168
Langmuir 1990,6, 168-172
homogenous nucleation process followed by a preferential growth at the periphery of the particles. This, in turn, resembles the surface-phase polymerization mechanism of styrene in aqueous emulsion polymerization.
Acknowledgment. We thank NSERC Canada for their support of this research. Registry No. PMMA, 9011-14-7; PS, 9003-53-6.
Effect of Dispersed Manganese Oxides on the Decomposition of Permanganate Solutions Liang-zhong Zhao,+ Jack G. Davis, Jr., and Vaneica Y. Young* Department of Chemistry, University of Florida, Gainesville, Florida 32611 Received July 21, 1988 The effect of manganese oxide particles dispersed on carbon substrates upon the decomposition of dilute, neutral permanganate solutions is investigated. Dispersed Mn,O, and MnO were prepared by vapor deposition, and dispersed MnO, was prepared by solution deposition. In all cases, dispersed oxides were more active than thick or bulk oxides.
Introduction The effect of dispersion of a metal upon its catalytic activity has been intensely investigated in recent years. The dispersed metal is usually prepared as a distribution of small particles supported on inert substrates, such as silica, alumina, or carbon. In most cases, it has been shown that the catalytic activity changes as the average particle size changes.'+ A fundamental goal of studies such as these is to characterize those properties of the particles, e.g., electronic structure, morphology, chargetransfer behavior, etc., which are responsible for the enhanced catalytic activity. Metal oxides also exhibit greater catalytic activity when discontinuously dispersed on a substrate. These materials are commonly prepared by chemical impregnation of powdered supports7 followed by activation, e.g., calcining, to give crystallites of the metal oxide on the surfaces of individual support particles. However, the activation step occurs a t such high temperatures that lateral diffusion on the surfaces of the support particles prevents precise control of the size distribution of the particles. As one example of many which could be cited, the results of Cimino et a1.' will be mentioned. These authors used X-ray photoelectron spectroscopy (XPS), X-ray diffraction, and optical measurements to show that CrOJSiO, catalysts contain molecularly dispersed Cr(V1) and a bimodal distribution of Cr,O, particles with sizes 120 and 2400 A. We are conducting research aimed a t ascertaining the electronic structure of small particles of metal oxides and its correlation with their catalytic activity. The former
'
Present address: Institute of Chemistry, Academia Sinica, Beijing, China. (1) Yates, D. J. C.; Sinfelt, J. H. J . Catal. 1967,8, 348. (2) Boudart, M.; Aldag, A. W.; Ptak, L. D.; Benson, J. E. J . Catal. 1968, 11, 35. (3) Lam, Y. L.; Sinfelt, J. H. J. Catal. 1976, 42, 319. (4) Dominguez, J. M.; Yacaman, M. J. J. Catal. 1980, 64, 223. ( 5 ) Fuentes, S.; Figueras, F. J. Catal. 1980, 61, 443. (6) Graydon, W. F.;Langan, M. D. J. Catal. 1981.69. 180. (7) Sinfelt, J. H. Science 1977, 195, 641. (8) Cimino, A.; DeAngelis, B. A,; Luchetti, A.; Minelli, G. J. Catal. 1976, 45, 316.
0743-7463/90/2406-0168$02.50/0
studies are facilitated by supporting the metal oxide particles on carbon substrates. Experimentally, the electronic structure of these systems is studied by using core and valence level XPS. The catalytic activity of these systems may be studied by choosing an appropriate reaction. In this paper, we report the effect of dispersed MnO, Mn203,and MnO, on the decomposition of dilute, neutral potassium permanganate solutions. Experimentally, this is a very simple reaction to study because the kinetics is slow enough that the course of the reaction can be followed spectroph~tometrically.~~~~ Although only a few will be cited, many researchers have investigated the kinetics of the oxidation of organic compounds by permanganate in acidic, neutral, and alkaline solutions by using spectrophotometry a t a wavelength corresponding to an absorption maximum in the first excitation band of permanganate.l'-l4 Here, neutral solutions have been used because the rate of reaction is thermally increased in strongly acidic15or strongly alkaline15*16solutions. The rate of reaction has also been reported to increase as a result of exposure to visible The decomposition product MnO has been shown to catalyze the decomposition reaction.'6 Mn(II1) in the solid has been proposed as an intermediate in the decomposition r e a ~ t i o n . ' ~ ~Manganous ~' ion has been shown to react with permanganate in 3 M HClO, s o l ~ t i o n . ' ~It is wellknown that it autocatalyzes oxidations by permanga(9) Kachan, A. A.; Sherstoboeva, M. A. Zh. Neorg. Khim. 1958, 3, 1089. (10) Iogansen, A. V.; Grushina, N. M. Khim. Fiz. 1982,1, 121. (11) Wiberg, K. B.; Fox, A. S. J. Am. Chem. SOC.1963,85, 3487. (12) Freeman, F.; Fuselier, C. 0.; Armstead, C. R.; Dalton, C. E.; Davidson, P. A.; Karchesfski, E. M.; Krochman, D. E.; Johnson, M. N.; Jones, N. K. J . Am. Chem. SOC.1981,103, 1154. (13) Toyoshima, K.; Okuyama, T.; Fueno, F.J . Org. Chem. 1980, 45, 1600. (14) Ogino, T. Tetrahedron Lett. 1980,21, 177. (15) Zimmerman, G. J. Chem. Phys. 1955,23, 825. (16) Narita, E.; Hashimoto, T.; Yoshida, S.; Okabe, T. Bull. Chem. SOC.Jpn. 1982,55, 963. (17) Duke, F. R. J. Phys. Chern. 1952, 56, 882. (18) Shafirovich, V. Ya.; Shilov, A. E. Kinet. Katal. 1979,20, 1156. (19) Adamson, A. W. J . Phys. Colloid Chem. 1951,55, 293.
0 1990 American Chemical Societv
Langmuir, Vol. 6, No. I , 1990 169
Decomposition of Permanganate Solutions nate in acidic solution,20but its catalytic behavior in neuOSO tral or basic solution has not been reported. What has been reported is that Mn(I1) is oxidized by potassium permanganate to form manganite precipitates,,l mixed salts of Mn(OH), and MnO, with a layered structure. This reaction is said to proceed autocatalytically.22 Thus, 0.500 it seems possible that both MnO and Mn203 can affect 0, the decomposition of permanganate, although neither can U C be regarded as true catalysts, in the sense that neither is 0 0 present a t the end of the reaction. The results of XPS L 0 studies on the electronic structure of the dispersed manUl ganese oxide particles have been reported e l ~ e w h e r e . ~ ~ , ' ~ 2 0450
Experimental Section The support used in these studies is 1 mm thick, light-tight carbon foil (99.8%, Goodfellow Metals) cut into 1 cm X 1.2 cm rectangles. The manufacturer reports that the 0.2% ash is comprised of 40% SiO,, 13% CaO, 13% Fe,O,, 13% Cr,O,, 10% P,O,, 7% A1,03, 6.5% NaOH, and 2.3% MgO and other oxides. The preparation and characterization of the various samples fall into one of three sets as follows: Samples of dispersed and thick MnO on carbon foil were prepared by vapor deposition from manganese chips (Alpha Products, M 3n7) heated in a masked tungsten filament basket of an evaporator probe system (HP 18612 A) attached to the sample preparation chamber (base pressure approximately lo-' Torr) of an HP 5950A ESCA spectrometer onto 1.0 cm X 1.2 cm fresh carbon foil rectangles whose surface oxygen has been removed by prior treatment.23s25Reducing residual gases26coupled with a thick layer of oxide present on the manganese chips lead to the deposition of MnO.n,m The geometrical relationship between the sample and the probe is such that deposition occurs in a circular area with a diameter of 0.9 cm on one side of the foil. The composition of these deposits is unchanged even after 4 months of storage in a drybox filled with ultrapure nitrogen. This observation is consistent with reports that MnO is quite stable to air exposure at room temperature for short periods of time.29.30Dispersed and thick deposits of Mn,O, on surfacedeoxygenated fresh carbon foil have been prepared by air oxidation of vapor-deposited MnO at 250 "C in a stable therm gravity convection oven for 2 h.31 Conversion has been confirmed by comparing the Mn 2p core level XPS spectra (recorded by using the HP 5950A spectrometer) with those of pellets of Mn,O, powder pressed at 35000 psi with a Beckman die and a Carver laboratory press. By use of variable-angle XPS, as described previou~ly,~~ the monolayer coverages, 9, of these samples have been estimated as 0.23 for dispersed MnO, 0.18 for dispersed Mn,O,, and 1.0 for thick MnO and thick Mn,O,. Samples of MnO, on surface-deoxygenated fresh carbon foil have been prepared in triplicate by exposing carbon foil rectF potassium permanganate angles to dilute, neutral 2.0 x for periods of 2, 4, 20, and 100 h, respectively. This method has previously been shown to result in MnO, decomposition on other supports.32 The rectangles were harvested, washed first with deionized water and then with reagent grade acetone, and (20) Adler, J.; Noyes, R. M. J. Am. Chem. SOC.1955, 77, 2036. (21) Buser, W.; Graf, P.; Feitknect, W. Helu. Chim. Acta 1954, 37,
noon LJLL.
(22) Polissar, M. J. J. Phys. Chem. 1935, 39, 1057. (23) Young, V.; Zhao, L. Chem. Phys. Lett. 1983,102,455. (24) Zhao, L.; Davis, J.; Young, V. J. Electron Specttosc. Relat. Phenom. (submitted). (25) Young, V. Carbon 1982,20, 35. (26) Handbook of X-Ray and Ultraviolet Photoelectron Spectroscopy; Biggs, D., Ed.; Heyden: London, 1978. (27) McCarroll, W. H. J.Appl. Phys. 1968, 39, 3414. (28) Wertheim, G.; Heifer, S.; Guggenheim, H. Phys. Reu. B 1973, 7 5.56 ., ---. (29) Oku, M.; Hirokawa, K. J. Electron Spectrosc. Relat. Phenom. 1975, 7, 465. (30) Foord, J.; Jackman, R.; Allen, G. Philos. Mag. A 1984,49, 657. (31) Helsop, R.; Robinson, P. Inorganic Chemistry; Elsevier: New York, 1964; p 681. (32) Shabalina, 0.;Ryzhen'kov, A.; Egorov, Yu.; Stotskii, V.; Popov, V.; Kotel'nikov, A. Deposited Document SPSTL 422 khp-D8 1980 CA 9 6 208570 a.
1
1
i 12
24 36 time (hours)
48
Figure 1. Absorbance versus time data for MnO and Mn,O, permanganate solution; (0)permanganate solusamples: (0) tion + carbon foil; (+) permanganate solution + dispersed MnO on carbon foil; (X) permanganate solution + dispersed Mn,O, on carbon foil; (A)permanganate solution + thick Mn203 on carbon foil.
finally dried at room temperature. Visual evidence of MnO, deposition has been observed for the 20-h and the 100-h samples as golden brown patches on both sides of the foil. The patches were denser on the 100-h samples. XPS analyses on these samples could not be performed until a much later date. The analyses have been performed on a Kratos XSAMSOO instrument using A1 Ka radiation. The presence of MnO, has been confirmed by comparison with spectra for pressed pellets of MnO, powder (Mallinckrodt, 99.5%). Because of the time lapse between preparation and XPS analyses, the monolayer coverages could not be obtained intervening surface contamination invalidates the values. Samples of MnO, on aged carbon foil (4 years old) have been prepared in duplicate by exposing carbon foil rectangles to 3.3 X F potassium permanganate for periods of 2,20, and 100 h, respectively. Attempts were made to deoxygenate the surfaces of these foil samples, but not all of the oxygen could be removed. XPS analyses on a Kratos XSAM 800 instrument revealed that the persistent oxygen is due to SiO,, which apparently has segregated to the surface from the bulk. The 20-h and 100-h samples do not show the same pattern of MnO, deposition as that observed for the surface deoxygenated fresh carbon foil. By use of variable-angle XPS, the monolayer coverages of these samples have been estimated as 0.055,0.133, and 0.073 for the 2-, 20-, and 100-h samples, respectively. Surface-deoxygenatedcarbon foil rectangles,the samples from sets 1and 2 (deposition side up), and pressed pellets of Mn,O, powder have been placed in dry 50-mL Erlenmeyer flasks and covered with 50 mL of approximately 2.0 X lo4 F neutral potassium permanganate. The unstirred solutions have been sampled with replacement at 12-h intervals, and the absorbances at 525 nm versus deionizedwater have been measwed on a Bausch & Lomb Spectronic 20 spectrophotometer. Thick MnO films became detached from the carbon foil, so these results have not been considered further, since the area of MnO exposed does not remain constant during the analysis and is indeterminant. The MnO, pellets repeatedly broke up in the solution, although the pieces did not dissolve. Only one of the Mn,O, pellets remained intact for 60 h. Thus, for consistency, we have not included the data obtained for the pressed pellets. Because product MnO, can deposit on these samples, the volume of the samples can change with time. However, these volume changes should be insignificant, because most of the volume is due to the carbon foil. More significant are the changes in the surface area, which we cannot measure. Typical results for these samples are shown in Figures 1 and 2. Aged carbon foil rectangles
170 Langmuir, Vol. 6,No. 1, 1990
Zhao et al. P e r e z - B e n i t ~regard ~ ~ this absorbance t o be due to light scattering by colloidal manganese dioxide. Thus, in the absence of intermediates which absorb at 525 or 526 nm, we can write the following equation for the absorbance:
A = ~ M ~-ob[MnOil , t~~o,b[Mn02] (1) where t is molar absorptivity and b is the cell path length. If the product colloidal manganese dioxide is not lost from solution, we have [MnO,] = [MnO;], - [MnO;] (2) where [MnO,-], is the initial concentration of permanganate. The first-order decay equation34 24
12
36
48
60
time (hours)
Figure 2. Absorbance versus time data for colloidal MnO, deposF potassium permanganate (see text): (0) ited from 2.0 x permanganate solution; ( 0 )permanganate + 100-h deposit; (+) permanganate + 20-h deposit; (A)permanganate + carbon foil; (X) permanganate + 4-h deposit; (A)permanganate + 2-h deposit.
(3)
d[MnO;]/dt = -k[MnO;l can be written in terms of A as dA/dt = -kA
+ ~&~~bk[MnO;]~
(4)
If the product colloidal manganese dioxide is completely lost from solution, say by deposition on the solid substrate, then
0 800
dA/dt = -kA (5) The second-order autocatalytic expression for the rate of disappearance of permanganate is given by34
0600 -
(6) d[MnO,-] /dt = -k[MnOi] [MnO,] When we express everything in terms of A, we obtain
(~M,,o,- tM,,o,)b(dA/dt) = kA2 - (E&o, +
aJ U
c
2
t~,o,-)bk[MnO;]d
0400-
b
-k
[MnOLl; (7)
This is a separable, first-order nonlinear differential equation, the solution of which is given by
IJl
n
Q
In {(t~,o,-b[MnO;]o - A)/(A - ~ ~ ~ o , b [ M n O i I=o ) l k[MnOilot + In {(fMno,-b[MnO,-10 - A,)/ 0
2
0 12
0
24
36
time
(hours)
5 48
Figure 3. Absorbance versus time data for colloidal MnO, deposF potassium permanganate (see text): (0) ited from 3.3 X permanganate solution; (A)permanganate + carbon foil; ( 0 ) permanganate + 20-h deposit; (+) permanganate + 2-h deposit; (x) permanganate + 100-h deposit.
treated to remove nonpersistent oxygen and samples from set 3 have been treated similarly, except that 3.3 X lo4 F neutral potassium permanganate has been used and the absorbances at 526 nm versus deionized water have been measured on an HP8450A diode array, reverse optics spectrophotometer. This spectrophotometer only allows even wavelengths to be read in the range 400-800 nm. Typical results are shown in Figure 3.
Results Apparent rate constants can be obtained from the absorbance data. Since various researchers have considered the rate of disappearance of permanganate to be first order with respect to permanganate, while others have considered it to be second-order autocatalytic, these two models must be considered. In addition, we also consider a zero-order model. Iogansen and Grushinal’ have published the absorption spectrum for colloidal manganese dioxide, one product of the decomposition; they show that its molar absorptivity is only slightly less than that of permanganate at 525 and 526 nm. Mata-Perez and
(Ao - t & ~ , b [ M n O i I ~ )(8) l where A, is the absorbance at t = 0. If product colloidal manganese dioxide is completely lost from the solution, then the equation has the form In {(tMno,-b[MnO,-]o - A)/A1 = k[MnO;Iot
+
In { ( ~ ~ ~ o , - b [ M n O - iAo)/Ao) Io (9) where eq 8 and 9 are valid for t # 0. Note that k’ = k[Mn04-], has the same units as k for the first-order decay. For zero-order decay, we have
A = -kt + C (10) where C is the constant of integration. In Table I, we show the k or k’ obtained for the fit (eq 4, 5,8, 9, or 10) which gives the largest correlation coefficient. The k values for the carbon foil samples are very consistent and give an average value of (11.1 f 1.8) X h-’. With the exception of the MnO, with CP = 0.133, all of the apparent k or k’ values are more than three standard deviations away from the average value. Thus the presence of manganese oxide on these samples enhances the permanganate decomposition relative to carbon foil. In order to compare the activities of the oxides, we assume that the oxides and carbon foil act independently of each other. Then the apparent k or k’ will be (33) Mata-Perez, F.; Perez-Benito, J. Can. J . Chem. 1985,63,988. (34) Frost, A,; Pearson, R. Kinetics and Mechanism, 2nd ed.; Wiley: New York, 1961; p 12-20,
Langmuir, Vol. 6, No. 1, 1990 171
Decomposition of Permanganate Solutions Table I. Apparent Rate Constants for Permanganate Decay
0
0
correlation sample fresh C 1 fresh C 2 aged C dispersed MnO dispersed Mn,03 thick Mn,O, MnO,; 2 h MnO,; 4 h MnO,; 20 + 100 h MnO,; 3 = 0.055 MnO,; 3 = 0.073 MnO,; 3 = 0.133
k or k’, h-’ coefficient
best model 1st order 1st order 1st order/MnO, 1 s t order 1st order 1st order autocatalytic autocatalytic autocatalytic 1st order/MnO, 1st order/MnO, 1st order/MnO,
loss
loss loss loss
9.2 X 12.7 X 11.3 X 17.2 X 22.9 X 23.0 X 51.3 X lo-’ 41.8 X 33.2 X 17.5 X 18.6 X lo-’ 13.9 X
0.980 0.992 0.988 0.982 0.996 0.995 0.997 0.994 0.984 0.998 0.992 0.995
Table 11. Weighing Factors and Calculated kolidefor Samples with Known sample dispersed MnO dispersed Mn,03 thick Mn,03 MnO,; 3 = 0.055 MnO,; = 0.073 MnO,; 3 = 0.133
Woiide
Wcarbon
0.061 0.048 0.265 0.132 0.175 0.319
0.939 0.952 0.735 0.868 0.825 0.681
koxider
h-’
110x 10-3
260 x 56 x 60 x 54 x 20 x
10-3 10-3 10-3 10-3 10-3
a weighed average of k or k’ for the oxide and k or k’ for carbon foil. The weighing factors are given by the fractional areas of the two phases. Thus = (Aoride/Atotal)koxide + (Acarbon/Atotal)kcarbon (11) Here, the total area, Atotal, is equal to 2.4 cm’, since both sides of the sample must be considered. The area for the vapor deposited oxides is given by
Aoxide= @IId2/4 (12) where @ is the monolayer coverage and d is the diameter of the circular deposition area. For the MnO, samples, the area for the oxide is given by Aoxide = 2.4@ (13) Weighing factors and calculated values for kOxideare shown in Table 11.
Discussion The goal of this study was to assess whether or not dispersion has an effect on the chemical activity of carbon-supported manganese oxides. The results for Mn,O, and MnO, show clearly that the more dispersed oxides are more active. However, these results also pose a number of interesting questions, the answers to which can be explored by further experimentation. For example, the results in Table I suggest that the MnO, on surface-deoxygenated fresh carbon foil is more active than MnO, on aged carbon foil. Clearly, the surfaces need to be further characterized by using other surface-sensitive techniques, such as ISS to confirm the zero-angle XPS results and SIMS to look for hydrocarbons, which cannot be detected a t low coverages on a carbon substrate. In our experiments, only the initial state of the manganese oxides is known. However, Mn(I1) is oxidized by atmospheric oxygen to Mn(1V) at pH 18.35 Since our solutions are not deoxygenated, dissolved oxygen can oxidize the MnO, although a t somewhat lower rates since pH = 7 . Also, permanganate can oxidize Mn(I1) to M I I ( I I I ) . ~Mn,03 ~ is stable to oxidation by air a t room
11 .’u “ O m
Mn02 + O 2
+
OH
k&/
Mn0;
\I
Mn02
O2
+
e-
Figure 4. Channels of decomposition for neutral permanganate.
temperature3’ and thus is expected to be stable to dissolved oxygen. Although it can be oxidized to MnO, by permanganate, its postulated role as an intermediate17’18means that it probably stays around in some form until the end of the reaction. No in situ method capable of providing information on the composition of these interfaces has been published. A very time intensive study can be done by harvesting the substrates after various exposure times to permanganate and characterizing the substance by using XPS, ISS, and SIMS. The results in Table I for MnO, on surface-deoxygenated fresh carbon foil and for MnO, on aged carbon foil also suggest that the details of the mechanism of permanganate decomposition are different. Unraveling these details would be a herculean task, because several channels are available for the decomposition of permanganate in neutral medium, and these channels are not necessarily independent. These channels are shown in Figure 4. In the Zimmerman channel,15 the permanganate ion absorbs a photon to form the excited state (Mn0,-)*, which rapidly relaxes nonradiatively to a vibrationally excited ground-state ion, (MnO,-)§. This ion can then thermally dissociate in a single step into MnO, and 0, or relax nonradiatively to the ground vibrational state of MnO,-. The location of the electron in the former process has not been determined; however, if it resides on 0,, this ion will undergo rapid disproportionation as follows:37
20,-
+ 2H2O
HZO,
+ 02 + 20H-
(14) If the electron resides on MnO, to give the hypothetical Mn(II1) species Mn0,-, this species can be oxidized to MnO, by permanganate. Manganese dioxide could accelerate decomposition of permanganate in this channel by increasing k, relative to k,. Note that the fate of the electron as proposed here causes formation of OH-. Channel 2 is the oxidation of water, which is thermodynamically favored, but which occurs a t a very slow rate in neutral medium because of the “overvoltage effect”.38 According to Douglas et al., this effect is not limited to electrolytic cells. Apparently it refers to the fact that gas evolution has a very high activation energy and thus will be the rate-determining step in the mechanism. Shafirovich and Shilov18have studied quantitatively the effect of hydrogen ion on the reaction, finding acceleration and strong dependence on the hydrogen ion concentration for pH 1 2 . They propose that the catalyzed decomposition is
Mn(1V) + HMnO,
-
--*
Mn(1V)-HMnO,
-
Hz0
Mn(1V) + Mn(II1) + O2 (15)
(36) Davies, G . Coord. Chem. Reu. 1969,4, 199. (37) Kim, S.; DiCosimo, R.; SanFillippo, J., Jr. Anal. Chem. 1979,
51,679. (35) Coughlin, R. W.; Matsui, I. J. Catal. 1976, 41, 108.
(38) Douglas, B.; McDaniel, D.; Alexander, J. Concepts and Models of Inorganic Chemistry; Wiley: New York, 1983.
Langmuir 1990, 6, 172-177
172
The reaction is also accelerated in strong base.15,16 A balanced equation for the oxidation of water by permanganate in terms of half reactions shows that it is really OH- which is being oxidized. It has been proposed that manganese dioxide catalyzes the oxidation of water by means of certain active sites which facilitate the reversible coordination of water molecules in the vicinity of Mn(VII).18 In channel 3, there is some reductant present which can be oxidized by permanganate. The reductant could be Mn(II), Mn(IV), or even contaminating hydrocarbons on the carbon substrates. Because OH- links all of these channels, it would be difficult to determine which one is the most important. Indeed, the experimental conditions probably determine the primary channel. I t seems likely that the reason why no single mechanism has been proposed for the autodecomposition of permanganate after almost 100 years of experimentation is exactly this complexity. The scenario is further complicated by the fact
that permanganate decomposition to M n 0 2 is only one reaction, and a set of reactions which involve particle growth, aggregation, dissociation, and precipitation3' must really be considered. It is clear that much more remains to be done.
Acknowledgment. This work has been performed in part at Texas A&M University, where it was supported by the Robert A. Welch Foundation (Grant No. A-771). Work done at the University of Florida has been supported in part by a grant from the National Science Foundation (Grant No. RII- 8508120). Registry No. MnO,-, 14333-13-2; MnO, 1344-43-0; MnO,, 1313-13-9; Mn,O,, 1317-34-6; C, 7440-44-0. (39) Heicklen, J. Colloid Formation and Growth: A Chemical Kinetics Approach; Academic Press: New York, 1976.
Structural Studies on Langmuir-Blodgett Films Containing Nitrostilbene and Hemicyanine Dyes H. Ancelin,t G. Briody,t7s J. Yarwood,*yt J. P. Lloyd,t M. C. Petty,$ M. M. Ahmad,$ and W. J. Feastt Department of Chemistry, University of Durham, Durham City, DH1 3LE, U.K., and Molecular Electronics Research Group, School of Engineering and Applied Science, University of Durham, Durham City, DH1 3LE, U.K. Received February 7, 1989. I n Final Form: June 28, 1989 Infrared spectroscopic studies on amidonitrostilbene and hemicyanine dye Langmuir-Blodgett layers (in multilayer and bilayer configurations) were performed to provide information about the molecular interactions responsible for enhanced nonlinear optical properties. In particular, hydrogen bonding at the carbonyl groups (and other head group interactions) provides a means of aligning and partially separating the hemicyanine molecules in the bilayer configuration. Attempts have been made to assess the degree and nature of microscopic ordering as the number of monolayers increases. These data are assessed in the light of RHEED and ellipsometric measurements, and possible sources of discrepancy are discussed.
I. Introduction There is much current interest in the exploitation of the nonlinear optical properties of organic materials',' for second-harmonic generation (SHG),3for electrooptical switching: and for modulators in optical signal processing. A wide variety of molecules with the necessary (second order nonlinear) properties have been produced,l*' and their noncentrosymmetric arrangement on a suitable substrate has been achieved in a number of ~ a y s . ' * ~One * ~ of the most elegant of these is the engi-
* Author t o whom correspondence should be addressed.
'Department of Chemistry.
* Molecular Electronics Research Group. 5 Present address: B.P. Research, Sunbury on Thames, Middx. TW16 7LN. (1) Ahmad, M. M.; Feast, W. J.; Neal, D. B.; Petty, M. C.; Roberts, G. G. J . Mol. Electron. 1986, 2, 123. (2) Girling, I. R.; Kolinsky, P. V.; Cade, N. A,; Earls, J. D.; Peterson, I. R. Opt. Commun. 1985,55, 289. (3) Neal, D. B., Petty, M. C., Roberta, G. G.; Ahmad, M. M.; Feast, W. J.; Girlina, I. R.; Cade, N. A.; Kolinsky, P. V.; Peterson, I. R. Elec-
tron Lett. 1@6, 22, 460. (4) Smith, P. W. Bell Syst. Tech. J. 1982,61, 1975. (5) Zyss, J.; Tsoucaris, G. Mol. Cryst. Liq. Cryst. 1986, 137, 303.
0743-7463/90/2406-Ol72$02.50/0
neering of molecular multilayers with the necessary macroscopic polarity using the Langmuir-Blodgett (LB) deposition technique.' Such multilayers have been characterized by using a variety of methods including SHG,2*3,8 X-ray and electron d i f f r a c t i ~ n , ~surface ~'~ plasmon and optical ~ p e c t r o s c o p y . l ~There - ~ ~ has been relatively little work published on the microscopic chemical interactions which presumably affect, and possibly control, the electronic properties of these materials. Infrared spectroscopy provides a powerful method for structural characterization of organic thin films as (6) Tweig, R. J.; Dick, C. W. J. Chem. Phys. 1986,85, 3537. (7) Blodgett, K. B. J . Am. Chem. SOC.1935,57, 1007. ( 8 ) Stroeve, P.; Srinivasan, M. P.; Higgins, B. G.; Kowel, S.T. Thin Solid Films 1987, 146, 209. (9) Earls, J. D.; Peterson, I. R.; Russell, G. J.; Girling, I. R.; Cade, N. A. J. Mol. Electron. 1986,2, 85. (10) Neal, D. B.; Russell, G. J.; Petty, M. C.; Roberts, G. G.; Ahmad, M . M.; Feast, W. J. J.Mol. Electron. 1986, 2, 135.
(11) Electromagnetic Surface Excitation; Wallis, R. F., Stegerman, G. I., Eds.; Springer: Heidelberg, 1986. (12) Prasad, P. N. Thin Solid F i l m 1987, 152, 275. (13) Rao, D. N.; Burzynski, R.; Mi, X.; Prasad, P. N. Appl. Phys. Lett. 1986, 48, 387. (14) Prasad, P. N.; Swiathiewicz, J.; Eisenhard, G. Appl. Spectrosc. Reu. 1982, 18, 59.
0 1990 American Chemical Society
