4158
J . Phys. Chem. 1989, 93, 4158-4167
Diffuse Reflectance Infrared Fourier Transform Spectroscopic Investigation of the Decomposition of Carbon-Supported Iron Carbonyl Clusters Jeremy J. Venter and M. Albert Vannice* Department of Chemical Engineering, The Pennsylvania State University, University Park, Pennsylvania 16802 (Received: July 26, 1988: In Final Form: February 27, 19891
The thermal decomposition of Fe3(C0)12has been studied for the first time by dispersing this cluster on an oxygen-free carbon surface and monitoring its behavior by diffuse reflectance infrared Fourier transform spectroscopy (DRIFTS). The Fe3(CO) clusters on carbon decomposed to Fe(CO), in either He or H,; consequently, the decarbonylation of both clusters could be followed. First-order rate constants of decomposition in He or H, were determined for each cluster and compared to literature k , values for nucleophilic substitution reactions in solution. The kl values for each reaction were very similar for Fe3(C0),*, but the k l value for decomposition was much higher for Fe(CO)5. The rate-determining step in either the substitution or the decomposition reaction appears to be the removal of the first C O ligand. No stable hydrido-iron clusters were formed. The activation energy of decomposition of Fe3(C0)12was near 18 kcal/mol in He and 21 kcal/mol in H2 while that for Fe(CO), was near 15.5 kcal/mol in either gas. These activation energies are lower than those observed for substitution reactions in solution, but this can be explained by Fe-Fe bond formation during the decomposition process to create very small Fe crystallites. Large amounts of chemisorbed CO confirmed the formation of well-dispersed Fe on carbon following decomposition at 673 K under either He or Hz; however, DRIFTS spectra of CO at 195 K detected no IR-active C O species and at 300 K only Fe(CO), was present. The heat of adsorption of CO on these small Fe particles was measured calorimetrically at 300 K and found to be 15.0 & 1.6 kcal/mol. This study represents a portion of the first successful application of an IR spectroscopic technique to characterize carbon-supported metal catalysts.
Introduction Carbon-supported iron catalysts are currently attracting significant attention.'-20 Although iron is one of the most frequently studied C O hydrogenation catalysts due to its high activity,'-'2'21-28
( 1 ) Chen, A.; Kami::;ky, M.; Geoffroy, G. L.; Vannice, M. A . J . Phys. Chem. 1986, 90, 4810. (2) Chen, A . A.; Vannice, M . A.; Phillips, J. J . Phys. Chem. 1987, 91. 6257. (3) Jones, V. K.; Neubauer, L. R.; Bartholomew, C. H. J . Phys. Chem. 1986, 90, 4832. (4) Vannice, M. A,; Walker, P. L.;Jung, H. J.; Moreno-Castilla, C.; Mahajan, 0. P. Proc. 7th Int. Congr. Caral. 1980, paper A31. (5) Jung, H. J.; Walker, P. L.; Vannice, M. A . J . Card. 1982, 75, 416. (6) Jung, H. J.; Mulay, L. N.; Vannice, M. A,; Stanfield, R. M.; Delgass. W. N. J . Catal. 1982, 76, 208. (7) Kaminsky, M.; Yoon. K. J.; Geoffroy, G. L.; Vannice, M. A. J . C a r d 1985, 91, 338. (8) Reinoso, F. R.; Gonzalez, J. D. L.; Castilla, C. M.; Ruiz, A. G.: Ramos, I . R. Fuel 1984, 63, 1089. (9) Reinoso, F. R.; Ramos, I . R.; Ruiz, A . G.: Gonzalez, J. D. L. Appl. Catal. 1986, 21. 251. ( I O ) Ruiz, A. G.; Gonzalez, J. D. L.; Ramos, 1. R.: Reinoso, F. R. React. Kine?. Catal. Lert. 1986. 31, 349. ( 1 1 ) Venter, J. J.; Kaminsky, M.; Geoffroy, G. L.; Vannice, M. A . J . Catal. 1987, 103, 450. (12) Sommen. A . P. B.: Stoop, F.; van der Wiele, K . Appl. Curd. 1985, 14. 277. (13) Venter. J . J.; Kaminsky, M.; Geoffroy, G . L.; Vannice, M. A. J . Catal. 1987, 105, 155. (14) Gatte, R. R.; Phillips, J. P. J . Catal. 1987, 104, 365. (15) Niemantsverdriet, J. W.; van der Kraan, A . M.; Delgass, W. N.; Vannice, M. A. J . Phys. Chem. 1985. 89, 67. (16) Phillips, J.; Clausen. B.; Dumesic. J. A . J . Phys. Chem. 1980, 84, 1814. ( 1 7 ) Phillips, J . ; Dumesic, J . A . Appl. Surf. Sci. 1981, 7, 215. (18) Lin, S. C.; Phillips, J. J . Appl. Phys. 1985, 58, 1943. (19) Bartholomew. C . H.; Boudart, M. J . Catal. 1973. 29, 278. (20) Tau, L. M.; Bennett, C . 0. J . Phys. Chem. 1986, 90, 4825. (21) Amelse. J A,; Butt. J. B.: Schwartz. L. H. J . Phys Chem. 1978. 82, 5.58. (22) Barrault, J.; Renard. C. Appl. Carul. 1985. 14. 133.
the preparation of highly dispersed metallic iron is still a challenging task. Interest in the preparation of such catalysts stems partly from reports of increased olefin selectivities over very small Fe ~ r y s t a l l i t e s l ~ " ~and ~ ~ "partly ~ ~ from the fact that well-dispersed Fe on typical oxide supports is frequently oxidized and difficult to Easily reduced, carbon-supported Fe catalysts can now be prepared with metal carbonyl clusters (MCC's) as precursor^.',^,^ This is a result of the surface chemistry of the carbon supports, since high-surface-area amorphous carbons can be prepared and thoroughly dehydroxylated and decarboxylated, thereby removing surface functional groups typically present on oxide support^.^^^ Although electronic interactions between iron and carbon have been p r o p o ~ e d , the ~ ~ exact ~ ~ - ~nature ~ of this surface chemistry remains unclear., The interaction of MCC's with different supports can be used to elucidate the chemistry associated with the surface of these
Dry, M. E.; Ferreira, L. C. J . Catal. 1967, 7 , 352. Dwyer, D. J.; Somorjai, G. A. J . Catal. 1978, 52, 291. Vannice, M. A. J . Card. 1975, 37, 449. Emmett, P. H.; Gray, J. B. J . Am. Chem. SOC. 1944, 66, 1338. Huff, G. A.; Satterfield, C. N. J . Catal. 1984, 85, 370. Krebs, H. J.; Bonzel, H. P. Surf. Sci. 1980, 99, 570. Hugues, F.; Basset, J. M.; Ben Taarit, Y . ;Choplin, A,; Primet, M.; Rajas, D.; Smith, A. K. J . A m . Chem. Soc. 1982, 104, 7020. (30) Hugues, F.; Dalmon, J. A.; Bussiere, P.; Smith, A. K.; Basset, J. M . J . Phys. Chem. 1982, 86, 5136. (31) Hugues, F.; Besson, B.; Bussiere, P.; Dalmon, J. A,; Bassct, J . M . Nouti. J . Chim. 1981, 5, 207. (32) Commereuc, D.; Chauvin, Y . ;Hugues, F.; Basset, J. M.; Olivier, D. J . Chem. Soc., Chem. Commun. 1980, 154. (33) Hugues, F.; Besson, B.; Bussiere, P.; Dalmon. J. A,; Leconte, M.; Basset. J. M.; Chauvin, Y . ;Commereuc, D.ACS Symp. Ser. 1982, 192, 255. (34) Hugues. F.;Besson, B.; Basset, J. M. J . Chem. Sor., Chem. Commun. (23) (24) (25) (26) (27) (28) (29)
1980, 719. ( 3 5 ) Phillips, J.; Dumesic, J. A . Appl. Carol. 1984, 9, 1. (36) Brenner, A.; Hucul, D. A. J . Am. Chem. SOC. 1980, 102, 2484. (37) Brenner, A,; Hucul, D. A. Inorg. Chem. 1979, 18, 2836. (38) Weatherbee, G. D.; Rankin. J. L.; Bartholomew, C. H. Appl. Catal. 1984. / I . 7 3 .
. (41 ) Rodriguez-Ramos, 1.; Rodriguez-Reinoso, F.; Guerrero-Ruiz, A.: Lopez-Gonzalez. J . J . Chem. Technol. Biotechnol. 1986, 36. 67.
e I989 American Chemical Society
Decomposition of Carbon-Supported Fe3(C0)12Clusters materials; for example, the behavior of Os and Ru MCC’s on carbon has been compared to that on various oxide supports in recent s t ~ d i e s . ~ ~Infrared - ~ ~ (IR) spectroscopy has been the predominant technique in most of these investigations, but few IR studies have been reported for either supported Fe MCC’s4’ or Fe catalysts in general,“* and no IR studies of carbon-supported Fe exist due to the opaqueness of the carbon. One reason for the limited number of studies is the ease of oxidation of Fe clusters on oxide surfaces; however, carbon surfaces free of oxygen-containing groups can be prepared to eliminate this problem, and the recent developments in D R I F T S (diffuse reflectance FTIR spectroscopy) along with improved controlled-environment cells have made it possible to characterize MCC’s on carbon supports for the first time.45,46Consequently, this provided an opportunity not only to examine the chemistry and kinetics of the thermal decomposition of Fe carbonyl clusters, but also to study the interaction between C O adsorbed on these carbon-supported Fe particles following complete decarbonylation. Surprisingly, there are no published studies of the decomposition of any iron carbonyl cluster although IR spectra of many Fe clusters in solution exist and kinetic data for nucleophilic substitution reactions involving Fe3(C0)i2and Fe(CO), have been reported. These studies allow a comparison with our kinetic results and provide a bridge between the chemistry of these clusters in solution and their behavior on a clean, solvent-free surface. W e report here rate constants and activation energies for decarbonylation of Fe3(C0)12and Fe(CO)5 in either H e or H,,as quantitatively measured by DRIFTS, as well as DRIFTS spectra after exposure to C O of the reduced Fe particles remaining after decomposition. In addition, the cluster-derived Fe/C catalysts were also characterized by C O and H 2 chemisorption, and the heat of adsorption of C O on these small, C-supported Fe particles was determined and compared to literature values. The catalytic behavior of these catalysts for CO hydrogenation is reported elsewhere.49
(ead)
Experimental Section
The amorphous carbon black (1400 m2/g) used as a catalyst support in this study was CSX-203 from Cabot Corp. (now available as Black Pearls ZOOO), and the Fe3(C0),2 (Strem Chemicals) was used as received. Sulfur and oxygen were removed and from the carbon by treatment in H, a t 1223 K for 12 h,2s5%6 this high-temperature-treated carbon was then stored in a glovebox to avoid air exposure. Before the support was impregnated with the carbonyl cluster, the carbon was heated to 673 K under dynamic vacuum ( kPa) for 8 h. Fe3(C0),2 was dispersed on the carbon by an incipent wetness impregnation technique using dry, degassed T H F as solvent, cooling the carbon in a NaCl/ice water bath, and impregnating under nitrogen using standard Schlenk technique^.^^ Using a previously employed method, the required Fe3(C0),, was dissolved in 120 cm3 of T H F , and three to five successive incipient wetness impregnations of 5 g of carbon, each followed by a 10-30-min evacuation, were used to disperse the clusters on the support.’ After the final impregnation, the solvent was removed by evacuating to kPa for 8 h a t 300 K and the catalyst was stored in a glovebox. The infrared spectra were collected on a N2-purged Mattson Instruments Sirius 100 FTIR using an extensively modified version of a Harrick Scientific HVC-DRP DRIFTS cell which was coupled to a praying mantis mirror assembly (Harrick-DRA-ZCS), (42) Venter, J . J.; Vannice. M . A. Carbon 1988, 26, 889. Venter, J . J. P1I.D. Thesis, Penn State University, 1988. (43) Venter, J. J.; Vannice, M . A . J . A m . Chem. Sor., in press. (44) Venter, J . J.; Vannice, M . A . J . Inorg. Chem., in press. (45) Venter, J . J.; Vannice, M. A. J . Am. Chem. Soc. 1987, 109. 6204. (46) Venter, J . J.; Vannice. M . A . Appl. Spectrosc. 1988. 42, 1096. (47) Gates, B. C. Metal Clusrers; Wiley: New York, 1986. (48) Sheppard. N.; Nguyen, T. T. Adc. Infrared Raman Spectrosc. 1978, 5 , 67. Heal, M. J.; Leisegang, E. C.: Torrington, R . G . J . Catal. 1978, 51, 314. (49) Venter, J. J.: Chen, A. A,; Phillips, J . ; Vannice, M. A. J . Cural. Submitted for publication. (50) Schriver, D.F. The Manipu1ation of Air Sensitice Compound,s: New York. 1969.
The Journal of Physical Chemistry, Vol. 93, No. I O , 1989 4159 as discussed e1sewhe1-e.~~ The DRIFTS samples were prepared by mixing the Fe/C catalyst with precalcined, H,-treated CaF, inside the glovebox using a CaF,:C dilution ratio of 200:l and then loading a portion of this sample into the DRIFTS cell inside the glovebox prior to connecting the cell to the FTlR without air exposure. Rates of decarbonylation were measured in H2 and also in H e by monitoring the D R I F T S peak intensities, after which the sample was heated to 673 K in the same gas for 10-16 h and then cooled in order to allow the collection of spectra during exposure of the fully decarbonylated catalyst to C O a t 195 and 300 K. Following this high-temperature reduction ( H T R ) , the Fe/C catalyst was also investigated under reaction conditions by collecting spectra a t 473, 523, and 573 K in flowing C O ( 1 1 Torr) m d H 2 (749 Torr). The catalyst was subsequently cooled to 300 K and spectra were recorded in flowing H, and CO a t that temperature. The fully decarbonylated catalyst was used as the background material in all cases. The FTIR parameters were set a t a resolution of 4 an-’, and decarbonylation spectra were obtained by averaging between 100 and 1000 scans, whereas the C O adsorption and reaction spectra were obtained by signal averaging 10000 scans. The background spectra in all cases were obtained by averaging I O 000 scans. Data manipulation consisted of ratioing the sample and background spectra to obtain the transmittance spectra, calculating the absorbance spectra from these spectra, base line correcting the absorbance files, and finally calculating the diffuse reflectance files (in Kubelka-Munk units) from the base-line-corrected absorbance files. All spectra presented in this paper are in Kubelka-Munk (KM) units, K M = ( 1 - R0),/(2RO) = k c / s , where Ro = R(sample)/R(background) is the transmittance spectrum, k is the molar extinction coefficient, c is the concentration, and s is a scattering c o e f f i ~ i e n t . ~In~ each case, a concomitant absorbance spectrum was routinely calculated to verify base line correction, and it was essentially identical with the K-M spectrum. The k and s values were constant over predetermined regions .of cluster loading.42 C O and H2 chemisorption on the fresh catalyst was measured in an U H V stainless steel after a low-temperature reduction (LTR) in flowing H2 for 2 h a t 473 K as well as a high-temperature reduction ( H T R ) in flowing H, for 16 h at 673 K. Used samples after the kinetic runs were given only an H T R pretreatment and characterized as described e l s e ~ h e r e .Irre~~~ versible C O uptakes were determined at 195 and 300 K by the difference between the initial isotherm and a second one obtained after a I-h evacuation step a t the adsorption t e m p e r a t ~ r e . ~ ~ Hydrogen desorption following a cooling step and evacuation at 300 K was used to overcome activated adsorption and obtain H, uptakes.53 Isothermal energy changes upon adsorption of C O a t 300 K were measured by using a modified Perkin-Elmer DSC-2C differential scanning ~ a l o r i m e t e r . ~The ~ - ~only ~ additional modification made was the installation of a small glovebox which allowed the transfer of catalyst into the DSC without air exposure. Identical, parallel pretreatments were conducted in the adsorption and calorimetric systems prior to the subsequent calorimetric and adsorption measurements on two samples derived from a single batch of catalyst. Details of the experimental procedure are described e l ~ e w h e r e . ~ ~ - ~ ~ Results
Chemisorption. Two Fe/C catalysts were prepared with metal loadings of 8.7 and 6.4 wt 5% Fe as determined by ashing of the sample. The chemisorption results are summarized i n Table I . Experiments with pure carbon showed that no irreversible uptakes (5 I ) Krishnan, K.; Ferraro, J . R. Fourier Transform Infrared Spectroscopy: Techniques (lying Fourier Transform Interferometry: Academic Press: New York, 1986; Vol. 3. p 149. ( 5 2 ) Emmett. P. H.: Bvnauer. S. J . Am. C h e w Soc.1937. 59, 310. (53) Amelse, J. A , ; Schwartz. L. H.; Butt. J. B. J . Cuta1. 1981. 7 2 . 95. (54) Vannice. 1.1. 4 . ; Sen. B.: Chou. P. Rec. Sci. Instrunl. 1987. 58, 647. (55) Sen. B.: Chou. P.: Vannice. M.A. J . Coral. 1986. I O / , 517. (56) Chou, P.: Vannice. M . A . J . Catal. 1987. 104. I .
4160
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989
Venter and Vannice
TABLE I: Chemisorption Measurements on 8.7% and 6.4% Fe/C Catalysts Derived from Fe,(CO)12
fresh catalyst uptake, pmol/g of catal LTR HZ(0-200 "C) CO(195 K ) CO(300 K ) CO(300 K)/C0(195 K) HTR H,(O-400 "C) CO(195 K) CO(300 K) CO(300 K)/C0(195 K)
dispersion' D,,c nm 8.7% Fe/C"
57 450 583
used catalyst uptake, pmol of CO/g of catal dispersion'
D,,c nm
0.7 0.58
1.6
0.38 1.3
84 469 395
0.11
0.60 0.25
1.6
86 37 I 368
0.8
0.1 1 1 .o
0.48 0.24
2.0
6.4% Fe/Cb LTR HZ(0-200 "C) CO( I95 K) CO(300 K) CO(300 K)/C0(195 K)
57 434
0.10
578
0.50
0.66
I .2
1.3
HTR
H(0-400 "C) CO( 195 K) CO(300 K) CO(300 K)/C0(195 K) 1551 pmol of Fe/g of catalyst.
79 403 367
0.14 0.70 0.32
1,3
83 356 312
0.14 0.62 0.27
1.5
0.9
0.9 1150 pmol of Fe/g of catalyst. 'Assuming C0:Fe = 1.2 at 195 K
occurred.",42 Highly dispersed Fe/C catalysts were obtained, and the C O uptakes at 195 and 300 K were in reasonable agreement, as ratios of these values were between 0.8 and 1.3 for both catalysts after either LTR or H T R . The H2 uptakes were significantly lower than the C O values, even though a n improved desorption technique was used to help overcome activated a d ~ o r p t i o n .Still, ~~ these results were an improvement over previously used static H 2 as highly dispersed adsorption measurements a t 300 K,5~6~38~57~58 Fe catalysts exhibit H2adsorption even more activated than that The two catalysts displayed similar on bulk Fe behavior in that the H, adsorption values increased after H T R , probably because of the larger particles present. CO chemisorption on the used samples decreased by IO-20% for both catalysts, slightly more than that observed for carbon-supported Os and Ru.43,44The H, adsorption values, however, were unchanged. No standard method exists for the measurement of iron surface area by chemisorption methods, but good agreement among different techniques has been obtained for estimates of larger crystallite sizes by assuming an adsorption stoichiometry of CO:Fe, = 1:2 for adsorption a t 195 K.6339 This ratio has also given consistent results for highly dispersed Fe on M g 0 . j 8 C O chemisorption a t 195 K on both catalysts indicated initial dispersions ranging from 0.6 to 0.8, with final values between 0.5 and 0.6 following the kinetic measurement. These dispersions correspond to Fe crystallite sizes below 2 nm. The low hydrogen uptakes give H/FetOtalratios between 0.07 and 0.14, but the H 2 / C 0 ( 1 9 5 K) uptake ratio has been shown to be a strong function of Fe crystallite size and these low ratios provide additional evidence for the presence of Fe particles in this size r a r ~ g e . ~ ~ -The ~ ' stoichiometry of H adsorption on small Fe particles has not yet been established and, in addition, with carbon supports H spillover can occur to obviate meaningful dispersion measurement^.^ W e have chosen C O adsorption at 195 K to estimate Fe particle size because no X R D pattern for Fe was obtained, T E M studies showed small, undiscernible Fe particles and no large particles, while MES results ( 5 7 ) Emmett, P. H.; Brunauer, S. J . A m . Chem. SOC.1937, 59. 319. ( 5 8 ) Topsoe, H.; Topsoe, N.; Bohlbro, H.; Proceedings o f t h e 7th International Congress on Catalysis, Tokyo, 1980 Elsevier: Amsterdam, 198 1;
p 247. (59) Yoon, K . J.; Walker, P. L.; Mulay. L. N.; Vannice, M . A. Ind. Eng. Chem. Prod. Res. Dea. 1983, 22, 519. ( 6 0 ) Hayward, D. 0.;Trapnell, B. M. W. Chemisorption. 2nd ed.; Butterworths: Washington, 1964. (61) McDonald, M. A.; Storm, D.A.; Boudart, M. J . Cutal. 1986, 102. 386.
TABLE 11: Isothermal (300 K), Integral Heats of Adsorption of CO on Reduced, Carbon-Supported Fe Derived from Fe3(CO)
catalyst
energy change, mcal/g of catal
chemisorption, pmol of CO/ Qad, kcal/ Coed/ g of catal (g mol of CO) FelOl,, 8,7% Fe/C
sample I no. I no. 2 no. 3
8573 6839 5254
632 458 379
13.6 14.9 13.9
8432 6858
632" 458"
13.3
0.4 I 0.30 0.24
sample 2 no. 1 no. 2
15.0
6.4% Fe/C sample I
no. 1 no. 2 no. 3 sample 2 no. 1 no. 2
6848 5457 4646
389 3IO 350
17.6 17.6 13.3
0.34 0.27 0.30
6614 5254
402 368
16.5 14.3
0.35 0.32
15.0 f 1.6b
"Assumed from sample I . bAverage value with standard deviation. also showed the presence of only small superparamagnetic Fe particles.2x62 Large Fe crystallites could readily be observed when present .2,62,63 Calorimetric Measurements. The integral heats of adsorption of CO a t 300 K are shown in Table I1 together with the CO chemisorption on each Fe/C catalyst. The chemisorption values typically decreased when successive chemisorption cycles were done on a single sample, as shown most clearly by the first 8.7% Fe/C sample. This decrease is presumed to be due to sintering during the extended H T R step; however, it had no discernible effect on the measured CO heat of adsorption and ten measurements of four samples gave an average value of 15.0 f 1.6 kcal/mol C O . DRIFTS Measurements. Prior to decarbonylation experiments, the spectrum of Fe,(CO),, dispersed on this carbon was recorded in flowing He to ensure the presence of intact Fe carbonyl clusters. The initial spectrum for each sample always showed the presence of both Fe3(CO),, and Fe(CO)S, and the relative abundance of (62) Chen, A . A.; Phillips. J.; Venter, J . J.; Vannice. M . A. J . C a r d . in press (63) Chen, A . A,: Vannice. M . A,; Phillips, J. J . Cora/.. in press
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 4161
Decomposition of Carbon-Supported Fe3(C0)12Clusters TABLE III: Reported IR Bands for Iron Carbonyl Clusters in Solution
cluster Fe(CO)5 Fe(CO),Lb Fe(CO)&Le Fe(COj;L2 [Fe(C0),l2[ HFe(CO),]FeAC0)9 [ Fe2(CO),] 2[HFe2(CO),]Fe3(C0)12 Fe3(CO)IIL Fe3(CO)loL2 Fe3(C0)9L3 [ Fe3(CO)I [ HFe3(CO)I,I[ Fe4(CO),,l2-
[HFe,(CO),,]-
wavenumber,' cm-I
ref
201 9-2023, 1995-2000
1923-1930 2013-2026 2008-2011,
64-7 1 65, 66, 66 66, 71, 73-76 73, 75, 78, 79 73, 77 73, 77 65, 67, 72, 86
1995-2003,
72, 86
1814-1829,
72, 86
1900-19 13,
67, 73, 77, 87
1970-1990,
28, 67, 73, 77, 84, 87-92 73, 77
2040-2099, 1970-2018, 1933-1963 2099. 2018. 2000
1877-1925 2005, 1962, 1730-1786 2008, 19 10-19 14, 1840-1897 2064-2065, 2015-2037, 1828-1847
1916-1 920, 1852-1 866 2068, 2048, 1980-1997, 2097-2103, 2043-2056, 2070-2092, 2025-2035, 1993-1 998 2062-2068, 201 3-2023, 1957-1978 2040-2049, 1974-1986, 1770-1788 2000-201 0, 1931-1943, 18 80- 18 84 2062-2075, 2000-2012, 1942-1961 2067-2068, 2021-2030,
1990-2009,
1938-1950 2020-2031, 1980-1984, 1961, 1936-1 942
71, 72 72 77
80-85
0
100 150 200 250 300 TIME (mid
Figure 2. Log (relative intensities) versus time for carbon-supported Fe carbonyl clusters in He at 303 K (8.7% Fe/C): 0, Fe3(C0)12(band at 2047
73,77
50
cm-I); 0,Fe(CO)5 (band at 2000 cm-I).
'Strongest bands are in italics. bAxial L. cEquatorial L. Fe(C0)5
1
TIME (mid
n
TIME (mid
0
h
N
0
0 0 0
8
45
63
c
.-0 .-v) >
a
105
111
Y 7
249
i,
y
165
A
261
536 2lR0
21c10
2060
2020
1980
1940
1900
Wave n u m b e r Figure 1. Spectra of carbon-supported Fe3(C0)12(8.7% Fe/C) in He at 303
K. The principal IR bands are at 2047, 2018, and 2000 cm-I.
these two species changed with storage time, with Fe3(C0)12slowly decomposing to form Fe(CO), and zerovalent iron particles.2-62 These initial spectra always exhibited bands a t 2047, 2020, and 2000 cm-I, in good agreement with the reported frequencies in Table 111 for Fe3(C0),2and Fe(CO), in solution.w2 The sample (64) (65) (66) (67)
Calderazzo, F.; L'Eplattenier, F. Inorg. Chem. 1967, 6, 1220. Ballivet-Tkatchenko, D.; Coudurier, G. Inorg. Chem. 1979, 18, 558. Haas, H.; Sheline, R. K. J . Chem. Phys. 1967, 47, 2996. Hanson, B. E.; Bermeaster, J. J.; Petty, J. T.; Connaway, M. C. Inorg.
Chem. 1986, 25, 3089. (68)
Jones, L. H.; McDowell, R. S.;Goldblatt, M.; Swanson,B. I. J. Chem.
Phys. 1972, 57, 2050. (69) (70) (71)
Cataliotti, R.; Foffani, A,; Marchetti, L. Inorg. Chem. 1971, 10, 1594. Guglielminott, E.; Zecchina, A. J . Mol. Catal. 1984, 24, 331. Bigorgne, M. J . Organomet. Chem. 1970, 24, 211.
2060
2020
1980
1940
1900
Navenu m b e r Figure 3. Spectra of carbon-supported Fe3(C0),2(8.7% Fe/C) in H2 at 292 K. The principal IR bands are at 2047, 2018, and 2000 cm-I.
was then heated or cooled rapidly in either He or H 2 to a chosen temperature and the decarbonylation process was monitored to (72) (73)
Shojaie, A,; Atwood, J. D. Organometallics 1985, 4, 187. Farmery, K.; Kilner, M.; Greatrex, R.; Greenwood, N. N. J . Chem.
SOC.A 1969, 2339. (74)
Alich, A.; Nelson, N. J.; Strope, D.; Shriver, D. F. Inorg. Chem. 1972,
11, 2976.
(75) Edgell, W. F.; Huff, J.; Thomas, J.; Lehman, H.; Angell, C.; Asato, G. J. Am. Chem. SOC.1960, 82, 1254. (76) Abel, E. W.; Stone, F. G. A. Q. Rev. 1969, 23, 325. (77) Edgell, W. F.; Yang, M. T.; Bulkin, B. J.; Bayer, R.; Koizumi, N. J . Am. Chem. SOC.1965, 87, 3080. (78) Fletcher, S. C.; Poliakoff, M.; Turner, J. J. Inorg. Chem. 1986, 25, 3597. (79) Kristoff, J. S.; Shriver, D. F. Inorg. Chem. 1974, 13, 499.
(80) Lazar, K.; Matusek, K.; Mink, J.; Dobos, S.; Guczi, L.; Marko, L.; Reiff, W. M. J . Catal. 1984, 87, 163. (81) Cotton, F. A.; Wilkinson, G. J . Am. Chem. SOC. 1957, 7 9 , 752. (82) Sheline, R. K. J. Am. Chem. SOC.1951, 7 3 , 1615.
4162
The Journal of Physical Chemistry, Vol. 93, N o . I O , I959
Venter and Vannice
1
0.1 T
I
5
0.01 :
X
4
I
r
z
-
0.001 :
0.0001 0.29
0.1
0
..
100
200
300
400
500
TIME b i n )
Fe(CO), (band at 2000 cm-I).
TABLE IV: Activation Energies and Rate Constants for Decomposition of Fel(CO),, and Fe(COL on Carbon in H, or He
band. cm-‘ 2047 2000
cluster Fe,(CO),, Fe(CO),
rate const at 350 K, min-’ He HI 2.3 0.52 0.15 0.37
0.33
lfr
Figure 4. Log (relative intensities) versus time for carbon-supported Fe carbonyl clusters in Hz at 292 K (8.7% Fe/C): 0 , Fe,(CO),, (band at
2047 cm-I);
0.31
He
H2 21.3 15.7
obtain spectra such as those shown in Figure 1 for the run in H e a t 303 K . The expected bands of Fe3(C0),* a t 2047 and 2020 cm-’ and of Fe(CO)5 at 2020 and 2000 cm-I are noted. Decomposition of Fe,(CO),, clearly leads to fragmentation and the formation of Fe(CO), as a primary product, and it occurs to a significant extent even a t room temperature. Because of the stability of Fe(CO),, the decarbonylation of both these species could be followed quantitatively and rate constants were determined from first-order rate plots, such as those in Figure 2, which nicely demonstrate the first-order behavior for the duration of the decomposition run. The decarbonylation of Fe3(CO),,, monitored by the decreasing intensity of the 2047-cm-I band, is more rapid than the disappearance of the 2000-cm-’ band of Fe(CO)5. The decarbonylation of Fe3(CO),2/Cin H2 proceeded in a similar manner, as shown in Figure 3 for the run at 292 K, and no formation of stable hydrido clusters was observed, in ,.~~ contradistinction to the Os and R u MCC’s on ~ a r b o n . ~The decarbonylation of Fe(CO), was also monitored by the change in intensity of the 2000-cm-’ band, as shown in Figures 1 and 3; but this was complicated by the fact that Fe(CO), is formed during the deconiposition of Fe3(C0)12. However, the kinetics for Fe(CO), could conveniently be determined from the 2000-cm-I band when the concentration of F e 3 ( C 0 ) i 2was low, i.e., when little overlap of the 2020-cm-’ band occurred and well-behaved firstorder plots were obtained, such as those shown i n Figure 4. (83) Cotton, F. A.: Hunter, D. L. Inorg. Chim. Acta 1974. 11, L9. (84) Hugues, F.; Smith, A. K.: Ben Taarit, Y . ;Basset. J. M.; Commereuc. D J . Chem. SOC.,Chem. Commun. 1980, 68. ( 8 5 ) Pierantozzi, R.; McQuade, K. J.: Gates, B. C.: Wolf. M.; Knozinger, H J . A m . Chem. SOC.1979, 101%5436. ( 8 6 ) Grant. S. M.: Manning, A . R. Inorg. Chim. Acta 1978. 31, 41. (87) Psaro. R.; Dossi, C.; Lgo, R. J , Mol. Carol. 1983, 21, 331. (88) Anders. U : G r a h a m , W . A.G. J . Chem. Soc.. Chem. Commun. 1966, 291. ( 8 9 ) Hodali. H. A.: Shriver. D. F.: Ammlung. C . A . J . A m . Chem. SOC. 1978, 100. 5239. (901 Effa. .I B. N.: Lieto, J.: Aune. J. P. Inorg. Chim. Acta 1982. 65, L105. (91) Wilkinson, J . R.; Todd. L. J . J . Orgonomet. Chem. 1976, 118, 199. ( 9 2 ) Iwamoto, V , :Nakamura. S. I . : Kusano. L. ./. Phys. Chem. 1986. 90, 5244.
0.37
0.39
.,
Figure 5. Arrhenius plots for the decarbonylation of carbon-supported Fe carbonyl clusters: 0, Fe,(CO)I, in He (band at 2047 cm-’); 0, Fe(CO), in He (band at 2000 cm-I); 0 , Fe,(CO),, in H, (band at 2047 cm-I); Fe(CO), in H, (band at 2000 cm-I).
4 Intensity x
activation energy. kcal/(g mol) 17.9 15.3
0.35
(liK+100)
-
i
r
0
0
; 0
C
.-0 In ._
>
6 7-
v
In .-c c
3
7 Y
I
/ \
1,
2200
2120
2040
1960
1880
1800
Wave n u m b e Y Figure 6. Spectra of CO adsorbed on Fe crystallites: (a) CaF,-diluted Fe,(CO),, at 300 K following HTR at 673 K for 3 h; (b) 8.7% Fe/C at 195 K in flowing He following HTR at 673 K for 16 h; (c) sample b after heating to 300 K i n flowing He and CO (in 1 1 Torr of CO): (d) 8.7% Fe/C at 195 K in flowing H, following HTR at 673 K for 16 h; (e) Sample d after heating to 300 K in flowing H, and CO (in 1 1 Torr of CO)
The activation energies of decarbonylation for these two clusters in either H 2 or H e were determined from the Arrhenius plots shown in Figure 5 . As indicated in Table IV, both the activation energy and the rate constant for decomposition of Fe3(C0)12are higher than their counterparts for Fe(CO), in either He or H2. The ambient gas had little effect on either the activation energy or the rate constant of Fe(CO), decarbonylation, whereas both parameters for Fe,(CO),, showed some dependence on the surrounding gas. The difference in rate constants for Fe3(C0)12under He or H2 is somewhat exaggerated in Table IV because these are extrapolated values at 350 K to allow comparison with the reported ~ ~ temperature region values for Os and R u on ~ a r b o n . ~ I,n. the
Decomposition of Carbon-Supported Fe3(CO) 1 2 Clusters investigated (270-320 K), the rate constants were similar. The adsorption of C O on the Fe particles formed after complete decomposition was also studied. To verify that chemisorbed C O could be observed and to obtain the C O band positions directly, solid Fe3(C0)12was diluted in CaF, with no carbon present, heated to 673 K in flowing H 2 for 3 h, cooled to 300 K, and exposed to 11 Torr of C O in flowing H2. The gas-phase C O was flushed with H2, and spectrum a in Figure 6 was then obtained. A very broad band was found with maximum intensity occurring in the 19502000-cm-' region, indicative of linear and bridged C O on bulk metallic Fe surfaces and very similar to spectra reported for Si02-supported Fe.48 Note that this spectrum was much more intense than the subsequent spectra. Following the decarbonylation run, the Fe/C sample was heated to 473 K in flowing H2, reduced at this temperature in H 2 for 2 h, and flushed in H 2 or He a t 473 K for 30 min prior to the adsorption of C O first at 195 K and then at 300 K. Although numerous attempts were made, after this LTR step no detectable adsorbed C O species were observed a t either temperature, either in the presence or absence of gasphase CO, even though the C O chemisorption studies clearly showed that significant uptakes occurred. After these cycles, the catalyst was heated to 673 K in flowing H2, reduced for 16 h, and flushed in H 2 or He at 673 K for 30 min prior to cooling, and spectra were obtained first at 195 K and then at 300 K, as shown in Figure 6. Spectrum b was recorded at 195 K after cooling the sample in flowing He to 195 K, exposing to 11 Torr of C O in flowing He, and flushing the gas-phase C O with He. A single peak was observed at 2050 cm-I; however, a similar peak existed on a carbon CaF2 blank; consequently, it is due to C O weakly chemisorbed on the carbon. Spectrum c was recorded in the presence of gas-phase C O after sample b was warmed to 300 K in flowing C O ( I 1 Torr) and He (749 Torr), and it illustrates that reformation of Fe(CO)S, a species not detected a t 195 K, had occurred. Purging in He removed all bands. Spectra d and e were collected under H2 by using the procedure just described, and they show that Fe(CO), is formed when these Fe particles are exposed to CO, independent of the ambient gas, which is consistent with Mossbauer spectroscopy studies following this same experimental procedure.2 Purging in H 2 again removed all bands. The catalyst was also investigated under reaction conditions after it was subjected to HTR and exposed to 11 Torr of C O and 749 Torr of H 2 at 473 K for 2 h, 523 K for 2 h, and 573 K for 2 h. The spectra obtained at this very high H 2 / C 0 ratio42revealed no detectable C O or CH, bands, but cooling to 300 K in flowing H 2 and C O produced the spectrum in Figure 7. Bands were observed at 2961, 2932, and 2874 cm-I, which are consistent with observed values of 2875 and 2960 cm-' for methyl groups and 2850 and 2925 cm-' for methylene g r o ~ p s . ~ The ~ - ~methyl ~ groups appear to be present in larger concentrations.
+
Discussion Numerous investigations to characterize iron carbonyl clusters have been reported.29~35*47~64-92,96-109 Many have pertained to (93) King, D. L. J . Carol. 1980, 61, 77. (94) Tipson, R. S. Infrared Spectroscopy of Carbohydrares; NBS Monograph 1 IO; Government Printing Office: Washington, DC, 1968. (95) McBreen, P. H.; Erley, W.; Ibach, H. Surf. Sci. 1984, 148, 292. (96) Zwart. J.; Snel, R. J . Mol. Cafal. 1985, 30, 305. (97) Boszormenyi, 1.; Dobos, S.; Lazar, K.; Schay, 2.; Guczi, L. Surf. Sci. 1985, 156. 995. (98) Boszormenyi, 1.; Dobos, S.; Guczi, L.; Marko, L.; Lazar, K.; Reiff, W. M. Proc. 8rh I n t . Congr. C a r d 1984, 5 , 183. (99) Dobos, S.; Boszormenyi, I.; Mink, J.; Guczi, L. Inorg. Chim. Acfa 1986, 120, 135. (100) Dobos, S.; Boszormenyi, 1.; Mink, J.; Guczi, L. Inorg. Chim. Acra 1986, 120, 145. (101) Alper, H.; Copal, M. J . Chem. S o t . , Chem. Commun. 1980, 821. (102) Basset, J . M.; Choplin, A. J . Mol. Cafal. 1983, 21, 95. (103) Rojas, D.; Bussiere, P.; Dalmon, J. A,; Choplin, A,; Basset, J. M. Surf. Sci. 1985, 156. 516. ( 104) Ballivet-Tkatchenko, D.; Coudurier, G.; Mozzanega, H.; Tkatchenko. I . Fundam. Res. Homogeneous Catal. 1979, 3, 257. (105) Bein, T.; Jacobs, P. A. J . Chem. Soc., Faraday Discuss. 1983, 79, 1819.
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989 4163
h
cu
0 0
8
8 II
C
.-0
.-v) >
i Y 7
v) .-c C
3
2 Y
5100
3050
3000
2950
2900
2850
2800
Wave n u m b e r Figure 7. IR spectrum of 8.7% F e / C catalyst after cooling in flowing H, and CO from reaction conditions to 300 K.
clusters in solution although some have focused on the interaction of iron MCC's with oxide surfaces. In addition, studies of fragmentation and substitution reactions, again involving iron carbonyl clusters in solution, have been conducted."*12' However, none has determined the kinetics of decomposition for any iron cluster, a situation that may be due to experimental limitations associated with the presence of the solvent. The approach used here allows the clusters to be dispersed on the carbon surface, the solvent to be removed by evacuation, and the IR spectra to be obtained as decomposition proceeds thereby allowing us to monitor the bond-breaking processes. The interaction of iron MCC's with oxide supports has been reviewed T o summarize briefly, it has been found that Fe2Ru(CO),2and Fe3(C0),2 decompose on hydroxylated SiO, to form iron ~ x i d e , ~but ~.'~ impregnation of Fe(CO)S and Fe3(C0)12has also produced intact Hydroxylated A1203invariably leads to the rapid transformation of Fe3(C0)12to either [HFe3(CO)l,]-29~67~84~101~102 or iron oxide^?^^'^ and when Fe(CO), is the precursor, formation of [HFe3(C0)l,]-29,84 as well as [HFe4(C0),3]-67has been reported. On hydroxylated MgO, Fe(CO)S and Fe3(CO),, form [HFe,(CO)' l]-,29*84,102 whereas on hydroxylated Na-Y zeolites Fe(CO), is physically adsorbed while Fe2(C0)9 and Fe3(CO),, form [HFe3(CO),l]-,92which is relatively stable and has been . (106) Nagy, J. B.; van Eenoo, M.; Derouane, E. G. J . Caral. 1979, 58, 230. , Discuss. 1984, 80, (107) Bein, T.; Jacobs, P. A . J . Chem. S O C . Faraday
1391. (108) Bein, T.; Schmidt, F.; Gunsser, W.; Schmeister, G. Surf. Sci. 1985, 156, 57. (109) Ballivet-Tkatchenko, D.; Tkatchenko, J. J . Mol. C a r d 1981, 13, I . ( 1 IO) Poe, A. J. Metal Clusters; Wiley: New York, 1986; p 53. (1 I I ) Cetini, G.; Gambino, 0.;Sappa, E.; Vaglio, G. A. Afri Accad. Sci. Torino 1965, 101, 855. (112) Shojaie, R.; Atwood, J. D. Inorg. Chem. 1987, 26, 2199. ( I 13) Candlin, J. P.; Shortland, A. C. J . Organomef. Chem. 1969, 16, 283. ( I 14) Weitz, E. J . Phys. Chem. 1987, 91. 3945. ( 1 15) Angelici, R . J.; Siefert, E. E. Inorg. Chem. 1966, 5 , 1457. (116) Schumann, H.; Opitz, J. J . Organomef. Chem. 1979, 166, 233. (117) Kumar, K. J . Organomef. Chem. 1977, 136, 235. ( 1 18) Siefert, E. E.; Angelici, R. J. J . Organomer. Chem. 1967, 8 , 374. (1 19) Johnson, B. F. G.; Lewis, J.; Twigg, M . V. J . Chem. SOC.,Dalron Trans. 1975, 1876. (120) Fox, J. R.; Gladfelter, W. L.; Geoffroy, G. L. Inorg. Chem. 1980, 19, 2514. (121) Baev. A. K.; Conner, J. A,; El-Saied, N . I.; Skinner, H. A. J . Organomef. Chem. 1981, 213, 151.
4164
The Journal of Physical Chemistry, Vol. 93, No. 10, 1989
Venter and Vannice
TABLE V: Reported Heats of Adsorption of CO on Fe Surfaces heat of adsomtion. kcal/mol of CO sample method" 8-0 8 = high integral 5% Fe/Si02 Cal 24 Fe Cal 27 5 21 Fe film Cal 46 Fe film Cal 32 15 Fe Cal 20 IO Fe Cal 31 2.2 6 Fe film 38.2 IO 25 Fe( 1OO)/C+O Is0 21.3 17.9 Fe( 1OO)/C Is0 18.4 18.4 Fe( I OO)/O Is0 17.7 17.7 Fe(l lO)/C+O Is0 19.1 10.5 Fe( 1 I O)/C Is0 19.8 15.8 11.4 2.9 Fe/A1203 Cal Fe/MgO Cal 26.3 6 17 Fe( 100) TPD 26.6
Fe( 100) Fe( 100)
TPD TPD
comments
ref
6'
E
0.6, reduced Fe304
6'
N
0.6
153 154 155 155 156 157 158 159
reduced Fe304 reduced Fe304 8 6' 6' 6' 6' 8
N
0.3 0.3, C and 0 precovered 0.3, C precovered 0.3, 0 precovered 0.3, C and 0 precovered
N
0.3, C precovered
E N E N
160 58 144-147 144-147 144-147
= 1; d, = 3.6 nm highly tilted CO 8hlgh
bridged CO
18.0
weakly held CO
12.8
Cal, calorimetry; Iso, isosteres; TPD, temperature-programmed desorption attached to functionalized polymersg0 and hydroxylated A120367 without decomposing to Fe(CO)5. The decomposition of these clusters in vacuum at high temperatures produces iron oxides on hydroxylated A1203 and MgO and a mixture of metallic and oxidized Fe on Si0,,67,80.84.97,98,~0~ Due to the ease of oxidation of Fe clusters, dehydroxylated oxide supports have been used. Impregnation of Si02with Fe3(C0),, led to physisorbed Fe(CO)S and Fe3(C0)12,'03while Fe(CO), and Fe3(C0),, on MgO and highly dehydroxylated A1203gave Fe(CO)s_,L, species, where L is a ligand such as a hydroxyl The impregnation of H-Y zeolites with Fe3(C0),,, Fe,(C0)9, or Fe(CO)5 led to the formation of Fe(C0)5-nL:s,104 or Fe(CO)5,105J06while only F e 2 ( C 0 ) 9 decomposed in Na-Y zeolites.92J07-'09The complete decarbonylation of these clusters gave and Mg0,84but metallic iron oxides on these two zeolites65*'04~'07.108 Fe on AI2O3,lo2 Mg0,709843102 and Na-Y zeolite^.^^^-^^^ In summary, Fe(CO)S or Fe3(C0)12on oxide supports frequently forms , and [HFe3(CO),I]-, but Fe(CO)5-,L,, [Fe(C0)4C02]2-Mg2+ [HFe4(C0)13]-have also been observed in addition to the physisorbed precursors. No studies of Fe carbonyl clusters have been reported on oxygen-free support surfaces. Thus the applicability of our DRIFTS system to characterize clusters on clean carbon surfaces allows the first opportunity to study decarbonylation processes in the absence of ligands which can alter the decomposition chemistry. The decomposition of F e 3 ( C 0 ) 1 2and Fe(CO)5 on carbon in He or Hz is shown in Figures 1 and 3, respectively, and the spectra imply that similar decarbonylation processes occurred. The initial spectra show that both Fe3(C0),, and Fe(CO), were present and indicate that the dodecacarbonyl clusters partially decomposed during catalyst preparation and continued to slowly decompose to Fe(CO)5 at room temperature, as also detected by Mossbauer effect spectroscopy ( MES)., N o subcarbonyl species or hydrido clusters were detected throughout the decarbonylation process under either H, or He, in contrast to the behavior of O S ~ ( C O ) , ~ and R u ~ ( C ~ ) ~ ~ . ~ ~ . ~ ~ The most probable pathway for either substitution or fragmentation reactions of Fe,(CO),, in solution involves the initial removal of a CO ligand from the Fe3(CO),, cluster.ll"-"z This initial step can be followed by either substitution or fragmentation, depending on the nature of the other nucleophilic ligand (L) that is present, as formation of Fe,(C0)9L3 is favored when L is triphenyl phosphite (P(OPh),) or trimethyl phosphite (P(OMe),) whereas triphenylphosphine (PPh,) or tributylphosphine (PBu,) leads to fragmentation into mononuclear specie^.^^,^^ This latter reaction proceeds more readily than with Os or Ru dodecacarbonyls due to the weaker metal-metal bonding in Fe,(CO)l,.l'o Second- and, especially, third-period MCC's in general retain the M-M bond following substitution whereas M-M bond scission prior to substitution occurs for some first-period MCC's such as
Fe(C0)aL + Fe(C0)SLz
Figure 8. Proposed reaction network for substitution and fragmentation of Fe3(C0)12clusters. TABLE VI: Activation Energies and Rate Constants for Nucleophilic Substitution into Fe3(CO)]*: Fe3(CO)11-nLn+l + CO Fe3(C0),,-,L, + L
-
AH*, n
L
0
P(OMe), P(OPh)3 PPhj P(CMe)3 P(Bu)3
0 0 0 0 0 1 1 a
PPh?
PPh3 P(OMe)3
solvent
rate constant k,,
kcal/ mol
T, K
29.5 29.1 29.1 29.1 29.1 22.5 29.1" 29.1"
303 303 303 303 303 303 303 303
s-l X
k , at 350 K,
lo5
4.0 5.6 6.3 5.7 5.0 3.3 22 17
min-'
ref
1.7 2.2 2.5 2.3 2.0 0.3 8.7 6.7
112 72 72 72 72 117 72 72
Assumed. *Products were Fe(C0)4L and Fe(CO),L2 exclusively.
Mn2(CO)lo."3.1'4For Fe3(C0),,, however, C O removal prior to Fe-Fe bond scission appears more probable, and an earlier report that Fe,(CO),, underwent direct fragmentation in the presence of high concentrations of PPh,l" is probably due to the fact that Fe,(CO), , L undergoes very rapid fragmentation to Fe(C0)4L and Fe(CO),L2, with this rate thus being equal to the initial rate of substitution into Fe3(C0)12.72Regardless, interaction of Fe3(CO),, with nucleophiles most frequently leads to a mixture of Fe3( C O ) l l L , Fe3(CO)loL2,Fq(CO)gL,, Fe(C0)4L, and Fe(CO),L2.72-'12-115-117 This behavior can be described by the reaction network shown in Figure 8.'1° The initial removal of C O can proceed via one of two mechanisms-a dissociative first-order mechanism in which the removal of C O precedes the attachment of L or an associative second-order mechanism in which the incorporation of L precedes the removal of CO. The total rate of substitution is the sum of the two paths k [ M C C ] = ( k , + k2[L])[MCC].i'o The relative importance of k , and k , depends greatly on the nature and concentration of L, which also determines the re!ative amounts of Fe(CO),, Fe(CO),L, and Fe(C0 ) 3 L 2 eventually formed.'I0 Reported rates for the initial substitution of L into Fe3(C0),, in solution are summarized in Table VI. The first-order k l values at 350 K are typically near 2.0 min-I, which is close to those obtained in this study for the decomposition of Fe3(C0)12in He (0.52 min-I) and in H 2 ( 2 . 3 min-I). This similarity strongly suggests that the slow step in the decomposition of Fe,(CO),,
SCHEME I: Thermodynamic Path To Estimate the Minimum Activation Energy of Decarbonylation of Fe3(CO),,'
TABLE VII: Activation Energies and Rate Constants for Nucleophilic Substitution into Fe(CO)$ L Fe(CO),+L,,+, CO Fe(CO)+.L. rate constant k , at
+
L solvent P(C6H5), decalin
PPh, a
decalin
kcal/ mol 42.5 42.5'
-
+
k,, T, K 452 433
s-I X
26 3
3nFeco,
I
350 K,
min-'
-
The Journal of Physical Chemistry, Vol. 93, NO. 10, 1989 4165
Decomposition of Carbon-Supported F e 3 ( C 0 ) 1 2Clusters
ref
1.6 X lo-* 118 1.4 X 119
Assumed.
clusters on carbon is the initial loss of a C O ligand, analogous to the dissociative substitution mechanism. The Fe(CO)S formed from Fe3(CO),, decomposition also decarbonylated following first-order kinetics. Substitution reactions with Fe(CO)s in solution have apparently not been studied, and the most analogous reaction reported is substitution into monosubstituted iron pentacarbonyl to form a disubstituted species, as shown in Table VII. These rates imply that the rate of substitution into Fe(CO)S via a dissociative mechanism is orders of magnitude lower than that for Fe3(CO),,. The increased stability of Fe(CO), relative to Fe3(C0),2is unique to this family of MCC's because R U ( C O ) and ~ R U ~ ( C O )as~ well ~ , as O S ( C O ) ~ and O S ~ ( C O ) ,undergo ~, nucleophilic substitution at comparable rates.72 This explains why Fe(CO)S is observed during the decomposition of Fe3(C0)12but R U ( C O )and ~ Os(CO), clusters were not detected during the decarbonylation of R U ~ ( C O )and , ~ Osj(C O ) z. 43*44 Phillips and Dumesic studied the effect of outgassing a Grafoil carbon support on both the rate of Fe(CO)S decarbonylation and the final Fe crystallite size.I7 From their data, rate constants at 383 K of 0.0009 min-l after outgassing at 383 K and of 0.07 mi& after outgassing a t 573 K can be calculated. Extrapolation of the rate under H e to 383 K in Figure 5 gives k = 1.0 min-]; consequently, a trend of higher rates with higher outgassing temperatures on carbon is apparent. These authors attributed the enhanced activity to the formation of "active" sites on this graphitic carbon during outgassing. The activation energy for the decarbonylation of carbon-supported F e 3 ( C 0 ) , 2was 17.9 kcal/mol in H e and 21.3 kcal/mol in H2, while it was nearly constant at 15.5 f 0.2 kcal/mol for Fe(CO)5. When these values are compared to those obtained in solution for substitution reactions-23-30 kcal/mol for Fe3(CO)12 and 42.5 kcal/mol for Fe(C0)s118,119-they are similar but somewhat lower for the dodecacarbonyl whereas they are markedly lower for the pentacarbonyl. A low activation energy of 13 kcal/mol has been reported for the fragmentation of H 2 R ~ 4 ( C 0 ) 1 3 in hexane to form R u ~ ( C O ) ,R~ U , ( C O ) ~and , H2,IB and the weaker Fe-Fe bondsl2I might be expected to give low values. Several studies have investigated the decomposition of supported Fe(CO)5. Bein et al. reported that Fe(CO)S decomposed on highly dehydroxylated zeolites with an activation energy of 24 f 5 k ~ a l / m o l . ' ~A~similar experiment by Brenner and Hucal with AI2O3yielded an activation energy of 30 kcal/mol,122and Carlton and Oxley reported a value of 20 kcal/mol for Fe(CO)s decarbonylation on steel, as determined by TPD.'23 Consequently, our values near 15.5 kcal/mol are quite consistent. It is very likely that the presence of metallic Fe in these carbon catalysts, formed during the impregnation step,2 lowers the activation energy, as noted by M i t t a ~ c h , 'and ~ ~ provides nucleation sites for the decomposition of Fe(CO)S, as proposed by Phillips and Dumesic.17 The lower activation energies can be explained by the subsequent formation of Fe-Fe bonds to form small Fe crystallites after C O removal. A thermodynamic pathway for the decarbonylation of these MCC's to metallic Fe and gas-phase C O can be constructed to estimate minimum activation energies for these endothermic processes, as illustrated in Schemes I and 11. Two limiting cases can be considered-the formation of Fe-Fe bonds
,
(122) Brenner, A.; Hucul, D.A. Inorg. Chem. 1979, 18, 2836. (123) Carlton, H . E.; Oxley, J. H.J . Am. Insr. Chem. Eng. 1965, 11, 79. (124) Mittash, A. 2. Angew. Chem. 1928, 4 1 , 831.
+
12nCOc0,
Fe3.(..1,
Qz
n[Fes I C O I I Z I
Heat of Transformation:
t
12nC0,s,
I
QI
CASE 1: Formation of Bulk Fe Metal. E c a o d = + ~ 3 n E c . h e r i v er n c r l y o a n s i t ~ l / l 2 n But ~ ~ ~ ~ i t u ~kcal/mole 9 8 . 7 (Ref. 1 2 5 ) E c o o a = +3(98.7)/12 = +24.7 kcal/mole Ql = -33.4 + 24.7 = -8.8 kcallmole CASE 2: Formation of Fe Clusters. = +(0.5*2*3n'Ere-rs I/12n = + ~ E F ~ - F C / ~ = +2.5z QI (per COI = -33.4 t 2.52 for QI (per COl=-18 to -21 kcal/mole, z = 5 - 6 neighbors Econd
nearest
aEFE-Fc is the Fe-Fe bond strength in Fe,(CO),,; EFdo is the Feis the energy released during the CO bond strength in Fe,(CO),*; EWnd addition of an Fe atom to form an Fe particle; z is the number of nearest neighbors in an Fe particle. SCHEME 11: Thermodynamic Path To Estimate the Minimum Activation Energy of Decarbonylation of Fe(CO)S' nFel.)
+ I + 5nC01~,
Fen (
Qz
B o I
+ 5nC01
)
t
Qi
n[Fe l C 0 ) 1~
Heat o f Transformation: Q1 (tot) = Qz + Q 3 = I - 5 n E ~ ~ - c o tl I E ~ o r m a t l o o o f s o QI (per CO) = -15nEic-co1/5n t I E c o n a1/5n But l E ~ * - c o I= 28 kcal/mole (Ref. 121) QI (per CO) = -28.0 + I E c o o dI/5n
i < d l
CASE 1: Formation of Bulk Fe Metal. Econd = + I n E c + a e a ~ vE~n e T I Y o c n ~ i t ~ l / 5 n But E c ~ r o e~r I y ~o e o S ~ t t u$= 9 8~.kcal/mole 7 ~ (Ref. 1 2 5 ) E c o n d = tl98.7)/5 = + 1 9 . 7 kcal/mole Qt = -28.0 t 19.7 = -8.3 kcal/mole CASE 2: Formation of Fe clusters. Econd
= +10.5*2*n*Ere-rc1/5n = +zE~e-~e/lO
= +1.962 QI (per CO) = -28.0 + 1 . 9 6 2 for QI (per C0)=-15.5 kcal/mole, 2=6.3 nearest neighbors
(IEFtC0is the Fe-CO bond strength in Fe(CO)S;Eandis the energy released during the addition of an Fe atom to form an Fe particle; EPtFe is the Fe-Fe bond strength in Fe,(C0),2; z is the number of nearest neighbors in a Fe particle.
similar to those in bulk Fe for which the condensation energy is taken as the cohesive energy density'2s or the formation of small Fe particles with an Fe-Fe bond strength the same as that in the original carbonyl cluster. The first case provides a lower limit for the activation energy and gives 8.8 kcal/mol for F e 3 ( C 0 ) 1 2 and 8.3 kcal/mol for Fe(CO)S. The latter is in good agreement (125) Kittel, C. Introduction to Solid State Physics; Wiley: New 1976.
York,
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The Journal of Physical Chemistry, Vol. 93, No. 10, 19619
with a reported value of I O kcal/mol for precisely such a is QI = -33.4 + 2.52 In the second case, the heat of reaction, Q,, for Fe,(CO),, and Ql = -15.3 + 1.962for Fe(CO),, where z is the number of nearest-neighbor Fe atoms. If Ead is assumed equal to Q,in these endothermic reactions, z is about 6 in each case, which is a very reasonable number for the surface Fe atoms predominating in these very small Fe particles. This exercise further indicates that the lower E,,, values, compared to substitution reactions in solution, are consistent with the chemistry that is occurring. The reduced Fe/C catalyst was investigated under reaction conditions at 0.1 MPa and 573 K, but spectra were observed only after cooling to 300 K in C O and H,, as shown in Figure 7 . N o adsorbed C O was observed, probably due to the high H 2 / C 0 ratio of 68,and the only detectable bands were those associated with C H , and CH, groups, with the latter predominating. This indicates that even on Fe/C catalysts the hydrocarbon chains are relatively short when very high H 2 / C 0 ratios are utilized. It is very possible that these species are adsorbed on the carbon support rather than on the Fe surface, similar to the behavior observed for oxide-supported Ru c a i a l y ~ t s . ’ ~ ~ ~ ’ ~ * The adsorption of C O on the fully decarbonylated clusters was investigated after either L T R or H T R pretreatments. No IRactive species were detected after LTR a t either 195 or 300 K, while following H T R no IR bands associated with Fe were detected at 195 K and only Fe(CO)S was detected a t 300 K under C O . No IR bands were observed after purging. These results were reproducibly obtained on different samples and are distinctly different from those obtained with comparable R u / C and Os/C catalysts, in which a band for chemisorbed CO was readily observed.43“ The observation of Fe(CO), by DRIFTS is consistent with Mossbauer studies on this same catalyst which also showed the reversible formation of the pentacarbonyl cluster,2*62and it demonstrates that IR-active species can be detected when present. The capability to form Fe(CO)S has been associated only with very small Fe particles as it has not been reported on bulk Fe surfaces or large Fe crystallites.2,62 The chemisorption results indicate that large amounts of C O adsorb but are not observed by IR spectroscopy. Considering that C O does not dissociate at 195 K on Fe single crystals and polycrystalline films,129-144 the quantities of adsorbed C O are large enough to be detected a t 195 K if they exist in an IR-active form. A tilted C O species on the more open Fe( 100) and Fe( 1 1 1) planes has been proposed to account for the peaks at 1180-1245 and 1530 cm-I in EELS spectra.i39.145-1s1 The suggestion that these tilted species pre(126) King, F. T.; Lippincott, E. R. J . A m . Chem. Soc. 1956, 78, 4192. (127) Ekerdt, J. G.; Bell, A. T. J . Catal. 1980, 62, 19. (128) Ekerdt, J. G.; Bell, A. T. J . Card. 1979, 58, 170. (I29) Vannice, M. A . In Catalysis-Science and Technology; Anderson, J . R., Boudart, M., Eds.; Springer-Verlag: Berlin, 1982; Vol. 3, Chapter 3. (130) Bell, A . T. Catal. Rea. Sci. Eng. 1981, 3, 203. (131) Kock, A . J . H. M.; Geus, J . W . f r o g . Surf. Sci. 1985, 20, 165. (132) Jona, F.; Legg, K. 0.;Shih, H . D.: Jepsen, D. W.; Marcus. P. M. f h y s . Rec. Letr. 1978, 40, 1466. (133) Brundle, C. R. I B M J . Res. Dec. 1978, 22, 235. (134) Benziger, J.; Madix, R. J. Surf. Sci. 1980, 94, 119. (135) Vink. T. J.; Gijzeman, 0. L. J.; Geus, J. W. Surf. Sci. 1985, 150. 14. (I36) Rhodin, T. N.; Brucker, C. F. Solid Stare Commun. 1977, 23. 2 7 5 . (137) Brucker, C. F.; Rhodin, T. ?ISurf. . Sci. 1979, 86, 638. (I38) Textor, M.; Gay, 1. D.; Mason, R. froc. R . Soc. London 1977, A356, 37. (139) Seip, U.; Tsai, M. C.: Cristmann. K.; Kuppers, J.: Ertl. G. Surf. Sci. 1984, 139, 29. (140) Wedler, G.; Ertel, J . Ber. Bunsen-Ges. Phys. Chem. 1983, 87, 469. (141) Borden, G.; Gafner, G.; Bonzel, H. P. Appl. f h y s . 1977, 1 3 , 333. (142) Borden, G . ; Gafner, G.; Bonzel, H. P. Surf, Sri. 1979. 84, 295. (143) Yoshida, K.; Somorjai, G . A. Surf. Sci. 1978, 75, 46. (144) Gonzalez, L.; Miranda, R.; Ferrer, S . Surf. Sci. 1982, 119, 61. ( I 45) Seip, C.;Bassignana. I. C.: Kuppers, J.; Ertl, G. Surf Sci. 1985, 160, 400. (146) Moon, D. W.; Bernasek, S. L.; Lu, J. P.: Gland, J . L.; Dwyer, D. J . Surf. Sci. 1987, 184, 90. (147) ‘Moon. D. W.: Dwyer, D. J.; Bernasek, S. L. Surf S r i . 1985, 163, 215. (148) Moon, D. W.; Bernasek, S. L.: Dwyer, D. J.; Gland, J . L. J . A m Chem. Soc. 1985. 107. 4363.
Venter and Vannice dominate on very small Fe particles, which have very open, rough surfaces, is strongly supported by the recent work of Cameron and Dwyer in which XPS, LEED, TPD, and UPS results indicated that on an Fe(100) surface the first C O molecules to adsorb dissociate, and the next type of strongly adsorbed C O species is bound through both the oxygen and the carbon end of the molecule.’52 Although speculative, this explanation is currently the most appropriate to explain the absence of a spectrum for CO chemisorbed on these small Fe particles. The adsorption behavior of C O on Fe a t 300 K and above is known to be complex because CO dissociation can occur.129J30 The adsorption of C O is dependent upon both structure and t e m p e r a t ~ r e , with ’ ~ ~ dissociation occurring readily a t 300 K on Fe( 100)‘32-137 and Fe(l11) surfaces138*139 as well as polycrystalline but it only occurs slowly at 320 K on the close-packed However, in these studies no C O dissociation Fe( 110) has been observed on these Fe surfaces a t or below 195 K. In addition to the problem of CO dissociation, various bonding configurations on Fe are possible, and adsorption states such as “on top” ( C 0 : F e = l : l ) , “shallow hollow” ( C 0 : F e = l:2),and “deep hollow” ( C 0 : F e = 1:3 and 1:4) have been specified on Fe( 1 1 1 ) surface^.'^^^'^^ Regardless, a reasonable stoichiometry at 195 K appears to be CO:Fe, = 1:2 because it has given estimates of particle size for large Fe crystallites in good agreement with other technique^,^^^^ and it has also provided consistent results on highly dispersed Fe/MgO catalysts.s8 Despite the uncertainty, the C O adsorption measurements illustrate that highly dispersed Fe/C catalysts were obtained after either LTR or H T R steps. The lower hydrogen uptakes also imply very small Fe particles, and when the correlation of the H2ad/CO(I95 K) ratio with Fe particle size reported by Topsoe et aL5*is used, it indicates particle sizes of 1-2 nm in our samples, which is very consistent with the sizes determined by C O adsorption at 195 K. The adsorption measurements also provide important additional information. First, the dispersion of the 8.7% Fe/C catalyst was slightly lower than that of the 6.4% Fe/C catalyst, which is reasonable due to the higher loading. Second, from the IR spectra at 195 K it is clear that no detectable amount of Fe(CO)S formed at that temperature, but heating the sample to 300 K under C O produced Fe(CO),. This could lead to concern that adsorption values at 300 K may not provide reliable estimates of Fe dispersion, but the irreversible C O adsorption at 300 K represents the residual uptake following a 1 -h evacuation between isotherms. DRIFTS showed that Fe(CO), was present a t 300 K in the presence of gas-phase CO but that its concentration diminished markedly after purging for 1 h with He or H,. Therefore, the irreversible adsorption a t 300 K represents C O adsorbed in forms other than Fe(CO),, and these concentrations of C O are easily high enough to be detectable by DRIFTS if they are present in IR-active forms.42-4s The integral heats of adsorption for C O on the 8.7 and 6.4% Fe/C catalysts, measured isothermally a t 300 K with a differential scanning calorimeter, are summarized in Table I1 and compared to reported literature values in Table V.1s3-’60 The average heat of adsorption of C O on these small Fe particles was 15.0 & 1.6 ~
~
~~~~
(149) Moon, D. W.; Cameron, S.; Zaera, F.; Eberhardt, W.; Carr, R.; Bernasek, J. L.; Gland, J. L.: Dwyer, D. J. Surf. Sci. 1987, 180, L123. (150) Kaesz, H . D.: Sailant, R. B. Chem. Rea. 1972, 72, 231. (151) Benndorf, C.; Kruger, B.: Thieme, F. Surf. Sci. 1985, 163, L675. (152) Cameron, S. D.; Dwyer, D. J. Langmuir 1988, 4, 282. (153) Zakumbaeva, G . D.; Beketaeva, L. A,; Uvaliev, T. Yu.; Khylstov, A. S.; Kuanyshev, A . Sh.; Sagov, Yu. M. React. Kinet. Catal. Lett. 1987, 34, 213.
(154) Dry, M. E.; Shingles, T.; Boshoff, L. J.; Oosthuizen, G. J. J . C a r d . 1969, 15, 190. (155) Toyoshima, I.; Somorjai, G . A. C a r d . Reo. Sci. Eng. 1979, 19, 105. (156) Kolbel, H.; Haubold, H. Ber. Bunsen-Ges. f h y s . Chem. 1961, 65,
421. (I5 7 ) Kolbel, H.; Roberg, H. Ber. Bunsen-Ges. fhys. Chem. 1977, 81, 634. (158) Walker, G.; Colb. K. G . ; McElhney, G.; Heinrich, W. Appl. Surf. S r i . 1978, 2. 30.
(159) Gijzeman, 0. L . J.; V i n k , T. J.; van Pruisen, 0. P.; Geus, J W. J . Vac. Sci. Technol. 1987. 5, 718. (160) Zakumbaeva. G . D.; Dostiyarov, A . M.; Nnidin, V . A,; Vozdvirhenskii. V F. Kinet. Catal. 1986, 27. 307.
J. Phys. Chem. 1989, 93, 4167-4173 kcal/mol CO, which is somewhat lower than most of the reported integral Qadvalues on bulk Fe or large Fe crystallites, as most values have ranged between 17 and 25 kcal/mol although values as low as 6 kcal/mol have been reported for bulk Fe304-derived catalysts. In the comprehensive study of Fe catalysts by Topsoe et al., the influence of particle size on the differential heats of adsorption was investigated, and it was proposed that two major types of sites exist which possess initial Qadvalues near 30 and 19 kcal/mol, with the latter type predominating on very small Fe particles.58 Our average value of 15.0 kcal/mol, which is close to the integral value of 17 kcal/mol calculated from the data of Topsoe et al. for 3.8-nm Fe particles on MgO, is consistent with this proposal. When compared to the work of Moon et al., who reported three adsorption states on Fe( 100) giving heats of desorption of 26.6 kcal/mol for highly tilted CO, 18.0 kcal/mol for bridged CO, and 12.8 kcal/mol for more weakly held C0,’46149 the value obtained from our experiments indicates a predominance of the last two species a t high coverages. Finally, the study of Geus and coworkers showed that Qadvalues for CO on 0-covered or C-covered Fe surfaces are lower than on clean surfaces,lsg and these values between 11 and 21 kcal/mol are similar to our value. One implication is that C O dissociation may be facilitated on these small Fe crystallites, thereby giving lower Qadvalues.
Summary DRIFTS has been applied to study carbon-supported Fe cluster catalysts for the first time and to measure decarbonylation rates of Fe3(CO),, and Fe(CO)5 on a clean carbon surface. Deposition of Fe3(CO) on this high-surface-area carbon yielded substantial amounts of Fe(CO)5, even a t temperatures below 300 K; however, this provided the opportunity to study the decomposition of both species. Under either flowing H e or H2, the same activation energy of 15.5 f 0.2 kcal/mol was determined for Fe(CO)5 while values of 17.9 and 2 1.3 kcal/mol, respectively, were obtained for Fe3(CO),,. This appears to be the first kinetic investigation of the
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thermal decomposition of Fe3(C0),, in the absence of other substituting ligands. The first-order decarbonylation rate constants for the dodecacarbonyl cluster are very similar to those for substitution reactions in solution, although the activation energies are somewhat lower. Comparable studies of Fe pentacarbonyl in solution could not be found. A similar rate-determining step for the thermal decomposition of the cluster on carbon and for substitution reactions in solution is suggested; Le., the loss of the first CO ligand. The formation of small metallic Fe crystallites, which does not occur in solution, accounts for the lower activation energies. The same decarbonylation process occurs in either H e or H,, and the formation of hydrido-iron carbonyl clusters was not observed, contrary to the behavior of Os and Ru dodecacarbonyl on this same carbon. I R spectra a t 300 K after cooling from reaction conditions showed only CH, and CH2 frequencies, and the more intense CH, bands indicated the presence of only short-chain hydrocarbons on the catalyst surface, very possibly on the carbon. Calorimetric measurements gave an integral CO heat of adsorption of 15.0 A 1.6 kcal/mol a t 300 K, which is consistent with values reported for small Fe crystallites. The high co uptakes, the low Had/COad ratios, the low Qadvalue, the absence of IR bands for adsorbed CO, the formation of Fe(CO)5 under CO, and the low decarbonylation activation energies are all consistent with the conclusion that decomposition of Fe3(C0),2 (and Fe(C0)5) produces only small Fe particles (C4 nm) on this carbon. This conclusion is also supported by additional MES, S T E M , and kinetic studies which are reported e l s e ~ h e r e . ~ , ~ ~ f ’ ~
Acknowledgment. This research was sponsored by the N S F Kinetics and Catalysis Program through Grant CBT-86 196 19 and the donors of the Petroleum Research Fund, administered by the American Chemical Society. Registry No. Fe, 7439-89-6; C, 7440-44-0; Fe,(CO),,, 17685-52-8; Fe(CO)S, 13463-40-6; CO, 630-08-0; H,, 1333-74-0.
Critical Concentrations and Compositions of Mixed Micelles of Sodium Dodecylbenzenesulfonate, Tetradecyltrimethylammonium Bromide, and Polyoxyethylene Octy Iphenols A. Graciaa, M. Ben Ghoulam, G. Marion, and J. Lachaise* Laboratoire de Thermodynamique des Etats MZtastables et de Physique MolZculaire, Centre Universitaire de Recherche Scientifique, Avenue de I’UniversitZ, 64000 Pau. France (Received: July 25, 1988; In Final Form: December 28, 1988)
The model for multicomponent nonideal mixed micelles used via the regular solution approximation allows accounting for the cmc of binary or ternary mixtures obtained with SDBS (anionic surfactant), TTAB (cationic surfactant), and OP(EO), (nonionic surfactants). In the binary mixed micelles, the anionic/nonionic interactions are slightly higher than the cationic/nonionic interactions, although of the same order of magnitude. This weak difference could be due to the formation of an oxonium salt. The anionic/cationic interactions are much higher and are probably due to the high electrostatic attraction between molecules whose polar heads have opposite charges. The introduction in the model of the pair interaction parameters, determined from the cmc of the binary mixtures, gives predictions for the cmc and for the micelle compositions of the ternary mixtures. Thus, synergism in the formation of the ternary mixed micelles would exist almost in the whole range of the compositions, except for mixtures quasi-exclusively composed of one of the ionic surfactants. Whatever the quantity of nonionic surfactant in the ternary mixture, the maximum of synergism would be obtained for equimolecular composition in anionic surfactant and cationic surfactant. Furthermore, in the large zone of synergism, the ternary mixed micelles would always be composed of an equal number of anionic molecules and cationic molecules; so they would be practically neutral. Measurements of the cmc of ternary mixtures, in agreement with these predictions, verify their validity.
Introduction Micelles composed of mixtures of surfactants with different structures (mixed micelles) are of great theoretical and industrial interest. Since surfactants used in practical applications are rarely pure, there is increasing interest in understanding the structure 0022-3654/89/2093-4l67$01.50/0
and properties of mixed micelles. Examples of such applications are detergency and e d ~ a n c e dOil recovery. Models for binary mixtures have been developed on the basis of pseudo phase separations. Generally they assume ideal mixing of the surfactants in the micelle for nonionic or ionic surfactants,
0 1989 American Chemical Society
