group-subphase interactions to be important. The head-group-subphase electrostatic interaction at large areas (relative to Aperp) become important when there is ion-pairing interaction, e.g., between the charged ammonium, -NH3+, head-group and a large acid counterion in the subphase.13 The organization of fatty amphiphiles at the air-water interface induces an organization of the subphase counteranions also at the interface14 (or vice versa) with its attendant constraints, including simple stereochemical constraints. Such a mutual organization of the charged species at the two sides of the air-water interface must influence the (9) See: Bibo, A. M.; Knobler, C. M.; Peterson, I. R. J. Phys. Chem. 1991, 95, 5591. (10) Bibo, A. M.; Peterson, I. R. Adv. Mater. 1990, 2, 309. (11) See: Bloch, J. M.; Yun, W.; Phys. Rev. A. 1990, 41, 844 and references therein. (12) See: Ruckenstein, E.; Li, B. J. Phys. Chem. 1996, 100, 3108 and references therein. (13) The problem is reminiscent of the changes in the pressure-area isotherms of fatty acids in the presence of large cationic complexes in the aqueous airface. (a) Murakata, T.; Miyashita, T.; Matsuda, M. J. Phys. Chem. 1988, 92, 6040. (b) Palacin, S.; Barraud, R. A. J. Chem. Soc., Chem. Commun. 1989, 45. (c) DeArmond, M. K.; Samha, H.; Dvorak, O. Langmuir 1994, 10, 343. (d) Samha, H.; DeArmond, M. K. Langmuir 1994, 10, 4157. (14) Cai, Z-H.; Rice, S. A. J. Chem. Phys. 1989, 90, 6716.
© 1997 American Chemical Society
5434 Langmuir, Vol. 13, No. 20, 1997
pressure-area isotherms. For convenience we introduce the nomenclature of airface and aquaface to distinguish, respectively, between the amphiphile layer on the air side of the interface and the counterion layer on the aqueous side of the interface.15 The lack of systematic investigations on the pressurearea isotherms of fatty amines16,17 must be partly because of the “bizarre” nature of the pressure-area isotherms of amines encountered in the early studies by Adams and co-workers.17 Among more important of these are the following: (i) There is sometimes a loss in pressure with time for the low bulk pH of the aqueous subphase. This effect is most prominent with HCl and is surprisingly absent or much less prominent with acids such as HClO4 or H2SO4. (ii) For stable fatty amine monolayers at the air-water interface, there is a marked increase in the area per amine molecule with time at a given surface pressure, π, even under highly acidic conditions or large areas. (iii) The area per fatty amine molecule shows a marked dependence on the nature of the anions in the subphase16 at low bulk pH. This is similar to the behavior of fatty acids spread on an aqueous subphase of high bulk pH (g10) of alkali metal or tetraalkylammonium hydroxides.18 The systematics of the nature and strength19 of the acid in the subphase in determining the characteristics of the π-A isotherms at high areas (>25-30 Å2; characteristic of the LC/LE region), for a given hydrocarbon chain length, have not been examined in detail. The changes in the head-group-subphase interaction may also be brought about by changing the dielectric properties21 of the aqueous subphase on the addition of small amounts (∼2 × 10-4 M) of alcohol. There are other factors that affect the nature of the π-A isotherms for the amines differently from that of the acids. For instance, head-group-head-group Coulomb repulsion effects between like charges20 are likely to be more important for less polarizable head-group cations such as -NH3+ than for anions such as -COO-. We report in this communication some aspects of our investigations on the π-A isotherms of fatty amines. We seek to highlight the way the relative strengths of the tail-tail interaction of the hydrocarbon chains and the airface-aquaface complex formation affect the nature of the π-A isotherm. It is emphasized that the π-A isotherms reported in this communication are obtained within a short time of spreading of the amphiphile. In the accompanying communication we shall discuss the long-time effects, which are considerable. Because of this time dependence, a comparison of our results with those published earlier cannot be made since the temporal parameters (such as the time of commencement of measurement, the compression rate) are not usually controlled and may vary. The results in this as well as (15) Such nomenclature is similar to but distinct from the subphase and surface nomenclature adopted by Simha and DeArmond (see ref 14d). They used the term subphase for complexes that dissolve in the subphase and surface for complexes formed at the air-water interface. (16) (a) Porter, E. F. J. Chem. Soc. 1937, 59, 1883. (b) Hoffmann, E.; Boyd, G. E.; Ralston, A. W. J. Am. Chem. Soc. 1942, 64, 498. (c) Petrov, J. G.; Kuhn, H. J. Colloid Interface Sci. 1980, 73, 66. (d) Gaines, G. L. Nature 1982, 298, 544. (e) Didymus, J. M.; Mann, S.; Benton, W. J.; Collins, I. R. Languir 1995, 11, 3130. (f) Bardosova, M.; Tredgold, R. H.; Ali-Adib, Z. Langmuir 1995, 11, 1273. (g) Angelova, A.; Petrov, J. G.; Dudev, T.; Galabov, B. Colloid Surf. 1991, 60, 351. Angelova, A.; Petrov, J. G. Langmuir 1992, 8, 213. (17) Adam, N. K. Proc. R. Soc. London A 1930, 126, 526. (18) Goddard, E. D.; Kao, O.; Kung, H. C. J. Colloid Interface Sci. 1967, 24, 297. (19) Albert, A.; Serjeant, E. P. Ionization Constants of Acids and Bases; Wiley: New York, 1932. (20) Marcus, Y. Ion Solvation; John Wiley & Sons: New York, 1985. (21) Damodaran, K. V.; Merz, K. M., Jr. Langmuir 1993, 9, 1179.
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the accompanying communication are discussed in terms of the recently proposed two-state model.12,22,23 II. Experimental Section The various fatty amines were bought from commercial sources and were greater than 98% purity. The C20 docosylamine was synthesized by Mlle, Bonnier. Unless otherwise mentioned the temperature at which the experiments were carried out corresponds to the room temperature, which was between 28 and 30 °C. The pressure-area isotherms were carried out under an ambient atmosphere using a Nima trough (Model 611). The effect of humidity and atmospheric carbon dioxide at low pH was specifically studied. There was no appreciable influence for low humidities or the usual atmospheric content of carbon dioxide. A piece of filter paper was used as the Wilhelmy plate for measuring the surface pressure. The experiments were carried out with double-distilled water as well as millipore water. The π-A isotherms were close to each other in all the cases studied. There is a continuous increase of the lift-off area ALO with time especially at low bulk pH even when the content of atmospheric carbon dioxide is reduced considerably. The isotherms reported here are those in which the compression was commenced within 20 min of spreading (see legends of figures for the exact time).
III. Results III.1. Stabilization of Langmuir Monolayers of Fatty Amines at Low Bulk pH. Adam, in his pioneering study,17 sought to control the bulk pH by using various buffer mixtures consisting of various anions. We have not used this approach as such because the nature of the acid counterion influences the isotherm and is a variable parameter, as we shall see later. The central feature of the π-A isotherms (20 min after spreading) as the bulk pH is increased (bulk pH > 5) is the continuous decrease in the lift-off area, ALO, and an increase in the slopes of the π-A isotherms. At bulk pH ) 10 the isotherm resembles that of fatty acids with an extrapolated area of ∼20 Å2, which is close to the cross-sectional area, Aperp, of the hydrocarbon chains. The pressure-area isotherms in this case resemble that reported by Lyons and Rideal24 for C18NH2 spread on an aqueous subphase of high bulk pH by the addition of NaOH. The monolayer is unstable on an aqueous solution of hydrochloric acid of low bulk pH ( H2SO4 ∼ H3PO4
(1)
Addition of ∼2 × 10-4 M n-butanol in the subphase of perchloric acid solution of bulk pH ) 2 causes an increase in ALO and a decrease in the slope of the pressure-area isotherm at low areas. However, the addition of 2 × 10-4 M concentration of methanol in the subphase does not change significantly the pressure-area isotherm. The larger solubility of methanol in water compared to that
Figure 7. Pressure-area isotherms of myristic acid spread on an aqueous solution of hydrochloric acid at a continuous compression rate of 10 (Å2/mol of acid)/min: c14ac11 (full line), bulk pH ) 2; c14acpa4 (broken line), pH ) 4; c14acb11 (dots and dashes), pH ) 2, 2 × 10-4 M n-BuOH in aqueous subphase.
Figure 8. Pressure-area isotherms of myristic acid spread on an aqueous solution of hydrochloric acid at bulk pH ) 6: (full line) without n-butanol in the subphase (5 min after spreading); (dotted line) without n-butanol in the subphase 30 min after spreading of fresh monolayer; (dots and dashes) 5 h after injection of n-butanol into the subphase to give a final concentration of 2 × 10-4 M n-BuOH. Compression isotherms are shown. Continuous compression rate of 10 (Å2/mol of acid)/ min.
of n-butanol apparently reduces the extent of segregation of the alcohol at the air-water interface. III.1.4. Studies with Fatty Acids. The pressure-area isotherm of myristic acid at bulk pH ) 2 and an ambient temperature of 26 °C is shown in Figure 7. The isotherm agrees rather well with that reported earlier,26 especially with respect to the features at the area, ALE/LC, which signals the classical liquid-expanded/liquid-condensed (LE/LC) transition. In the presence of 2 × 10-4 M n-butanol in the aqueous subphase (measured surface tension ∼62 dynes/cm) there is a marked change in the pressure-area isotherm (bulk pH ) 2, temperature ) 26 °C). The lift-off area, ALO, shifts to higher areas while ALE/LC shifts to lower values compared to that in the absence of alcohol. The pressure-area isotherm of myristic acid at bulk pH ) 4 is also shown in Figure 7. The effect of small amounts of n-butanol in the subphase at pH ) 2 seems to be similar to that of increasing bulk pH as far as the LE/LC transition is concerned. The pressure-area isotherms show a decrease in area with time when myristic acid is spread on double-distilled (26) Adam, N. K.; Jessop, G. Proc. R. Soc. London, Ser. A 1926, 112, 364; J. Am. Chem. Soc. 1939, 61, 1182.
π-A Isotherms of Fatty Amines. 1
water (pH ) 6; Figure 8), indicating a “dissolution” of the acid into the bulk. At longer times after spreading, the pressure-area isotherm does not show the “kink” characteristic of the LC/LE transition (Figure 8). On injecting n-butanol into the subphase (to make a 2 × 10-4 M solution), there is a reexpansion of the pressure-area isotherm with time for a given area. After about 5 h, a stable pressure-area isotherm is obtained, which is reproducible even after 90 h (Figure 8). The value of ALO is now very much larger compared to that obtained in the initial stages in the absence of alcohol in the subphase. IV. Discussion IV.1. Equilibria at the Air-Water Interface. The above studies show that any model for the surface pressure, π, at large molecular areas typical of the LE phase (such as those which deal with chain defects alone) and which does not take the head-group-subphase interaction into account may be considered to be incomplete. The equilibrium at the airface-aquaface interface may be written as
(1 - x)RNH2(airface) + (z - x)H+(bulk, solvated) + (z - y)X-(bulk, solvated) T (y - y′)RNH3+(airface)X-(aquaface, solvated) + yRNH3+X-(bulk, solvated) + (x - y)RNH3+(airface) (2) where z is the initial bulk concentration of the acid. y e x, to take into account the presence of (x - y) moles of another anionic species, X′-, present at the interface. X′may, for example, be OH- so that RNH3OH (the equivalent of basic NH4OH) or the hydrated RNH2·H2O may be formed. In effect, this takes into account the possibility (as we shall discuss later) that there may be a significantly large number of amine molecules that are not complexed with the acid even if the pKa of the amine is high and even if the bulk pH is typical of highly acidic character. The term, yRNH3+X-(bulk, solvated), in the right in eq 2 incorporates the possibility of dissolution of the acidbase ion pair in the bulk. RNH3+(airface)X-(aquaface, solvated) is the acid-base ion pair or complex at the interface that is of interest. Both the airface (RNH3+) and aquaface (X-) ions in the ion pairs are likely to be solvated. Equation 2 takes into account the possibility that all the amine molecules at the airface need not be protonated at low bulk pH. We consider a two-state model12,22,23 for the monolayer in which there are “individual” molecules (contributing to liquid-expanded-like behavior) and clusters of close-packed molecules (liquid-condensed-like) in an airface equilibrium
(1 - x)[RNH2}M(clusters) T (xM)RNH2(individual) (3) in the monolayer. The increase with time in the slope of the pressure-area isotherms of the C18NH2 (after normalizing for ALO) spread on HCl solution at pH ) 2 (Figure 2) suggests that “individual” molecules are likely to be “dissolved” first into the subphase. The equilibrium involving the “dissolution” of the complex
(y - y′)RNH3+(airface)X-(airface, solvated) T y′ (RNH3+)X-(bulk, solvated) (4) shifts to the right as the bulk solvation energy of the anion increases for a given hydrocarbon chain length. For a given subphase acid, eq 4 is expected to be shifted to the left with increasing hydrocarbon chain length.
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IV.2. Subphase Properties. IV.2.1. Acidity and Hydration Energies. The primary amines with pKa ∼ 10 are stronger bases than the anilines (pKa ∼ 4.5). We have used mainly strong acids whose bulk pKa values are expected to decrease in the order18
H2SO4- (pKa ) 3) > HCl- (7) > HClO4- > HI (11) (4) The values of the pKa of these acids at the air-water interface is expected to be different from that in the bulk although the above order main remain. The standard enthalpies of hydration, ∆H°hydr (see Table 5.10 of ref 21), of the ions in the bulk are (values in brackets are in kJ mol-1)
Cl- (1403) > HSO4- (1383) > I- (1340) > ClO4- (1326) < SO42- (3326) (5) IV.2.2. Sizes of Acid Counteranions. The actual sizes of the ions are difficult to evaluate, especially when they are hydrated. The “absolute” molar volumes (cm3 mol-1; given in brackets at 25 °C; see ref 21, Table 5.8) of some of the anions used in this study decrease in the order
Cl- (24) < I- (36) ∼ H2PO4- (36) < HSO4- (42) < ClO4- (50) (6a) from which the “absolute” radii, rabs, and cross-sectional areas, Aabs, are (given in brackets in Å and Å2)
Cl- (2.11, 14) < I- (2.42, 18) < H2PO4- (2.4, 18) < HSO4- (2.55, 20.5) < ClO4- (2.7, 23) (6b) while the hydrated volumes (for hydration numbers given in square brackets) vary as
Cl- (94) [3.9] ∼ I- (93) [2.9] ∼ ClO4- (100) [2.6] (7a) Although the hydration volumes are nearly the same, the hydration numbers decrease with increasing size. The corresponding values for the “hydrated” radius, rhydr, and area, Ahydr, are close to each other, being given by
rhydr ∼ 3.4 Å
Ahydr ∼ 36 Å2
(7b)
The ratio of Ahydr/Aperp ∼ 2 for the hydrated monovalent anions ClO4- or HSO4-, while for the divalent SO42- ion (see ref 21; Table 5.8, molar volume at infinite hydration ) 123 cm3 mol-1) the ratio of Ahydr/2Aperp ∼ 1.05. From these considerations one may expect, first of all, that Ahydr/ 3Aperp < 1 for trivalent anions such as PO43- (even if it exists). Since the SO42- ion has a larger enthalpy of hydration than the HSO4- ion (see eq 7) it is likely to be more stable in the “bulk”. One may expect that at large areas (area per mole of amine, Amol > 40 Å2) only the singly charged form of the anion (e.g., ClO4-, HSO4-, H2PO4-) complexed to the charged amine group at the airface. We may define dmin as the minimum distance between the acid-base ion pair RNH3+(airface)X-(aquaface, solvated), below which the attractive interaction energy overcomes any kinetic energy of random thermal motion. It is known20 that for monovalent ions at room temperature,
dmin ≈ 28 nm
(8)
in the continuum limit. dmin for monovalent species is
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Ganguly et al.
expected to be close to 3.4 Å in bulk water (dielectric constant ∼ 80 at room temperature) so that dmin < 2rhydr (see eq 7b), the separation between the hydrated anionic species. In this case, as the area per amine molecule, Amol, is decreased, contact ion pairs (see ref 21, p 221) that constitute a dipole may be formed. However, the value of the dielectric constant at the airface-aquaface interface in the presence of the amphiphile and the acid counterion is likely to be considerably less than that of bulk water. IV.3. Head-Group-Head-Group Interactions. IV.3.1. Airface Interactions. Damodaran and Merz21 find through molecular dynamic simulation experiments, that at distances of the order of 5.5 Å, repulsive electrostatic interactions between charged -NH3- groups become important. Such repulsive head-group-head-group interactions are therefore expected to take place when the area becomes less than a characteristic value, AH+ (∼(5.5 Å)2 ∼ 30 Å2), for the given system. Experimentally, the pressure-area isotherm for the C20NH2/HCl system becomes steep around 30-35 Å2/amine (Figure 2). IV.3.2. Aquaface Interactions. The size of the hydrated anion, Ahydr, in the bulk is nearly 40 Å2/anion, which is close to the estimate of AH+. It is then not clear whether the increase in energy calculated by Damodaran and Merz21 for protonated amine layers at a N‚‚‚N distance less than ∼5.5 Å is due to Coulomb repulsion between -NH3- head groups or is a consequence of the dehydration of the acid counteranions at low areas. In the latter case the acid counterions at the aquaface may be driven into the “bulk” in order to remain fully hydrated. The -NH3+ head group would now be deprotonated to conserve charge neutrality. In either case, the limiting area, AH+ ()35 ( 5 Å2/mol) seems therefore to be a realistic protonationdeprotonation limit for fatty amines at low pH. The surface pressure π is expected to be a maximum when Amol is close to AH+ (∼35 ( 5 Å2/mol). As the pH is increased or the pKa of the amino group is decreased, the equilibrium in eq 2 or 3 is shifted to the left with a decrease in ALO as well as the extent of dissolution because of deprotonation/ decomplexation of the head group. Coulomb repulsive interactions between the protonated head groups may not be important with amines such as 4-hexadecylaniline with low pKa so that ALO may be less than the limiting value of AH+ of 35 ( 5 Å2/mol. IV.4. Tail-Tail versus Head-Group-Subphase Interactions. IV.4.1. RT/C, Ratio of Tail-Tail vs HeadGroup-Subphase Interactions. We use a qualitative parameter, RT/C, which is a measure of the ratio of the strengths of the tail-tail interaction to that of the headgroup-subphase interactions (e.g., airface-base/aquafaceacid complex RNH3+(airface)X-(aquaface, solvated)). For example, we expect RT/C(C18NH2) < RT/C(C20NH2) for a given subphase acid. RT/C may also be varied by changing the basicity of the -NH2 group for a given length of the hydrocarbon chain. At high pH the main features of the (initial) pressure-area isotherms of the amines (such as the extrapolated area per molecule) are expected to be rather independent of RT/C or the length of the hydrocarbon chain. The initial (short-time after spreading) values of ALO of C20NH2 spread on various subphases (Figure 5) seem to increase in the order
HCl < H3PO4- ∼ H2SO4- < HClO4
(9)
although the size of the hydrated singly charged anions of HI, HCl, H3PO4, H2SO4, and HClO4 are nearly the same (eq 6). ALO is thus not determined by the size of the anion.
The sequence in eq 9 is directly related to pKa (eq 4) so that ALO increases as RT/C decreases or the strength of the acid-base complex at the air-water interface increases. One could also conclude from eq 5 that a high solvation energy of the anions drives the acid counteranion to the bulk. In this case the sequence in eq 9 is due to a decomplexation at the interface with the equilibrium in eq 2 being shifted to the left. The hypothetical pressure-area isotherm of C18NH2 on an aqueous subphase of hydrochloric acid solution (at 28 30° C) with pH ) 2 is not expected to be very different from that shown in Figure 2. From Figures 2, 4, and 5 we find that ALO values of the amines decrease in the sequence
ALO(C18NH2)(∼50 Å2) > ALO(C20NH2)(∼35 Å2) > ALO(C16PhNH2)(∼25 Å2) (10) on an aqueous hydrochloric acid subphase (pH ) 2) so that ALO increases as RT/C decreases. The slope of the pressure-area isotherms decreases with increasing RT/C with
C16PhNH2 ∼ C20NH2 > C18NH2
(11)
at a given (short) time after spreading. The presence of n-butanol (Figure 6) influences the nature of complexation and is equivalent to a decrease in RT/C since ALO increases. The charged acid/base ion pair (RNH3+(airface)X-(aquaface, solvated)) formation at the air-water interface is favored as a consequence of a decrease in the dielectric constant at the aquaface, which would cause an increase in dmin (eq 8) and thereby favor ion-pair formation. IV.5. On a Possible Two-State Model for Langmuir Monolayers of Fatty Compounds. The high-area region in the case of fatty amines spread on an acidic subphase is likely to be similar to an LE phase although the isotherms presented in this communication may not actually show any evidence for an LE/LC coexistence region. A two-state model has been proposed by Israelachvili22 and further developed by Ruckenstein and Li12,23 for the LE/LC transition of fatty acids. The latter have proposed that the insoluble surfactant monolayer at the air-water interface is made up of singly dispersed or clusters of ordered (all-trans) surfactant molecules and disordered (with gauche defects in the hydrocarbon chain) individual surfactant molecules. The “gauche” defects cause tilting of the hydrocarbon chains and an increase in area per molecule. The model of Ruckenstein and Li12 does not take into account the influence of any headgroup-subphase interaction on the energetics of the monolayer. This is inconsistent with our observations. The airface-aquaface interactions affect the chemical potential of the molecule whether in the isolated state or in clusters. The equilibrium constant (see Israelchavali22) for the isolated molecule (e.g., in eq 2) would depend on the nature of these head-group-subphase interactions. The high-area liquid-expanded state is predominantly driven by head-group-subphase interactions. This is suggested from the dependence on the nature of the subphase acid counterion as well as the experiments with myristic acid monolayer spread on an aqueous subphase of high pH (Figure 8), in which the high-area “individual” molecules are “dissolved” first leaving the “clusters” at the surface (see section III.1.1). The distinction between individual molecules that are electrically bound to subphase monovalent ions including
π-A Isotherms of Fatty Amines. 1
protons and “condensed” state appears in the study of Bloch and Yun.11 According to these authors, the divalent metal ions are condensed onto the carboxylate groups in the monolayer while the monovalent ions, including protons, are electrically attracted to the monolayer. Bloch and Yun11 could not reproduce the LE/LC phase transition in the Poisson-Boltzmann-Stern formalism. Our results suggest that a connection between their two-state model
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involving individual species electrostatically bound to monovalent subphase ions and condensed states could be made. Acknowledgment. We are thankful to the CEFIPRA Indo-French program for funding. LA960981M
