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Universal Functional Group Activity Coefficient Model in Electrosorption. 3. Thermodynamical Excess Mixing Functions of the 1-Propanol-Water System Adsorbed at the Hg Electrode Marian Karolczak Universytet im. Adama Mickiewicza, Wydziat Chemii, Zakhd Chemii Fizycznej, ul. Grunwaldzka 6, 60-780 Poznafi, Poland Received April 6, 1989 Electrosorption of 1-propanol at the Hg electrode from aqueous sodium sulfate, of constant salt activity a* = 0.07087, at the temperature 298.15 K at various electrode potentials has been investigated. Excess electrochemical Gibbs energy of mixing, excess electrochemicalenthalphy of mixing, and excess electrochemical entropy of mixing of the adsorbed 1-propanol and water in the surface solution have been determined. These excess mixing functions have been shown as functions of concentration, temperature, and electrode potential. The excess partial molar quantities additionally supplement the_presentation. Thermodynamical implications of the obtained-results lgad us to the conclusion that PHE # 0 and A@ # 0 as well as in a diluted solution region TIASEI > IhHEI, answering thus important questions of adsorption modeling raised 12 years ago by Mohilner et al. The results have been qualitatively interpreted in the terms of structure breaking-structure making balance in the surface solution providing some physical insights in the interpretation of the interfacial region structure. It has been shown that this balance strongly depends on the electrodepotential. In particular, excess electrochemical enthalpies and excess electrochemical entropies of mixing, as well as their partial molar equivalents, show a significant dependence on electrode potential leaving excess electrochemicalGibbs energy of mixing and excess electrochemical potentials of the adsorbed 1-propanol and the adsorbed water rather insensitive in this respect. It has been found that for the investigated system there is no simple correlation of surface and bulk behavior, which could be expressed in terms of the thermodynamical excess mixing functions.
Introduction The necessity of complex thermodynamical studies for solving some fundamental problems in electrosorption has been noted and emphasized in the earlier works of Mohilner et a1.l and Nakadomari et aL2 It has been particulary interesting for these authors to find whether or not afiE = 0, i.e., whether the mixing in surface solution is an athermal process or, alternatively, whether excess entropy of mixing equals zero, ASE= 0, and if not so, to establish whether the properties of surface solution are under “entropy control” (Le., whether TIAsEl > I@(. This ambitious goal had stimulated very demanding experimental p r ~ c e d u r e ~and - ~ remarkable improvements of the precision of the interfacial data measurements. Thus Mohilner and Kakiuchi? Krishnan and de Levie,G and also Saffarian snd de LevieZ6achieved a higher, exceptional a t the present time, level of the measuring technique. Almost simultaneously, theoretical foundations of the electrosorption of organic compounds have been and subsequently further developed.lO-14 In (1) Mohilner, D. M.; Nakadomari, H.; Mohilner, P. R. J.Phys. Chem. 1977,81, 244. (2) Nakadomari, H.; Mohilner. D. M.: Mohilner. P. R. J.Phvs. Chem. 1976,80,1761. (3) Mohilner, D. M. In Electroanalytical Chemistry; Bard, A. J., Ed.; Marcel Dekker: New York. 1966: Vol. 1. DO 241-409. (4) Mohilner, D. M.; Nakadomari, H.*3. Electroanal. Chem. 1975, 65,843. ( 5 ) Mohilner, D. M.; Kakiuchi, T. J. Electrochem. SOC.1981, 128, 350. (6) Krishnan, M.; de Levie, R. J. Electroanal. Chem. 1982, 131, 97. (7) Karolczak, M.; Mohilner, D. M. J. Phys. Chem. 1982,86, 2840. (8) Karolczak, M.; Mohilner, D. M. J. Phys. Chem. 1982,86, 2845. (9) Mohilner, D. M.; Karolczak, M. J . Phys. Chem. 1982,86, 2838. (10) Karolczak, M. J. Colloid Interface Sci. 1984, 97, 284.
the next stage, specific objects for crucial studies have been selected. They comprised comprehensive measurements of the interfacial tension a t the mercury/lpropanol, aqueous sodium sulfate (of constant activity) interface, at various electrode potentials and temperatures. The results have already been reported for the temperature 293.15 K.15 The purpose of the present paper is to supplement previous investigations with the results at standard temperature, 298.15 K, together with some results which follow from the temperature dependence, i.e., the excess electrochemical enthalpy of mixing and excess electrochemical entropy of mixing in the surface solution at standard temperature. Thus, we can provide also the answer to the principal questions set in 1977 by Mohilner et al.1 Experimental Section Reagents, solutions, purity, experimental procedure, hardware, measuring technique, data eleboration, etc., have been described in detail in a previous p~b1ication.l~The results reported in this paper follow them in all respect, except the temperature, which is 298.15 K, and the modifications of weighing factors in the final fitting procedure, as described below. The original interfacial tension data have been collected at the Department of Chemistry, Colorado State University, by using a computer-controlled capillary electrometer. Data Analysis Calculations from the experimental data of interfacial tension, -In aads,values and calculated surface mole frac(11) Karolczak, M. J. Electroanal. Chem. 1984, 181, 21. (12) Karolczak, M. Electrochim. Acta 1985,30, 325. (13) Karolczak, M. Electrochim. Acta 1986,31, 1177. (14) Karolczak, M. Electrochim. Acta 1987, 32, 79. (15) Karolczak, M. Langmuir 1990, 6, 863.
0 1990 American Chemical Society
1046 Langmuir, Vol. 6, No. 6, 1990
Karolczak
tions, ~ A a d s ,have been fitted by a nonlinear leastsquares method, referred to in ref 15, to the equation
+
-ln a;ds = -In (1- X L ~ ~ ) In (A~X;~'
+ I -x;~')
-
electrochemical Gibbs energy of mixing from the equation
AGE/RT = - ~ X A ' & In + (1 - XA'~')/AZ]~ X A ' ~ ln ' [XA'~' + (1- XA'~')/A~]- (1- XA'~') x In [AIXAadS+ (1 -xAads)] + 3 X A a d s h [(I + 3xAads)/4] (3) derived in ref 16 and the electrochemical enthalpy of mixing from the equation
derived in ref 16, providing parameters AI, Az, and A3 and the estimations of their confidence intervals. The values of -In a@& have been calculated, from experimental data according to ref 17, by using the equation 0
-In awads - -- - y In (awlawe) nRTI?, where yo and y are interfacial tension without and with adsorbing matter, respectively; aw and awe are activities of water in bulk and supporting electrolyte, respectively; n is the stoichiometric coefficient in the adsorption process; and rmis the maximum, theoretical value of surface concentration. During extensive studies, we have realized that the numerical values of the fitting parameters depend (to some extent) on the form of the equation used for the fitting, more precisely, on the coordinates in which the fitting was made. Therefore, we have decided to modify our original fitting technique by incorporating in the weighting factors the errors resulting from the possible variance of the independent variable, The modified procedure proved to be successful and especially efficient in damping of the discrepancies of the estimated parameters resulting from the differences in the fitting ranges. It provided also the solution to the problem of the self-consistency of the estimated adsorption isotherm parameters, raised in ref 11. We have recalculated our previous datal5 accordingto the modified weighting procedure to make them coherent with the present data and use them in our further analysis. The magnitude of the weighting factors wi varied significantly with ads, and the equation used was 1/wi2 = 1.39013 X
+ 2.13689 X 10m7(l+ 3xtds)*X
The numerical constants in this equation result from admitted errors of the interfacial tension, viz., Ay = i O . O 1 pJ C ~ - ~ the , ~ relative O surface excess AI'Aw N A r A = 10-l2 mol cm-2, and the assumed values of molecular constants n = 4 and rm= 5.4082 X mol cm-2. Contrary to our previous estimation^,^' we have decided to take a constant value as the possible error of AI'A and estimate it as the upper limit of the error of the tabulation as provided by the computer printout of the program we used. Comparisons with ref 17 have shown that this is an overestimation. The partial derivative
can be calculated analytically from eq 1. Having determined parameters AI, Az, and A3 at different electrode potentials and temperatures, we could calculate excess (16)Karolczak, M.J.Electroanab Chem. 1989,262,263. (17)Karolczak, M.J . Phys. Chem. 1985,89,1556.
Simultaneously, excess electrochemical entropy of mixing has been determined from the thermodynamical relation
TASE = @ - A G E (5) Since the temperature range was rather narrow (5 deg), we have approximated the derivatives with respect to the temperature in eq 4 by respective finite differences AAi/ AT. In a similar manner, all partial molar excess functions have been calculated. Although original interfacial tension data have been measured in the range f0.6 V with respect to the potential of maximum adsorption, we decided to analyze more carefully and present in this paper only the data from the (arbitrarily selected) range i 0 . 3 V. Our decision in this respect resulted from consideration of the following factors: (1)It follows from Payne's paperla that for potentials more positive than +0.3 V an appreciable amount of sulfate ions may coadsorb. We have estimated from Payne's datal8 surface concentration of specifically adsorbed sulfate ions: I'so,z-/mol cm-2 equals 3.11 X 10-l' (at +6 gC c m 3 , 2.07 X lo-" (at +4 g C cm-2), and 1.03 X 10-l' (at +2 KC cm-2). The extreme left end of Payne's graph (Fig ure 11)corresponds to the activity at = 0.07078 (i.e., the same value we consequently kept constant in our investigations), and Payne stated: "As frequently found for inorganic ions the specifically adsorbed charge becomes asymptotic to the value of the electrode charge at low electrolyte concentrations". Since we intended to interpret the results in terms of binary mixtures, to avoid the complications in the calculations and the interpretation due to eventual coadsorption of sulfate ions, we had to bind the analysis to the range where the coadsorption can be reasonably neglected. (2) For the potentials more negative than -0.3 V, coadsorption of sodium ions cannot be excludedla-it follows from the deficiencies of the diffuse layer theory and the observed noncoincidence of the electrocapillary curves2 as well as other findings.'s Thus, one can expect somewhat similar problems as noted above. (3) Additionally, for the potentials outside the f0.3-V range a progressive desorption of 1-propanolbegins. Thus, only within the maximum adsorption region is there a sufficiently large amount of 1-propanol present in order to neglect effects mentioned above and assure a uniform error distribution and subsequently the same statistical significance of the fittings. (18)Payne, R.J. Electroanol. Chem. 1975,60,183.
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Thermodynamical Excess Mixing Functions
Results and Discussion Before examining surface behavior of aqueous solutions of 1-propanol, we should recall the most important features of aqueous solution of the related compounds in the bulk. Although water and alcohols resemble each other in hydrogen bonding, it is well established that the structures to which these bonds lead are very dissimilar and are not compatible. Thus, it has been widely accepted that liquid water structures are quasi-crystalline, tridimitelike clusters (of fickering structures in some theories) which are in equilibrium with the monomeric molecules. At elevated temperatures, close-packed structures have been postulated. One of the most popular along this line is the “iceberg” concept of Frank and Evans.lg In spite of the fact that the “iceberg” model cannot quantitively interpret all anomalous properties of liquid water, it created the basis of common convincing that liquid water in reality, even at room temperatures, is in fact highly structured. Pure alcohols are also to a great extent associated liquids although not as highly structured as water; dimers, trimers, tetramers, and even cyclic polymeric structures have been successfully ad~anced~’-~~-the latter explained some peculiar physical properties, e.g., abnormal heats of evaporization and boiling point, in comparison to normal latent heats of fusion and melting point. We think that a meaningful discussion of alcohol + water in surface mixtures cannot neglect evidently competing effects of noncongruent structures of parents liquids and ignore the differences mentioned above. Enthalpy effects should be considered: when a mixture is made, some hydrogen bonds are broken endothermically, new ones are created exothermically, and excess enthalpy of mixing will be the net result of two much larger contributions. As the composition of the solution varies, the relative magnitude and perhaps even the origin of such contributions may significantly change, disabling a clear, unambiguous interpretation. Quite similary, excess Gibbs energy of mixing function can be viewed as a contribution of opposite enthalpy-entropy effects. Nonetheless, the interpretation of the thermodynamical mixing function is not much advanced by simply noting that they are the effect of the opposing contributions, and some general conclusions can be derived from their concentration, temperature, and potential dependence. For a l-propanol-water system adsorbed at the mercury electrode, the thermodynamical excess electrochemical mixing functions in surface solution are shown in Figure 1 as functions of surface mole fractions for two temperatures, 293.15 and 298.15 K, and at a constant (19) Frank, H. S.; Evans, M. W. J. Chem. Phys. 1945,13,507. (20) Rowlinson, J. S. Liquids and Liquid Mixtures; Butterworth. London, 1959. (21) Glew, C. Can. J. Chem. 1960,30,208. (22) Jeffrey, G. A. Proc. Conf. Desalination Res. 1961; Woods Hole, MA, Publ. 942, NAS-NRC, Washington, 1963; p 156. (23) McMillan, W.; Mayer, J. J. Chem. Phys. 1945,13, 276. (24) Robinson, R. A,; Stokes, R. H. J. Chem. Phys.1961,65,1954. (25) Munday, E. B.; Mullins, J. C.; Edle, D. D. J. Chem. Eng. Data 1981, 25, 191. (26) Saffarian, M. H.; de Levie, R. J. Electroanal. Chem. 1985, 189, 325. (27) Piekara, A. J. Chem. Phys. 1962,36, 2145. (28) Malecki, J. J. Chem. Phys. 1962, 36, 2144. (29) Prigogine, I. J. Chim. Phys. 1948, 45,17. (30) The original interfacial tension data, as they are tabulated by the computer, show the precision ca. AT = *0.002 LJ cm-2; however, we
have observed that the subsequent smoothing and differentiation programs, using a two-dimensional moving least-squares path,2 provide the smoothed data for which the admitted value AT was in our judgement the best estimate at the present time.
i
:.
08
-.
10
A
Figure 1. Thermodynamical excess mixing functions of l-propanol-water system adsorbed at the mercury electrode at the potential of maximum adsorption and the temperatures 293.15 and 298.15 K. Dashed lines represent the data in the bulk solution. electrode potential of +0.0142 V (maximum adsorption potential). It follows from Figure 1 that the temperature has a very pronounced effect on the excess enthalpy and excess entropy of mixing while showing a weakly perceptible effect on the excess Gibbs energy of mixing. At 293.15 K, TIASEl > I m l in the whole composition range, which confirms typically aqueous behavior20 of l-propanol, supports the earlier hypothesis of Mohilner et al.,l and thus answers their second question. It simultaneously answers their first question because # 0 and ASE # 0 (except the trivial points X A = ~0 and ~ X A = ~1). ~The~ last conditions are sufficient, according to ref 1, for eliminating Flory-Huggins and regular solution models as well as other athermal models from considerations in adsorption modeling. The following remark seems to be relevant in this context. From a strictly thermodynamical point of view, the physical interpretation of Frumkin’s isotherm (the most popular isotherm in the electrosorption studies) can be extended beyond the regular solution model: a conceivable so-called simple model in Guggenheim’s sense. Thus, the conclusion above should not be extended to such a generally understood interpretation of Frumkin model. In other words, such a generalized model cannot be simply rejected just on the basis of the above findings. This point has been overlooked in ref 1, which appears too restrictive in this respect. In the water-rich region, the mixing is exothermic, i.e., A i P 0, but as the surface mole fraction increases it first decreases, goes through a minimum, and next increases. In the mixed solution region, the mixing is endothermic, i.e., m > 0, and passes through a maximum; a part of this region at the temperature 298.15 K is under enthalpy control, Le., I m I > T l A P I ; thus, it presents typically nonaqueous behavior.20 Finally, in the alcoholrich region, the mixing is again exothermic, under entropy control, and passes through a minimum. These observations confirm a very complex nature of the excess electrochemical enthalpy of mixing, reflecting also a delicate balance of enthalpy-entropy contributions to the final values of excess Gibbs energy as well as the balance of a variety of complex interactions or bonds in the excess enthalpy. The negative entropy effect which overcomes the exothermic enthalpy effect in the water-rich region can be
~
1048 Langmuir, Vol. 6, No. 6, 1990
Figure 2. Excess electrochemical Gibbs energy of mixing in surface solution at 298.15 K plotted as a function of concentration a t several electrode potentials.
__ 03
c2
c-.
~___ 26 OCS
a
Figure 3. Excess electrochemical enthalpy of mixing in surface solution at 298.15 K plotted as a function of concentration
at several (indicated)electrode potentials. Dashed lines represent the data in the bulk solution. explained on the basis of the Frank-Evans proposition.19 According to this proposition, which also received independent support from the investigations of methane hydrate and other clathrate structures,21,22the presence of apolar solute (which we recognize here as the hydrocarbon part of the 1-propanol molecule) causes an increases in the order of the water surroundings of the solute. The authorsls labeled this the region of increased order, i.e., the region of relative large negative entropy, and gave an “iceberg” warning not to understand this term too literally. If one considers an enthalpy effect in the region where the apolar part of the alcohol molecule promotes the water structure, it is reasonable to suppose that due to the structure-making effect, mentioned above, the hydrogen bonding of water is enhanced and induces an increasing exothermic enthalpy effect. As the concentration of the solute increases, progressive breaking down of the original structure must develop resulting in the increasing endothermic effects. The net enthalpy should than pass through a minimum and again increase. This is because the ratio of the hydrogen bonds being broken to the hydrogen bonds being formed (actually no hydrogen bonds with the apolar part can be a_ssumed)is increasing; increasingly more endothermic AI-P is what we can actually observe in the plots in Figure 1.
Karolczak In the mixed solution region, addition of 1-propanol completely disintegrates the original water structure. However, it is not conceivable that l-propanol and water will exist independently in such mixtures. Thus it seems highly possible that 1-propanol-water hydrogen bonds (hydratation) will be formed. This effect will oppose an uncontrolled endothermic increase of the enthalpy (and the corresponding uncontrolled increase of the entropy), concurring with the “iceberg” structure-breaking effect considered above. As a consequence, a maximum on the enthalpy and entropy curves develops. This rather simplified picture finds its support in the temperature dependence of the position and the relative value of such a maximum observed in Figure 1. Since lower temperature promotes the “iceberg” structures, the maxima at lower temperatures are lower and appear at higher l-propanol concentrations. As classified in ref 20, I@( > ITAsEl is typical for nonaqueous behavior. Two different interpretations have been offered to explain quantitively properties of such solutions. The first is based on the McMillan-Mayerz3 theory of solutions in which solute-solute interaction is the predominant factor in the description of various mixtures. The second interpretationz4 uses the concept of hydrogen-bonded complex formation, tentatively considered above; Le., solute-solvent interaction is the predominant factor in this interpretation, and solute-solute interactions are of minor importance and are bypassed. Independent physical evidences such as NMR and dielectric relaxations support the latter approach. However, it should be noted that there is evidence for subtle inadequacies of this d e ~ c r i p t i o n . ~ ~ In the alcohol-rich mixtures, any highly ordered structures of water probably do not exist; it is commonly believed that water exists in such mixtures as a monomer and eventually associates to some extent (in slightly polar solvents). The hydration effects become gradually balanced by bonding and structural effects of the alcohol itself in favor of the latter. Thus, a second minimum, which has not been observed in the corresponding bulk mixtures, develops. The last observation indicates a specific influence of the electrical field on l-propanolwater interactions and the related structures. In a more sophisticated and quantitative description, the proton-accepting facility of various hydrogen bonds must be also taken into account together with the role played in structure making-structure breaking balance (especially in diluted solutions) by the “third”, geometrical, factor of the solute. Since we are dealing here with only one solute, we must leave these nuances for future investigations. The influence of the electrode potential, measured with respect to an appropriate (reversible to an ion of the solution) reference electrode (called an indicator electrode2J5)on the electrochemicalmixing function has been shown in Figures 2-4. Examination of Figures 2-4 reveals that electrode potential significantly influences the concentration profile of excess electrochemical entropy and excess electrochemical enthalpy of mixing but has a very small impact on the excess electrochemical Gibbs energy of mixing. Such results allow us to foresee the congruence of the electrosorption isotherm and are consistent with our previous findings for the temperature 293.15 K. We may also conclude that although the components in fact change with the potential their net effect remains almost unchanged. If we recall that the dependence of the excess electrochemical Gibbs energy of mixing on the electrode potential is rigorously related to the problem of congru-
Langmuir, Vol. 6, No. 6, 1990 1049
Thermodynamical Excess Mixing Functions
t 00
32
C'
,ads
05
08
-8 10
02
Figure 4. Excess electrochemical entropy of mixing in surface solution at 298.15 K plotted as a function of concentration at several (indicated) electrode potentials. Dashed lines represent the data in the bulk solution.
06
OL
98
10
.ad5
A
Figure 6. Excess partial molar enthalpy of mixing at 298.15 K plotted as a function of concentration at several (indicated) electrode potentials. Continuous lines refer to adsorbed l-propanol; dotted lines refer to adsorbed water.
-81
02
06
04
08
10
Yids
Figure 5. Excess electrochemical potentials of the adsorbed 1-propanol (curves running down from left to right) and of adsorbed water (curvesrunning up from left to right) at 298.15 K plotted as a function of concentration at several electrode potentials. ence of an isotherms (and the resulting problems'), the conclusion emerges that the real reasons for the apparent congruence or "almost congruent" behavior observed sometimes in the literature may be far more complicated than had been originally thought; i.e., by no means does a mere examination of the net effect provide sufficient insight to the problem. Except for the two extreme potentials, the characteristic features of the concentration shapes are conserved in Figures 3 and 4. Moreover, the influence of the electrode potential on the diluted solution, water-rich region which is under entropy control is not as significant as it is on the mixed solution region and first of all on the alcohol-rich region. It is evident from both, the excess entropy and excess enthalpy dependence. It suggests in turn that "iceberg" forms are relatively more stable than subsequent hydratation forms of 1-propanol-water complexes and perhaps 1-propanol structures. It may also indicate the enhancement of the latter and weakening of the former. The observed magnitudes of the maxima and the second minima support the interpretation along this
Figure 7. Excess partial molar entropy of mixing at 298.15 K plotted at several (indicated)electrode potentials. Continuous lines refer to adsorbed 1-propanol;dotted lines refer to adsorbed water. line. In this respect, the influence of electrode potential on the structures of the solutions differs from the influence of temperature. In Figure 1,a general trend of structure breaking seems to follow for both variables. It is well-known for the bulk solutions and can be confirmed from Figure 1 that temperature has a much smaller effect on the pure alcohol than on the structure of water and diluted solution region. A distinct departure from the described behavior begins around *0.3 V of the potential of the adsorption maximum. For these potentials, the mixing is endothermic in the whole composition range. The same trend is reflected in the positive entropy values. These observations confirm a hypothesis made above that both water and 1-propanol exist in completely disintegrated structures, most likely as monomeric, repulsive species. It is interesting to note that both considered potentials almost coincide with the beginning of adsorption-desorption peaks on the differential capacitance curves. For comparison, dashed lines in Figures 2, 3, and 4 represent the corresponding thermodynamicalmixing functions in the bulk solutions. They were calculated on the basis of data and equations published in ref 25. The results
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presented in Figures 2-4 suggest that the correlation of the bulk and interfacial behavior in the terms of thermodynamical functions appears to be not particularly usefull, if promising a t all. From general theoretical considerations, some of the excess electrochemical partial molar quantities are more sensitive in reflecting the concentration dependence of various structures in the solution than are the integral quantities. Therefore, we included in this paper the results of calculations concerning excess electrochemical potentials, excess electrochemical partial molar enthalpies of mixing, and excess electrochemical partial molar entropies of mixing for the adsorbed water and the adsorbed 1-propanol. They are presented in Figures 5-7. Continuous lines in Figures 6 and 7 refer to 1-propanol; dotted lines refer to water adsorbed at the mercury electrode. The results presented in Figures 5-7 confirm the general conclusions concerning the structure of 1-propanol-water surface mixtures derived above. Electrode potential has very pronounced effect on the excess electrochemical partial enthalpies and entropies of mixing of 1-pro-
panol, particularly when the concentration of the latter is low; its effect on the corresponding partial molar properties of water is much smaller. In conclusion, we note that both the above description of 1-propanol-water structures and presented shapes of the excess enthalpies and entropies are derived from the experimental data which are only indirectly related to the changes of structures. The initial thermodynamical rigor of the experimental data is lowered because of approximations we had to introduce in our calculations (c.f. Experimental Section) and especially because of a linearization of the temperature dependence of the fitting parameters. However, since the errors of these parameters, estimated a t the confidence level of 95 5% , fluctuate, 0.152% of A1, 0.01-0.12% of A2, and 0.03-0.3% of As within the investigated electrode potential range, we believe that the calculated data represent real data in a faithful manner. Registry No. PrOH, 71-23-8; Hg, 7439-97-6; HzO, 773218-5.
Vibrational Characterization of Langmuir-Blodget t Monolayers and Evaporated Films of a Bisphenylene-Substituted Perylene-3,4,9,10-tetracarboxylic Acid Derivative P. Aroca, Jr., and R. Aroca Department of Chemistry & Biochemistry, University of Windsor, Windsor, Ontario, Canada N9B 3P4
G. J. Kovacs* and R. 0. Loutfy Xerox Research Centre of Canada, 2660 Speakman Dr., Mississauga, Ontario, Canada L5K 2L1 Received September 28, 1989. I n Final Form: December 18, 1989 The vibrational spectra of a bisphenylene-substituted perylene-3,4,9,10-tetracarboxylicacid derivative (BPTCA) in thin solid films and Langmuir-Blodgett (LB) monolayers have been studied with IR and Raman spectroscopy. The molecular spectra of BPTCA evaporated at monolayer coverage were obtained by using surface-enhanced resonant Raman scattering (SERRS) on Ag-coated Sn spheres. Evaporated BPTCA films did not show any signs of preferred molecular orientation and were physisorbed onto the Ag surfaces. Stable LB layers of BPTCA mixed with arachidic acid were prepared and transferred to glass and to glass coated with the SERS active surface. An edge-on molecular orientation of BPTCA in the mixed monomolecular layer was identified from the SERRS spectra on Ag; the spectra indicated chemisorption and a well-defined metal-molecule interaction. SERRS of BPTCA detached from the surface was also obtained by inserting an LB spacer layer of fatty acid between the BPTCA and the Ag surface; evidence for the chemisorption disappeared in the spectra of the detached layer. Introduction
BPTCA is one of the PerYlene imidazole derivatives that is photoactive in the visible region and thereby a suitable candidate for xerographic applications using vis-
* Author t o whom correspondence should be directed.
ible 1ight.l Tang2 has incorporated the BPTCA together with copper phthalocyanine in the development of a photovoltaic cell based on a two-layer structure of organic films. In the latter work, it was shown that the interface between (1)Popovic, Z. D.; Loutfy, R. 0.;Hor, A. M. Can. J . Chem. 1986,63, 134. (2) Tang, C. W.; A p p l . Phys. Lett. 1985, 48,183.
0743-7463/90/2406-1050$02.50/0 0 1990 American Chemical Society
