0 Copyright 1986 American Chemical Society
SEPTEMBER/OCTOBER, 1986 VOLUME 2, NUMBER 5
Art ic 1es Effect of Pressure on the Adsorption of Tris (2,2’-bipyridine)ruthenium(2+) from Solution Mary E. Zawadzki and Arthur W. Adamson* Department of Chemistry, University of Southern California, Los Angeles, California 90089-1062
Monty Fetterolf and Henry W. Offen Department of Chemistry, University of California-Santa Barbara, Santa Barbara, California 93106 Received March 6, 1986 The system Ru(bpy),2+-Ti02(anatase) (bpy = 2,2’-bipyridine) shows reversible, Langmuir-typeadsorption from dry acetonitrile solution. The maximum adsorption and the b constant of the Langmuir equation both decrease with increasing hydrostatic pressure with respective coefficients of -9.8% and -32% per lo00 bar. The effect of pressure is thus to decrease both the apparent adsorption capacity and the strength of adsorption. If adsorption is simply treated as a chemical equilibrium, the application of simple thermodynamics gives a AV value of 7.7 cm3 mol-’ at 25 O C for the adsorption process. The significance of this positive value is discussed. We believe these observations to be the first reported ones on the effect of pressure on adsorption from solution.
Introduction The first-named authors have been interested in the phenomenon of adsorption from solution onto both powdered solids1 and smooth metal surfaces.2 Particular emphasis has been on the study of cases of irreversible adsorption, that is adsorption which gives normal appearing (Langmuirian) adsorption isotherms but with desorption nil or very slow and incomplete on dilution with the same solvent. The effect is not due to chemical bond formation or change; desorption in such systems is typically rapid and complete on changing to a better solvent, and moreover, typical calorimetric heats of adsorption are small, in the 2-4 kcal mol-’ range. In the course of a study with Ru(bpy)32+(bpy = 2,2’bipyridine) on Ti02 (anatase), we found the adsorption to be irreversible in the above sense if occurring from aqueous (1) Namnath, J. S. Ph.D. Dissertation, University of Southern California, Los Angeles, Jan 1983. (2) Tamura, K.;Tse, J. T.; Adamson, A. W. J. Jpn. Pet. Inst. 1983, 26,309-317. Tamura, K.;Tse, J. T.; Adamson, A. W. J. Jpn. Pet. Inst. 1984,27, 385-391.
or from aqueous 1%’ acetonitrile (AN)solutions. Since the adsorption was, however, entirely reversible on dilution of the system if pure AN was the solvent, it was of interest to investigate the effect of high pressure on this system. For the reversible case, a t least, nominal use of conventional thermodynamic relationships should be allowed. It was of additional interest that a full search of Chemical Abstracts revealed no prior publication on the effect of high pressure on adsorption from solution. There was thus the possibility of developing a new phenomenology. This initial report is based on work carried out collaboratively a t the high-pressure laboratory of the last-named author. (The effect of pressure on electrode processes has been reported by Conway et al.9 Experimental Section The high-pressure equipment is that described elsewhere3and could be used for the present experiments without modification. Initial and final solution concentrations were measured spec(3) Dawson, D. R.;Offen, H. W. Reu. Sci. Znstrum. 1980, 51, 1349.
0743-7463/86/2402-0541$01.50/00 1986 American Chemical Society
542 Langmuir, Vol. 2, No. 5, 1986
Zawadzki et al.
I
2
4
to4 c ,
6 M
8
(o4c, M
Figure 1. Effect of pressure on the adsorption of Ru(2,2'-bpy)2+ by TiOBfrom AN solution at 25 "C. (a) Effect of increasing pressure on the adsorption isotherms. (b) Data plotted according to the linear form of eq 1.
trophotometrically, after allowing the powdered adsorbent to settle, which it did cleanly. The TiOpwas prepared from a hydrous slurry supplied by Gulf & Western Natural Resources Group, Division of Gulf & Western Industries, Inc. The slurry was dried under mild conditions to remove free water and then progressively to a final treatment at 110 "C for 4 h. BET adsorption measurements using N2 gave a specific surface area, 2 , of 139 m2 g-'. The Ru(bpy),2+was recrystallized as the chloride salt and its UV-visible absorption spectrum agreed to within 1%with the published Dry AN solutions were made up in a moisture-free glovebag, and the mixture of solution and adsorbent were loaded into the highpressure cell under moisture-free conditions. The general procedure was to take an absorption spectrum of the equilibrium system under atmospheric pressure and then to take a succession of spectra of the system equilibrated at 25 " C under increasing applied pressure at room temperature. The sequence was then repeated with a different initial concentration. All measurements were in duplicate. Reversibility was verified within 5% by measuring the solution concentration before and after each pressure experiment. For the highest concentration points, it was necessary to use tandum neutral density filters in the reference beam of the spectrophotometerand this may have introduced a small systematic error.
The slopes and intercepts in Figure l b allow calculation of n, and b. The values of the former are 6.31 X lo", 5.34 X and 4.83 X and the values of b are 3.32 X lo4, 1.97 X lo4, and 1.31 X lo4 for 1atm, 1500 bar, and 3000 bar, respectively. The n, values show some curvature if plotted against pressure but give a more linear semilogarithmic graph; this plot has a slope corresponding to 9.8% decrease in n, per 1000 bar. Taking Ru(bpy),2+ to be spherical with a radius of 7.0 A, the n, value corresponds to an effective 2 of 59 m2g-' at 1atm and 45 m2g-l at 3000 bar. It appears that the BET surface area at 1atm is about half-occupied and occupancy decreases with increasing pressure. Both the effect of pressure on n, and the discrepancy with the BET value could be accounted for if the adsorbate molecule flattens somewhat in the adsorbed state, the more so with increasing pressure. Also, the adsorbent should be nonporous, but SEM photographs show aggregates of very small particles and there may be areas accessible to N2 in the BET measurement but not to R ~ ( b p y ) , ~ + . The b values also show an essentially linear behavior if plotted semilogarithmically against pressure, the slope corresponding to a 35% decrease per 1000 bar. Since b is assumed to be a measure of adsorption strength, the simple conclusion is that the adsorption interaction is weakened with increasing pressure. The qualitative behavior of both n, and b is evident in Figure l: both adsorption capacity and adsorption strength decrease with increasing pressure. Since we are dealing with adsorption from solution, solvent must be displaced as solute is adsorbed, and the adsorption process might better be written6 A(so1ute in solution a t C) + B(adsorbed solvent, N l s ) = A(adsorbed solute, N z S )+ B(solvent in solution, N,) (2) where N," and N,"are the mole fractions of adsorption sites occupied by solvent and solute, respectively. Because the solutions are dilute, we take N1 = 1. Equation 2 may be rearranged to the form of eq 1with the equilibrium constant for eq 2, K , equal to b. From thermodynamic^^,^
(y ) T =
AV
(3)
Results and Discussion The results are summarized in Figure l a , which shows the sequence of adsorption isotherms a t 25 O C for up to 3000-bar pressure on the system. Figure l b displays the same data plotted according to the linear form of the Langmuir equation: bC n = .-,n C / n = l/n,b +C/n, (1) 1 + bC' Here n denotes moles adsorbed per gram of adsorbent, n, is the maximum adsorption, and C is the supernatant concentration of adsorbate in moles/liter. The isotherms were Langmuirian within experimental error, except perhaps for the lowest concentration point at 1500 and at 3000 bar. The above results represent true adsorption data in the sense that the apparent concentration changes have been corrected for small pressure shifts in the absorption spectrum of the solute and that the increase in solvent density with pressure has been allowed We thus obtained the actual amount of adsorbate in solution in each case and, by material balance, the amount adsorbed. (4) Demas, J. N.; Adamson, A. W. J. Am. Chem. SOC. 1971,93,1800. ( 5 ) Fetterolf, M. L.; Offen, H. W. J. Phys. Chem. 1985, 89, 3320.
and our data give AP = 7.7 cm3 mol-'. Qualitatively, one might expect a negative A T for adsorption on the ground that an adsorbent-adsorbate complex should be more compact than the free surface plus free adsorbate. Equation 2 reminds us, however, that we are dealing with an exchange of solvent for charged solute. The positive A T could be interpreted in terms of electrostrictive effects.s High pressures favor separated ions, Le., desorption of the cationic adsorbate from the surface. The desorption process would favor greater compaction around the ions in solution, thus decreasing the total volume, and be favored by pressure. Alternatively, solvent AN may be more compact in the interfacial region than in the bulk phase. Various approximations have been made, of course. Activity coefficients and changes in them have been ne(6) Adamson, A. w. The Physical Chemistry of Surfaces, 4th ed.; New York, Wiley: 1982; p 371. (7) Adamson, A. W. A Textbook of Physical Chemistry, 2nd ed.; Academis Press: New York, 1979. (8) Isaacs, N. S. Liquid Phase High Pressure Chemistry; Wiley: Chichester, 1981. (9) Conway, B. E.; Currie, J. C. J. Chem. Soc., Faraday Trans. I 1978, 74, 1390.
Langmuir 1986,2, 543-548 glected. Pressure effects are often difficult to relate to any particular volume change; they result from all density changes, including possible pressure effects on the surface structure of the adsorbent. Finally, eq 2 assumes a 1:l stoichiometrv: it is more likelv that each adsorbed Ru(bpy)32+molecule displaces more than one AN molecule from the interfacial region. The conclusion remains that relatively modest pressures can have a significant effect on equilibrium adsorption
543
from solution. We plan further studies in this area, including systems showing irreversible adsorption. As a general comment, there should be many adsorption systems of practical value where the effect of pressure is imDortant.
Acknowledgment. This investigation was supported in part by the Science FoundationRegistry No. Ru(bpy),2+,15158-62-0; Ti02,13463-67-7.
Vibrational Spectroscopic Examination of Langmuir-Blodgett Mono- and Multilayer Fatty Acids by High-Resolution Electron Energy Loss Spectroscopy Joseph H. Wandass and Joseph A. Gardella, Jr.* Department of Chemistry, University at Buffalo, SUNY, Buffalo, New York 14214 Received August 5, 1985. In Final Form: January 6, 1986 HREELS analysis of Langmuir-Blodgett mono- and multilayer systems was conducted at a primary electron energy of 6.5 eV and a specular angle of 60’. Results from the analysis of multilayers of stearic acid on silver for up to 15 layers (approaching 400 A in thickness) demonstrate that successful HREELS spectra can be obtained with extreme surface sensitivity to the outermost methyl groups. Selective enhancement of relative intensities is observed for certain numbers of stearic acid layers, intimating that resonance scattering may be involved in the sampling process. HREELS analysis of oleic acid on Au, Ag, Ge, and A1 demonstrates a complex relationship between the role of surface oxide and morphology of the substrate in the character of the observed spectra. Radical differences are observed in the relative intensities of loss features as a function of substrate. Clear detection of olefinic C-H stretching features in these samples indicates that the substrate has no effect on the detection limit for this loss feature. The ability to detect and discriminate this olefinic stretch is demonstrated for a 3:l molar mixture of stearic/linolenic acids transferred to a Ag substrate. Various loss features are assigned to C-H vibrations, deformations, and twists in light of IR and Raman data for fatty acids and other recent results of HREELS analyses.
Introduction In a previous paper,l we reported the first high-resolution electron energy loss spectroscopy of Langmuir-Blodgett synthesized monolayers of various fatty acids on polycrystalline substrates. I t was shown to be possible to obtain useful HREELS spectra on such substrates with a vacuum/sample interface layer of methyl groups and to be able to deduce the orientation and relative unsaturation of the monolayer constituents. In this paper, we extend the results to the investigation of multilayer systems (where the total layer thickness approaches 375 A) and the influence of various substrates (both conducting and nonconducting) on the characteristics of the HREELS spectra. Recently, the HREELS technique has been used to examine a wider variety of adsorbed overlayers and substrates than ever before. This trend has been toward experiments which analyze polycrystalline and insulating materials and adsorbates which are “large” molecules (number of carbon atoms >4). For instance, insulating materials such as MgO(001) and a-A1203(0001)and semiconducting GaAs have been studied by the HREELS technique using external charge compensation. Pireaux and Caudano et a1.2 examined MgO(001) with HREELS and recorded signals from Fuchs-Kliewer phonons with high resolution (27 cm-’). In addition, they were able to calculate (using the dielectric theory) various optical constants of the media. In a related paper3 by the same
* Author to whom correspondence should
be addressed.
0743-7463/86/2402-0543$01.50/0
group, a-Al,O~(OOOl)surfaces were used to test the validity of the dielectric theory of HREELS sampling. In both reports, an auxiliary defocused electron gun was used to overcome effects of sample charging during the analysis. Dubois et al.4 used HREELS to investigate the mechanism of growth of thin polycrystalline Ag overlayers on GaAs(100). Rao et al.5 also studied the adsorption of CO on various polycrystalline transition metals and alloys such as Mn, Ag, and Pt. Thus, fruitful HREELS analysis of polycrystalline and insulating materials is possible. In the area of adsorbed organic monolayers much has been done to extend HREELS observations beyond small molecules such as C2H2,NO, and CO. Koel and Somorjai et al.! for example, studied benzene adsorbed on Rh(ll1) crystal faces and determined the orientation of the organic ring through angle-resolved HREELS. Richardson and Hoffman have also used HREELS to study the vibrational structure of pyridine on Pt(l10)7and phenol on Cu(110).8 Demuth and Avouris were able to use .rr 7 ~ *electronic
-
(1) Wandass, J. H.; Gardella, J. A., Jr. Surf. Sci. 1986, 150, L107. (2) Thiry, P. A.; Liehr, M.; Pireaux, J. J.; Caudano, R. Phys. Rev. B 1984,29,4024. (3)Liehr, M.;Thiry, P. A.; Pireaux, J. J.; Caudano, R. Phys. Reu. B 1986,31,42. (4) Dubois, L. H.; Schwartz, G. P.; Camley, R. E.; Mills, D. L. J. Vac. Sci. Technol., A 1984,2,1086. ( 5 ) Vishnu, P. K.; Rao, C. N. R. Indian J. Chem., Sect. A 1984,23A (12), 973. (6) Koel, B. E.; Crowell, J. E.; Mate, C. M.; Somorjai, G. A. J. Phys. Chem. 1984,88,1988. (7)Richardson, N.V. Vacuum 1983,33,787. (8)Richardson, N.V.; Hofmann, P. Vacuum 1983,33,793.
0 1986 American Chemical Society
