© Copyright 1998 by the American Chemical Society
VOLUME 102, NUMBER 38, SEPTEMBER 17, 1998
LETTERS Decomposition of Dimethyl Methylphosphonate (DMMP) on Alumina-Supported Iron Oxide Teweldemedhin M. Tesfai, V. N. Sheinker, and Mark B. Mitchell* Department of Chemistry, Clark Atlanta UniVersity, Atlanta, Georgia 30314 ReceiVed: January 6, 1998; In Final Form: July 22, 1998
The adsorption of dimethyl methylphosphonate (DMMP) on alumina-supported iron oxide has been examined. DMMP reacts on adsorption at room temperature, apparently through cleavage of the phosphorus-carbon bond. This bond is observed to be extremely resistant to cleavage when DMMP is adsorbed on oxides such as alumina, magnesia, and lanthana. The phosphorus-methoxy bonds, which are the most readily cleaved on the other oxides, appear at least initially to remain intact on the alumina-supported iron oxide. The hypothesis proposed to account for the unusual activity of the iron oxide surface is an oxidation pathway involving the Fe(II)/Fe(III) redox couple.
Introduction An understanding of the interaction of phosphonate esters with surfaces is critical for the development of catalysts for the demilitarization of chemical warfare agents or the decontamination of surfaces exposed to those agents.1 This same information is important for the development of models of the migration of pesticides through soils. A number of groups have been involved in these studies, including Weinberg et al.,2,3 Klabunde et al.,4-8 and White et al.9,10 Recently, our group has had good success studying the adsorption and decomposition of dimethyl methylphosphonate (DMMP) on a variety of metal oxide surfaces, including aluminum oxide, magnesium oxide, lanthanum oxide, and iron oxide.11 DMMP, illustrated below, is a widely used simulant for Sarin and other toxic organophosphonate esters. Infrared diffuse reflectance spectroscopy was used
in our earlier study to examine the successive elimination of the carbon-containing fragments as a function of temperature.
On aluminum oxide, magnesium oxide, and lanthanum oxide, the decomposition of DMMP was found to proceed in a clear sequence, involving the elimination of one methoxy group at approximately 100 °C, followed by elimination of the second methoxy group at approximately 300 °C. The phosphoruscarbon bond was observed to remain intact even at 400 °C in the presence of oxygen when DMMP was adsorbed on aluminum oxide. The initial steps apparently involve an acidcatalyzed nucleophilic displacement reaction, Scheme 1, with surface oxide groups replacing one of the methoxy groups. In this mechanism, however, the surface is not acting as a true catalyst since it is also the nucleophile that replaces the CH3Ogroup, and the active site is not recycled. This mechanism accounts for the resilience of the P-CH3 bond, in that the CH3group is a very poor leaving group, especially compared to the CH3O- group. Decomposition on iron oxide was observed to proceed in a different fashion than on the other oxides, with no obvious preference for eliminating the methoxy groups or the methyl group attached to the phosphorus. The numbers of both types of methyl groups were observed to decrease as the temperature was increased, with virtually complete elimination of all the carbon-containing fragments by 300 °C. We postulated that the mechanism by which the iron oxide surface cleaves the
S1089-5647(98)00690-7 CCC: $15.00 © 1998 American Chemical Society Published on Web 08/27/1998
7300 J. Phys. Chem. B, Vol. 102, No. 38, 1998 SCHEME 1
P-CH3 bond is through an oxidative mechanism involving the Fe(II)/Fe(III) redox couple. The Mars and Van Krevelen mechanism12 is an oxidation mechanism whereby an oxide ion on an oxide surface is made available via a change in the oxidation state of the metal. This ion can react with adsorbed DMMP to cleave the P-CH3 bond as shown below. Oxidation of the surface by gas-phase oxygen can regenerate the oxide sites on the surface.
Fe2O3 + 2e- f 2FeO + O2(H3CO)2P(O)-CH3 + O2- f (H3CO)2P(O)ads + H3COads
2FeO + 1/2O2 f Fe2O3 The work described here represents the initial results of a study to clarify the activity of iron oxide by examining a supported iron oxide material. The support chosen was aluminum oxide because the chemistry of the aluminum oxide surface is understood and reactivity due to the presence of iron oxide could easily be discriminated from activity due to the alumina support. It was envisioned that at least two iron atoms in close proximity would be needed to accomplish the proposed reaction sequence of the Mars and Van Krevelen mechanism, but that higher activity would be realized with networks of iron oxide on the surface. The networks would sustain and delocalize the charge depletion that results from the proposed reaction sequence. Experimental Section Catalyst Synthesis and Characterization. The catalysts were prepared using a successive impregnation protocol using a nonaqueous impregnation method and an Fe(III) complex. The complex was dissolved in acetonitrile, and a fixed amount of γ-alumina was added. The mixture was stirred at room temperature for 24 h and the solid filtered from the solution. The solid was then calcined at 500 °C in air. An initial determination of the iron loading was accomplished by measuring the visible absorption spectrum of the impregnation solution before and after impregnation and determining the amount of complex adsorbed by difference. The impregnation process was repeated as needed to accomplish the desired loading. The final determination of the iron loading on the samples was accomplished by Galbraith Laboratories, Knoxville, TN. The surface areas of the samples were determined from multipoint BET surface area analyses using a Micromeritics Gemini 2360 Surface area analyzer.
Letters Infrared Diffuse Reflectance Studies. Before study, the alumina and the alumina-supported iron oxide were calcined at 500 °C in air in a muffle furnace for 24 h and were stored in sealed containers in a desiccator prior to use. The pure solid being examined was sieved and placed in a Harrick Scientific HVC-DR infrared diffuse reflectance controlled environment cell, using a Harrick Scientific DRA-2 optical accessory. The cell was closed and evacuated using an Alcatel diffusion pump and heated to 400 °C under vacuum to remove any adsorbed species. After the cell was cooled to room temperature under vacuum, a dilute mixture of DMMP in helium (1:1000) was allowed to flow through the cell. The pressure in the cell during this time was approximately 500 mTorr. Spectra were measured during the adsorption process to monitor for DMMP adsorption using a Nicolet Magna 750 FT-IR and a liquid nitrogen cooled mercury-cadmium-telluride (MCT) detector. During adsorption, 100 scans at 8 cm-1 resolution (to minimize collection time) were coadded to form the spectra shown. Results The results that are presented in this Letter concern two different materials: (1) pure γ-alumina and (2) an aluminasupported iron oxide prepared by the successive impregnation of γ-alumina with iron acetylacetonate seven times to yield a solid that measured 11.6 wt % Fe. On the basis of the number of OH sites on the (111) surface of alumina, which corresponds to the number of aluminum sites on the surface,13 the aluminasupported iron oxide has approximately 80% of its surface covered with iron oxide, provided no three-dimensional crystals are formed. We are not reporting the adsorption results from solids with intermediate loadings of iron oxide. It was thought that the solid with the highest iron concentration was the most likely to show behavior significantly different from that observed for pure alumina. Further characterization of the intermediateloading solids is needed, and those results will be presented in a future publication. The initial surface area of the alumina substrate was measured to be 148 m2/g, while that of the iron oxide material was determined to be 106 m2/g. This loss in surface area is greater than what would be expected if it is assumed that the iron oxide adds only mass and no surface area, and indicates that there probably is some blocking of small pores. The infrared absorption spectra of DMMP and several of its 2H- and 13C-substituted derivatives have been assigned by van der Veken and Herman14 and by Moravie et al.,15 making the infrared spectra of adsorbed DMMP relatively easy to interpret. In previous work from our laboratory,11 it has been shown that the methyl stretching modes of adsorbed DMMP are excellent indicators of the progress of adsorption and decomposition. The methyl groups do not participate directly in interactions with the surface, and their absorption frequencies are not significantly modified on adsorption. The relative intensities of the different methyl group vibrational absorptions can be used very effectively to determine the extent of decomposition of DMMP on surfaces. In this work, we have not included spectra of the lower frequency region, 1500-900 cm-1, which show the methyl deformation mode, δ(CH3P), at 1317 cm-1. However, the absorptions in that frequency region and the behavior of the deformation mode are entirely consistent with the results observed for the methyl (CH3P) stretch absorptions that are shown. Figure 1A,B shows a sequence of spectra measured as the DMMP/He mixture was flowing through the cell over pure aluminum oxide. Figure 1A shows the spectra acquired as a
Letters
Figure 1. Spectra of DMMP on γ-alumina acquired as a function of exposure to flowing mixture of DMMP in He. The spectra are labeled according to their total exposure to DMMP. In (A) the spectra are shown with the same absolute intensity scale. In (B) the spectra are overlaid and normalized to the intensity of the band at approximately 2960 cm-1.
function of exposure to a flowing mixture of DMMP in He at room temperature. The spectra are labeled according to the amount of DMMP that has passed through the sample. In Figure 1B, the spectra are shown overlaid, with the intensity of each spectrum normalized to the intensity of the most intense band, the infrared absorption at 2960 cm-1. Figure 1B clearly shows that all of the bands due to the methyl stretching modes, the pair due to the two methoxy (CH3O) groups (at 2960 and 2850 cm-1) and the pair due to the methyl (CH3P) group (at 2990 and 2930 cm-1) increase with exposure to DMMP with the same relative intensity. These results show that DMMP is adsorbing molecularly at room temperature, without loss of the methoxy or the methyl groups. This result was expected on the basis of our earlier work and the work of others.2-11 Figure 2A,B shows a sequence of spectra taken as the DMMP/He mixture was flowing through the cell over the supported iron oxide. The format for the figure is the same as was discussed for Figure 1. In Figure 2A, it is clear that the bands due to the methyl (CH3P) group are not present during the initial exposure of the sample to DMMP. In Figure 2B, the methyl (CH3P) bands can be seen to grow in, relative to the methoxy absorptions, after exposure to larger amounts of DMMP. If DMMP were adsorbing molecularly, the relative intensities of all of the bands would be constant, regardless of the total amount of DMMP adsorbed. This is not the case in the spectra of Figure 2A,B, which indicates that the molecule is adsorbing dissociatively, presumably by cleavage of the phosphorus-carbon bond. The lack of absorption by the methyl (CH3P) modes during initial stages of DMMP adsorption on alumina-supported iron oxide is shown dramatically in Figure 3. In this figure we have shown spectra of the two solids after exposure to comparable
J. Phys. Chem. B, Vol. 102, No. 38, 1998 7301
Figure 2. Spectra of DMMP on alumina-supported iron oxide (11.6 wt % Fe), acquired as a function of exposure to a flowing mixture of DMMP in He. The spectra are labeled according to their total exposure to DMMP. In (A) the spectra are shown with the same absolute intensity scale. In (B) the spectra are overlaid and normalized to the intensity of the band at approximately 2960 cm-1.
Figure 3. Spectra representing exposure of the two solids to comparable amounts of DMMP are shown overlaid, with a common intensity scale.
amounts of DMMP. The same vertical scale is used for each spectrum. The spectrum of alumina exposed to 0.55 µmol of DMMP clearly shows the presence of the two methyl (CH3P) stretching modes. These modes are absent in the spectrum of the alumina-supported iron oxide exposed to 0.59 µmol of DMMP. Also notable from this figure is that the methoxy (CH3O) stretching modes for DMMP on the two solids have similar intensities after exposure to comparable amounts of DMMP. In this work, the only evidence presented for the decomposition route of DMMP on the iron oxide surface is that from infrared spectroscopy. However, our group is not the only one to suggest that cleavage of the phosphorus-carbon bond occurs when DMMP is adsorbed on iron oxide surfaces. Henderson
7302 J. Phys. Chem. B, Vol. 102, No. 38, 1998
Letters exposed to approximately 10 µmol of DMMP, and as a result infrared absorptions due to molecular DMMP can clearly be observed. The spectra in Figure 4 are very similar to the spectra shown in Figure 1 (DMMP on alumina), showing the good correspondence between the molecular DMMP infrared absorptions on the two surfaces. Conclusions DMMP adsorbs onto the alumina-supported iron oxide surface with cleavage of the phosphorus-carbon bond at room temperature. After adsorption of significant amounts of DMMP, presumably corresponding to saturation of the surface sites and passivation of the material, molecular DMMP can be adsorbed.
Figure 4. Spectrum of DMMP on alumina-supported iron oxide after adsorption of multiple layers of DMMP.
et al., in 1986,9 suggested that phosphorus-carbon cleavage of DMMP occurs on iron oxide at low temperature. They used Auger electron spectroscopy and temperature-programmed desorption to examine R-Fe2O3 dosed with DMMP and found that the adsorbed phosphorus species was fully oxygencoordinated, consistent with a phosphate or phosphite species, and inconsistent with the presence of a methyl group bound to the phosphorus. They also observed no molecular DMMP desorption peak at low exposures. A possibility that would account for the loss of the CH3P methyl stretch absorption is cleavage of one of the carbon-hydrogen bonds during DMMP adsorption. This might account for a shift in the methyl stretching modes but would not be consistent with the results of Henderson et al. The material examined by Henderson et al. was a single-phase, low surface area material, while the material examined in this study is probably amorphous, since no attempt was made to convert it to a single phase, and has a high surface area. While the two materials are physically distinct, the evidence indicates that they are acting in chemically similar ways. We conclude that DMMP is adsorbed on the supported iron oxide material with cleavage of the P-C bond, yielding an adsorbed fragment similar to the one proposed in the mechanism above, (CH3O)2PO-, and an adsorbed methoxy group. Eventually, the surface sites are saturated and molecular DMMP is adsorbed on the supported iron oxide. This is shown in Figure 4. In that experiment, the supported iron oxide was
Acknowledgment. The authors acknowledge the help of Mr. Gautam Saha, who carried out the surface area measurements. The authors also thank Professor Mark G. White for many helpful discussions and much insight. We gratefully acknowledge the financial support of this work by the Army Research Office through Grant DAAL03-90-G-0208, the support of ERDEC through Grant DAAA15-94-K-0004, and NASA through Grant NAGW-2939. References and Notes (1) Ekerdt, J. G.; Klabunde, K. J.; Shapley, J. R.; White, J. M.; Yates, J. T., Jr. J. Phys. Chem. 1988, 92, 6182. (2) Templeton, M. K.; Weinberg, W. H. J. Am. Chem. Soc. 1985, 107, 97. (3) Templeton, M. K.; Weinberg, W. H. J. Am. Chem. Soc. 1985, 107, 774. (4) Li, Y. X.; Schlup, J. R.; Klabunde, K. J. Langmuir 1991, 7, 1394. (5) Lin, S. T.; Klabunde, K. J. Langmuir 1985, 1, 600. (6) Atteya, M.; Klabunde, K. J. Chem. Mater. 1991, 3, 182. (7) Li, Y. X.; Klabunde, K. J. Langmuir 1991, 7, 1388. (8) Li, Y. X.; Koper, O.; Atteya, M.; Klabunde, K. J. Chem. Mater. 1992, 4, 323. (9) Henderson, M. A.; Jin, T.; White, J. M. J. Phys. Chem. 1986, 90, 4607. (10) Hegde, R. I.; White, J. M. J. Phys. Chem. 1986, 90, 2159. (11) Mitchell, M. B.; Sheinker, V. N.; Mintz, E. A. J. Phys. Chem. B 1997, 101, 11192. (12) Mars, P.; Van Krevelen, D. W. Chem. Eng. Sci. Suppl. 1954, 3, 41. (13) Kno¨zinger, H.; Ratnasamy, P. Catal. ReV.- Sci. Eng. 1978, 17, 31. (14) Van Der Veken, B. J.; Herman, M. A. Phosphorous Sulfur 1981, 10, 357. (15) Moravie, R. M.; Froment, F.; Corset, J. Spectrochim. Acta 1989, 45A, 1015.
