J. Phys. Chem. C 2007, 111, 9417-9426
9417
Room Temperature Reaction of Ozone and Dimethyl Methylphosphonate (DMMP) on Alumina-Supported Iron Oxide Mark B. Mitchell,* Viktor N. Sheinker, and Woodrow W. Cox, Jr. Department of Chemistry and the Center for Surface Chemistry, Clark Atlanta UniVersity, Atlanta, Georgia 30314 ReceiVed: October 4, 2006; In Final Form: April 10, 2007
The reaction of dimethyl methylphosphonate (DMMP) with ozone on a reactive adsorbent, alumina-supported iron oxide, has been examined at room temperature. The surface oxidation reaction was found to follow the Langmuir-Hinshelwood mechanism, involving a surface adsorbed and activated DMMP molecule or fragment at a Lewis acid site and an active oxygen species on the surface formed from ozone. Carbon dioxide is the primary reaction product observed, and the amount formed was found to plateau as a function of ozone concentration after a concentration of approximately 300-400 ppm. The surface reaction with ozone on the 10 wt % Fe2O3/Al2O3 adsorbent generates approximately 2.7 times as much gas-phase carbon, in the form of CO2 and CO, as does the same surface in the absence of ozone, in the form of methanol. The ozone appears to increase the number of active sites for DMMP decomposition by liberating some of the reaction products from the surface and recycling the surface sites. The supported iron oxide material is a more effective reactive surface than alumina for the reaction with ozone primarily because of its greater facility for ozone decomposition. The gas-phase reaction between DMMP and O3 for the formation of CO2, while not the focus of the study, was found to be of fractional order in O3, indicating a multistep reaction mechanism.
Introduction Dimethyl methylphosphonate (DMMP) is a widely used simulant for chemical warfare agents and pesticides. Several recent publications from our laboratory have discussed examinations of the surface reactions and decomposition efficiencies of a variety of different metal oxides for the adsorption and reaction of DMMP with metal oxide surfaces at ambient and elevated temperatures.1-6 All of those reports, and studies from a variety of other labs as well,7-10 have shown that the decomposition reactions of DMMP on oxide surfaces at these temperatures (25-400 °C) are stoichiometric reactions, not catalytic ones, and though the reactive adsorbents may be very active initially, their total capacity for decomposition is limited to some fraction, less than one, of the number of surface adsorption sites. There is not enough energy available at these relatively low temperatures, in the absence of an energetic coreactant or an additional source of energy, to accomplish turnover of the active sites once decomposition has occurred. Methods of maximizing the DMMP decomposition yield per unit weight of solid reactant include using very high surface area substrates for dispersed nanodimensional oxidative reactants as has been carried out in our laboratory or using nanodimensional reactants as developed by Klabunde and co-workers.7,11 There have been at least two studies that have shown evidence of pathways that circumvent the surface site limitation. For the DMMP decomposition reaction at 500 °C, Jiang et al. observed what appeared to be a unique transport property involving an iron oxide-coated MgO nanoparticle assembly.12 The core-shell particles showed a much higher capacity for destruction of DMMP than did the normal MgO nanoparticles or Fe2O3 itself. It was postulated that the Fe2O3 shell was able to cause the * To whom correspondence should be addressed. E-mail: mmitchell@ cau.edu.
active MgO oxidant to be more mobile and facilitate the participation of virtually all of the MgO present in the core of the particle in the oxidative decomposition of DMMP. Although the reaction is still stoichiometric, the decomposition capacity is twice that observed for nanodimensional MgO alone since interior magnesium species participate in the reaction. Wagner and co-workers observed a similar effect when using alumina nanoparticles in aqueous solution for the decomposition of VX, GB, GD, and HD.13 They found that, in the presence of large amounts of agent, virtually all of the aluminum in nanosized aluminum oxide particles can react. However, in those studies, the reaction of the aluminum atoms in the particle interior occurs as the phosphonic acid reaction products erode the outer layers of the particle, exposing the inner aluminum species for reaction. This process takes several days to occur, two weeks for GB and GD, and seven weeks for VX. For lower amounts of agent, the total decomposition capacity was still limited by available surface reaction sites. The fact that the DMMP decomposition capacity of the reactive adsorbents at low temperatures is limited by surface stoichiometry has led us to examine the use of a co-reactant to provide additional chemical energy and increase the decomposition yield by either (a) making it possible recycle (turnover) the active decomposition sites or (b) increasing the number of surface sites that are able to facilitate the decomposition. One promising co-reactant is ozone. Ozone is a powerful oxidizing agent that is easily generated with a source of electricity and dry air, which makes its use as a co-reactant at remote locations straightforward. Additionally, excess ozone can be easily and efficiently decomposed using a metal oxide catalyst at ambient temperatures, so that its use need have no negative environmental consequences. This manuscript reports the results of our studies of the reaction of DMMP with ozone on reactive adsorbents using a
10.1021/jp066533x CCC: $37.00 © 2007 American Chemical Society Published on Web 06/13/2007
9418 J. Phys. Chem. C, Vol. 111, No. 26, 2007
Figure 1. Experimental setup for determining the decomposition of DMMP on oxides in the presence of O3. This is a modification of an earlier setup for examining the decomposition of DMMP on oxides alone (see refs 3 and 6).
variety of adsorbents including silica, alumina, alumina-supported cerium oxide, and alumina-supported iron oxide, with the focus on the iron oxide materials. The reason for using alumina-supported cerium oxide was earlier success found with this adsorbent in the absence of an added oxidant.4-6 Similarly, the reason for the examination of supported iron oxide was earlier experience with supported iron oxide that showed interesting surface reaction chemistry with DMMP.2 Additionally, iron oxide and the oxidant hydrogen peroxide are wellknown in aqueous systems to form an efficient reactant combination, in the form of Fenton’s reagent, for the oxidation of organic contaminants.14
Mitchell et al. Figure 1. The setup is a modification of the one used to examine the decomposition in the absence of an added co-reactant, and a detailed description of the original apparatus and typical reaction conditions and procedures can be found in earlier references.3,6 Approximately 60 mg of the sample was placed in the u-tube, and the sample was preheated at 350 °C in O2 for 1 h followed by pure helium for 30 min. The sample was then cooled to 25 °C in helium flowing at a rate of 30 standard cm3/min (30 sccm). The pure He flow was turned off, and the flow of the DMMP/He mixture to the u-tube was established at 30 sccm and was mixed with an appropriate amount of the O3/O2 mixture to develop the concentrations needed. The total flow was typically 36 sccm. The DMMP bubbler and aerosol trap are kept in a thermostatted enclosure to ensure delivery of the calibrated amount of DMMP (47.0 µmol/L). After the microreactor, the gas stream flowed to a long-path gas cell (Ultra-Mini Long Path Cell from Infrared Analysis, Inc.) with an effective path length of 2.4 m and an internal volume of 100 mL. Infrared spectra (Thermo-Nicolet Avatar 360 FT-IR), at 2 cm-1 resolution, were collected approximately every 2 min. The data analysis and product quantitation were carried out as in the earlier study. The ozone generator used is a model Lab 11 from Pacific Ozone Technologies and the ozone monitor is a model 450H from Advanced Pollution Instruments (API) Inc. The oxygen used to generate the ozone was 99.994% minimum purity UHP O2 from Airgas South, containing less than 2 ppm H2O and less than 1 ppm CO2. The flow from the generator is adjusted using a splitter to send the appropriate amount to the reactor, with the remainder of the flow directed to an ozone destruction catalyst. The surface area determinations were carried out using a Micromeritics Gemini model 236 surface area analyzer. The X-ray diffractometer used to analyze the materials is a Phillips X’Pert model RW-3040 using a Cu KR source at 1.54 Å at a voltage of 45 kV and 40 mA current.
Experimental Methods The helium used for the study was a 99.9995% ultrahigh purity grade from AirGas South. DMMP was purchased from Aldrich and distilled under vacuum before use. The reactive adsorbent samples were prepared using a commercially available Bayerite material from UOP called Versal B as the substrate, which is transformed into η-Al2O3 on calcining. The η-Al2O3 has a final surface area of approximately 360 m2/g. The impregnated η-Al2O3 samples were prepared using precipitation deposition, which, in the case of cerium oxide, results in the formation of small (d < 6 nm) domains on the alumina surface.6,15,16 The precursor for the cerium oxide materials was Ce(NO3)3‚6H2O from Aldrich and that for the iron oxide materials was Fe(NO3)3‚9H2O from Aldrich. For the preparation of the impregnated solid reactants, three grams of the η-Al2O3 were added to 500 mL of water while stirring vigorously. Enough of the nitrate precursor was then added to the solution to make the desired weight loading of the corresponding oxide on the substrate. With continuous stirring, a 0.1 N solution of NH4OH was added dropwise to the suspension to obtain a final pH of approximately 9.5. After an additional 4-5 h of stirring, the solution was left to sit overnight. Most of the liquid was then decanted and the solid filtered from the suspension, dried at 80 °C overnight, and then calcined overnight at 400 °C (with a heating ramp of 10 °C/min). The powder was then sieved, and the particles in the range of 90250 µm diameter were used for testing. The experimental setup used to determine the decomposition yield of DMMP in the presence of added ozone is shown in
Results Gas-Phase Reaction. In order to examine the formation of CO2 by metal oxide surface reactions between DMMP and ozone, it is important that the CO2 formed by other paths, including the gas-phase reaction between ozone and DMMP, be accounted for. The formation of CO2 in the absence of a metal oxide reactive adsorbent was examined as a function of ozone concentration using the setup shown in Figure 1, but without the oxide adsorbent. In these experiments, there is a period prior to stabilization of the formation rate during which the CO2 concentration builds, and which is attributed to DMMP adsorption on the interior surfaces of the microreactor and passivation of the interior surfaces by ozone. As a check, the empty reactor was replaced with a length of Teflon tubing for the 500 ppm O3 experiment, yielding the same CO2 formation rate as that found for the empty reactor experiment after the production had stabilized, within experimental error, but without any noticeable induction period. Thus, the CO2 formation observed in the absence of a reactive adsorbent is due to the gas-phase reaction between ozone and DMMP, with no significant contribution from any surface reaction on the interior walls of the reactor. Figure 2 shows the equilibrated formation rate of CO2 as a function of O3 concentration in the absence of a solid reactive adsorbent. This rate shows a monotonic increase with ozone concentration, but not a linear one, and has an apparent reaction order that is fractional in O3 (see Figure 2), consistent with a multistep reaction mechanism. It is useful to compare the results
Room Temperature Reaction of Ozone and DMMP
Figure 2. CO2 formation rate observed from the decomposition of DMMP in the absence of a reactive adsorbent. The graph shows the equilibrated formation rate for CO2 observed as a function of the O3 concentration, with a fixed DMMP concentration and fixed total flow rate. Two sets of units are provided for each of the axes to allow for comparison to the other data.
J. Phys. Chem. C, Vol. 111, No. 26, 2007 9419
Figure 4. Cumulative CO2 production as a function of different metal oxide and oxide loading after exposure to a total of 650 µmol DMMP in the presence of 500 ppm O3.
TABLE 1: Products Formed from DMMP Decomposition on 10 wt % Fe2O3/Al2O3, After Exposure to 650 µmol DMMP
Figure 3. X-ray diffraction patterns observed for the alumina substrate and five different weight loadings of iron oxide. Each diffraction pattern has the same vertical scale. The individual diffraction traces have been vertically offset to provide for easier viewing.
observed here to those found by others for similar systems. Under the experimental conditions used here, the gas-phase reaction between DMMP and O3 only converts approximately 2.5% of the DMMP to CO2 on a per molecule basis (assuming one DMMP molecule yields one molecule of CO2) so that the concentrations of the reactants are virtually constant during the measurements. The steady-state rate for CO2 formation can be used to calculate a lower limit to the bimolecular rate constant for the reaction between O3 and DMMP. This rate constant is found to be 7.3 × 10-20 cm3/mol‚s, where the CO2 formation rate was calculated by using the measured amount of CO2 produced in the infrared gas cell volume with the time scale provided by the gas flow rate through the cell and using the initial gas-phase concentrations of the reactants. For comparison, reactions between ozone and alkenes exhibit rate constants 1-2 orders of magnitude higher under similar conditions, illustrating that the DMMP/O3 reaction is relatively slow.17 The calculated lower limit is, however, 2-10 times faster than that observed for ozone with acetaldehyde,18,19 and Reed and Oyama found no measurable gas-phase reaction of ozone with acetone at temperatures less than 400 °C.20
O3 conc (ppm)
CO2 produced (µmol)
CO produced (µmol)
CH3OH produced (µmol)
0 50 100 200 300 400 500 1000 1500
0 16.7 35.6 44.4 49.1 50.0 50.6 61.3 69.7
0 0 0 0 0 2.1 3.4 5.2 7.9
14.1 4.5 2.5 0 0 0 0 0 0
Surface Reaction. The surface areas of the alumina substrate and the impregnated alumina materials were all determined to be within the range 364 ( 5 m2/g, with no apparent trend in surface area with iron loading. The powder X-ray diffraction patterns observed for several of these materials are shown in Figure 3. The η-Al2O3 substrate shows a number of sharp features that are due to the more crystalline nature of this material as compared to γ-Al2O3. The sharp features are obscured as the iron oxide loading is increased, but there are no features observed that can be correlated with crystalline Fe2O3. In past experiments examining the decomposition of DMMP on oxides in the presence of oxygen but not ozone, the primary decomposition product was found to be methanol with a minor product being dimethyl ether. In the current studies of ozone reacting with DMMP on Fe2O3/Al2O3, methanol is formed at low ozone concentration, less than 200 ppm, but only CO2 and CO are observed as products at higher ozone concentrations (see Table 1 above). From Table 1, it seems clear that CO2 is replacing methanol as the primary reaction product at 25 °C when ozone is added. However, in experiments with more than 50 ppm ozone, there is much more CO2 produced than could be accounted for by assuming that the methanol formed in the absence of ozone is converted to carbon dioxide when ozone is present. Figure 4 shows the total yields of CO2 observed after exposure of a series of adsorbents to 650 µmol DMMP in the presence of 500 ppm ozone. As demonstrated by the results for the empty reactor, the gas-phase reaction between DMMP and ozone
9420 J. Phys. Chem. C, Vol. 111, No. 26, 2007 accounts for a significant fraction of the CO2 generated in these experiments. The silica MCM-41 material showed little activity for the decomposition of DMMP with ozone over that observed with no oxide present (Figure 4). This result is not particularly surprising. We have conducted a number of previous studies exposing DMMP to a variety of forms of silica including MCM41, SBA-15, and fumed silica and have used silica as a support for more reactive oxides such as iron oxide, cerium oxide, or copper oxide. In the absence of ozone, none of these materials has shown any significant DMMP decomposition.21 This lack of reaction on silica has, in fact, been examined as a strategy for pre-filtering chemical agents to enhance the selectivity of sensors for these agents.22,23 This lack of surface reactivity is most likely due to the absence of Lewis acid sites on silica. When DMMP initially interacts with alumina, it adsorbs on Lewis acid sites through interaction with the phosphoryl oxygen to form an activated DMMP species.1,24,25,26 This activated form is the proposed precursor to DMMP decomposition at low temperature and is likely the only significant pathway for the decomposition at these temperatures. From Figure 4, it can be seen that the 15 wt % Fe2O3/Al2O3 material is the most effective solid, with the 20, 25, and 30 wt % Fe2O3 materials showing less production of CO2 than the bare alumina substrate. One possible explanation for this result is that the 20 wt % and higher Fe2O3/Al2O3 materials represent a surface structure that is not particularly active for DMMP decomposition and that dispersion of the iron oxide on the alumina substrate may be an important determining factor. The 10 wt % Fe2O3 material has a higher decomposition capacity than bare alumina, though not the highest, but transitional structures that might be present on the 15 wt % Fe2O3 material should not be as likely to form on the lower weight loading material. For this reason, the 10 wt % material was chosen for the remainder of the study. The ceria on alumina material, even though it had shown high DMMP conversion to methanol in past studies without ozone,6 showed only slight improvement over the unimpregnated alumina material with regard to conversion to CO2. In fact, the amount of CO2 formed with the ceria on alumina reactive adsorbent is within experimental error of the amount of methanol formed in the absence of ozone, so that for ceria on alumina the net effect of adding ozone is to convert methanol to CO2, with no obvious increase in DMMP decomposition. In the results that follow, the CO2 production rate and total production values for the gas-phase reaction have been subtracted from the data presented. The plots showing the CO2 formed due to the surface reaction for 100 ppm ozone and for 1000 ppm ozone, as a function of DMMP exposure, are shown in Figure 5. The period between the beginning of the experiment and the observation of CO2 in the infrared cell is due in large part to the adsorption of CO2 on the adsorbent surface. CO2 is known to adsorb on the alumina surface,27-29 forming primarily bicarbonate and carbonate species depending on the temperature of the experiment, the dehydroxylation temperature of the solid, and the availability of hydroxyl and aluminum cation sites. Given the relatively low dehydroxylation temperature used for the reactive adsorbent and the low reaction temperature, there will be a significant population of hydroxyl groups on the surface, and the CO2 will preferentially react with the surface hydroxyl groups to form monodentate bicarbonate species. The emergence of CO2 from the microreactor corresponds to saturation of these CO2 adsorption sites.
Mitchell et al.
Figure 5. Data for the CO2 production in the presence of 100 ppm O3 (plot A) and 1000 ppm O3 (plot B). The plots show the corrected total CO2 production curves, obtained by subtracting the production of CO2 observed without a reactive adsorbent from that obtained with a reactive adsorbent. The small circles correspond to the experimental data, while the thin lines correspond to the fit of the data to eq 1. The arrows indicate the point beyond which the data points were used to determine the fit to eq 8.
Figure 5 shows the conversion of DMMP to CO2 as a function of DMMP exposure. The small open circles are the observed data, whereas the solid line is a fit to the data, described below. The slope of the conversion curve is called the fractional conversion, defined as the incremental change in CO2 produced for an incremental exposure to DMMP (∆CO2/∆DMMP). Initially, there is a relatively rapid increase in the amount of CO2 observed as a function of DMMP exposure and a relatively high fractional conversion. As the exposure to DMMP increases, the slope of the conversion curve, and hence the fractional conversion, decreases. The flow of CO2 from the reactor can be fit to an equation of the form
CO2 produced )
a{DMMP exposure} 1 + b{DMMP exposure}
(1)
The term {DMMP exposure} refers to the total amount of DMMP that has flowed into the microreactor, and the variables a and b are fitting parameters.
Room Temperature Reaction of Ozone and DMMP
J. Phys. Chem. C, Vol. 111, No. 26, 2007 9421
After breakthrough of CO2 from the adsorbent bed, the CO2 flow increases rapidly to a maximum value and then slowly declines. The initial increase in the flow rate of CO2 from the reactor is determined by the initial surface reaction and transport of CO2 through the adsorbent. However, once the maximum is reached, the slow decrease in CO2 flow is fit well by the simple function of eq 1. The fitting parameters are adjusted to optimize the fit to times after the maximum flow rate is reached. These points are indicated by arrows on the plots in Figure 5. All experimental data points to the right of the arrows are included in the fits. The initial data points, those points to the left of the arrows, reflect a lower fractional conversion than that predicted by eq 1, due to the relatively low initial CO2 concentration in the flow from the reactor. Surface Mechanisms. Two different reaction mechanisms dominate the discussion of surface reactions: the LangmuirHinshelwood mechanism and the Eley-Rideal mechanism.30,31 The Langmuir-Hinshelwood mechanism involves a reaction between two adsorbed species, whereas the Eley-Rideal mechanism is a reaction between an adsorbed species and a gas-phase co-reactant. The usual expression for the rate of a heterogeneous bimolecular reaction that takes place between two adsorbed species, A and B, via the Langmuir-Hinshelwood mechanism is
rate ) kθAθB
(2)
where θA and θB are the surface coverages of reactants A and B, respectively. The surface coverage, θ, for each compound is taken to be determined by the Langmuir isotherm, involving an equilibrium between the gas-phase and adsorbed forms of a compound. The Langmuir isotherm predicts the equilibrium surface coverage of a compound as a function of its vapor pressure and assumes that adsorption does not occur beyond monolayer coverage. The mathematical expression of the Langmuir isotherm for one-component adsorption is
θ)
Kp 1 + Kp
(3)
in which θ, the surface coverage of the compound, is dependent on p, the vapor pressure of the compound and K, the equilibrium constant for adsorption, K ) kadsorption/kdesorption. In the experiments described here, a reactant gas mixture is formed containing a fixed gas-phase concentration of DMMP mixed with varying amounts of ozone, from approximately 1/15 to 2 times the DMMP concentration. Over this 30-fold increase in ozone concentration, from 50 to 1500 ppm, the DMMP breakthrough point changes by a little more than 10%, and the small change that is seen is in the direction opposite that expected for surface site competition. This strongly suggests that the DMMP surface coverage is determined by the vapor pressure of DMMP, and its equilibrium constant for adsorption and is only minimally, if at all, affected by the presence of ozone. For the interaction and decomposition of DMMP on oxide surfaces at low temperature, the accepted mechanism for the interaction leading to decomposition is adsorption of DMMP at Lewis acid sites.1,24 Adsorption of DMMP at other sites including Brønsted acid sites leads to physisorption. Lundie et al. have determined that there are 5.5 × 1013 medium and strong Lewis acid sites per cm2 on the surface of η-Al2O3 from pyridine adsorption and desorption studies.32 The alumina used in that study was pretreated by heating to 350 °C, similar to the pretreatment used in this study. For our samples, this would
correspond to 1.25 × 1019 sites present in the sample (total), or 20.7 µmol, for the pure alumina material. The breakthrough point for DMMP in these experiments occurs after exposure to between 230 and 260 µmol DMMP. This indicates that there exist a large number of adsorption sites on the surface, a relatively small fraction of which, on the order of 10%, potentially lead to unimolecular decomposition. Our earlier studies of DMMP adsorbing on similar surfaces have shown that even when very small amounts of DMMP are dosed onto the adsorbent surface, with long periods of observation between doses, DMMP does not appear at the exit of the adsorbent bed until a significant fraction of the surface adsorption sites, including sites that do not lead to reaction, are occupied.6 These observations indicate that the DMMP equilibrium constant for adsorption is very high. The ozone interaction with the surface is more dynamic than that of DMMP, with ozone rapidly adsorbing, desorbing, decomposing, and regenerating the surface sites. Ozone adsorbs on Lewis acid sites as shown by Thomas et al., who determined that for ozone the primary sites for adsorption on alumina at 77 K are Lewis acid sites.33 When ozone adsorbs, it blocks the site only temporarily, forming molecular oxygen along with a surface oxide or peroxide species, and eventually regenerating the surface site. The process is catalytic to a degree but results in the slow deactivation of the alumina surface and loss of ozone decomposition activity.34,35 When DMMP adsorbs and reacts on a Lewis acid site, it poisons the site for further reactive adsorption by DMMP or by ozone. Sequential Adsorption Experiments. Two sets of experiments using sequential exposures of the reactive adsorbent surface to DMMP and ozone were carried out. In the first experiment, the reactive adsorbent was exposed to DMMP without added ozone, the adsorbent was then purged with pure helium to remove weakly adsorbed DMMP, and then the adsorbent was exposed to ozone (in O2) alone. The results are shown in Figure 6 (top). Methanol is observed from the decomposition of DMMP on the reactive adsorbent in the absence of ozone, as is expected, and then once ozone is added after the purge by He, CO2 is produced. If the sequence is carried out in the reverse order (Figure 6, bottom) and the adsorbent is exposed first to ozone, ozone breakthrough is not observed due to the surface decomposition reaction. However, there is a small amount of CO2 produced once DMMP (in He) is admitted to the microreactor, followed by the production of methanol as was observed in the absence of ozone. The results from Figure 6, top panel, demonstrate unequivocally that the CO2 production is from the reaction of adsorbed DMMP or decomposition products, not gas-phase DMMP, with ozone. The results shown in the bottom panel suggest that the reaction to form CO2 takes place via a surface-active oxygen species formed from the reaction of ozone with a surface site. Rate Expression. In order to evaluate the observed ozone dependence, an expression that describes the dependence of the reaction rate on the ozone concentration is needed. For two species that competitively adsorb on a reactive surface, the Langmuir-Hinshelwood rate expression, eq 2, becomes30
rate ) k
KDMMPpDMMPKO3 pO3 (1 + KDMMPpDMMP + KO3 pO3)2
(4)
Because the DMMP surface coverage is insensitive to the presence of ozone, and given the high affinity of DMMP on these surfaces, it makes sense to develop an expression for the
9422 J. Phys. Chem. C, Vol. 111, No. 26, 2007
Mitchell et al.
Figure 6. Top graphs show the flow and accumulation of products from the microreactor when the adsorbent is exposed to DMMP in He first, followed by purging with pure He, followed by exposure to O3 in O2. The bottom graphs show the flow and accumulation of products then the adsorbent is exposed to the reactants in the reverse order, first O3 in O2, followed by DMMP in He. BTP indicates the DMMP breakthrough point.
rate at a specific DMMP surface coverage. Calling this specific s and the corresponding rate rates, eq coverage of DMMP θDMMP 4 becomes
rate
s
)kθsDMMP
KO3 pO3 (1 + KDMMP pDMMP + KO3 pO3)
(5)
By introducing a modified equilibrium constant for ozone adsorption, K′O3 ) KO3/1 + KDMMP pDMMP, eq 5 can be written as
rates )k
θsDMMP
KO3 pO3
(1 + KDMMP pDMMP) (1 + K′O3 pO3)
(6)
When the DMMP gas-phase concentration is held constant, pDMMP is constant and rates only varies with ozone concentration,
and that variation is shown in the last term on the right of eq 6. This can be shown more explicitly by rewriting eq 6 as
rates ) k′
K O 3 pO 3 1 + K′O3 pO3
(7)
s )/(1 + KDMMP pDMMP), and k′ is constant at where k′ ) k(θDMMP a fixed DMMP surface coverage and pDMMP. In these experiments, the formation of CO2 has been measured, not the disappearance of ozone or DMMP. Thus, the rate results and the conclusions refer to the rate-limiting step for CO2 production, and it is not strictly correct to refer to these determinations as the rate of the reaction. The rate of formation of CO2, d(CO2)/dt, as a function of ozone concentration, with constant DMMP vapor concentration, can be estimated from the DMMP to CO2 conversion curves for the different ozone
Room Temperature Reaction of Ozone and DMMP
J. Phys. Chem. C, Vol. 111, No. 26, 2007 9423
Figure 7. Fractional conversion of DMMP (µmol CO2 out/µmol DMMP in) observed during the initial reaction period prior to DMMP breakthrough, after exposure to 100 µmol DMMP, and after breakthrough, after exposure to 300 µmol DMMP, plotted as a function of ozone concentration. The fractional conversion was determined by evaluating the slope of the CO2 formation as a function of DMMP from the data, not fitted curves. The contribution to CO2 production by the gas-phase reaction has been subtracted from these values.
concentrations (the 100 and 1000 ppm O3 conversion curves are shown in Figure 5). The finite rate (∆CO2/∆t) at a particular DMMP surface coverage is proportional to the fractional conversion, (∆CO2/∆DMMP), since the flow of DMMP is constant and ∆DMMP is proportional to ∆t. A plot of the fractional conversion determined after exposure to 100 µmol DMMP and 300 µmol DMMP is shown in Figure 7, which shows a plateau in the fractional conversion of DMMP as a function of ozone concentration. The plateau appears to begin in the neighborhood of 300-400 ppm ozone, although the data suggest that the plateau is reached at lower ozone concentration for higher DMMP exposures. The fractional conversion values used in Figure 7 were values calculated from the fit of the data to eq 1 (see Figure 5). For the 100 µmol DMMP exposure, the plateau value for the fractional conversion is determined to be 0.12 ( 0.01 µmol(CO2)/µmol(DMMP). The data were fit to an equation of the form
fractional conversion )
m[O3]g 1 + n[O3]g
(8)
in which m and n are fitting parameters. Equation 8 contains the same ozone dependence as does eq 7, and the agreement between the data and the curve supports the assertion that the reaction is taking place via the Langmuir-Hinshelwood mechanism. The Langmuir-Hinshelwood mechanism has been the mechanism proposed in a number of studies for heterogeneous reactions between ozone and organic compounds. Kwamena et al. found that the reaction between anthracene preadsorbed on a pyrex surface and ozone takes place via the LangmuirHinshelwood mechanism.36 An earlier study by this same research group (Kwamena et al.37) also invokes the LangmuirHinshelwood mechanism for the reaction between benzo[a]pyrene and ozone on an organic aerosol. These two studies are similar to the current one in that they examine a heterogeneous reaction between an adsorbed organic molecule and ozone, and reaction products remain on the surface, so the reactions are
not catalytic. Studies of catalytic reactions between organic molecules and ozone have found that the Langmuir-Hinshelwood mechanism is active for those systems, as well. Studies by Oyama’s group (Reed and Oyama20 and Xi et al.38) suggest that the reaction between ozone and acetone on manganese oxide is between two adsorbed species. One potential complicating factor is the adsorption of reaction products on the adsorbent surface. CO2 and CO also adsorb on these surfaces, and both can react with an aluminum cation and an adjacent hydroxyl group to form bicarbonates.27,29 But the formation of a monomeric bicarbonate species should not affect the availability of surface Lewis acid sites,28,39 and to the extent that it does, it will reduce the total number of available surface sites and should not affect the rate dependence on the coverages of the reactants. One experiment at a lower DMMP vapor concentration was carried out by lowering the temperature of the thermostatted enclosure containing the DMMP bubbler and aerosol trap to 20 °C. The vapor pressure of DMMP at 25 °C is, based on our measurements, 0.87 Torr. Using the reported heat of vaporization for DMMP of 52.33 kJ/mol and the Clausius-Clapyron equation, the DMMP vapor pressure at 20 °C is calculated to be 0.61 Torr. The fractional conversion after exposure to 100 µmol DMMP was determined to be 0.11 µmol(CO2)/µmol(DMMP), equal to the plateau value determined for the higher DMMP vapor pressure experiments within experimental error. There are other experimental observations that are important to the development of an understanding of the surface reaction between ozone and DMMP. One is that, in the absence of DMMP, ozone is never observed in the flow from the reactor indicating that ozone is completely decomposed on the oxide surface. When DMMP is present, however, ozone is observed in the exit flow from the reactor prior to the observation of DMMP (after exposure to approximately 140 µmol DMMP). The ozone breakthrough changes only slightly with ozone concentration, showing a slight trend toward earlier breakthrough with higher ozone concentrations. The observation of ozone in the flow from the microreactor indicates that DMMP is adsorbing on and poisoning the surface sites that lead to the decomposition of ozone. The ozone concentration at the exit of the reactor does not approach the inlet concentration until DMMP breakthrough is observed. DMMP breakthrough, on the other hand, occurs much later than that of ozone and shows a stronger dependence on breakthrough with ozone concentration. Breakthrough occurs after exposure to 205 µmol DMMP with no ozone present, but with only 50 ppm ozone present the breakthrough point increases to 230 µmol DMMP. The breakthrough point appears to plateau at 260 µmol DMMP for ozone concentrations in excess of 300-400 ppm. This seems to indicate that one of the effects of ozone is to increase the apparent number of surface sites for DMMP adsorption. Another observation is that when methanol in helium is used instead of DMMP to react on the surface with ozone at room temperature, no methanol or ozone is ever observed at the reactor exit. Methanol is quantitatively converted to CO2. The corresponding gas-phase reaction between ozone and methanol shows only minor conversion of methanol to CO2 indicating that the heterogeneous reaction between methanol and ozone is catalytic on this surface. Discussion The iron oxide appears to disperse well over the alumina substrate, as indicated by the lack of additional features in the powder XRD results that might be attributed to the formation
9424 J. Phys. Chem. C, Vol. 111, No. 26, 2007 of crystalline Fe2O3 or other crystalline iron oxide forms. Additionally, the surface area is virtually constant, suggesting that the iron oxide forms a surface structure that resembles that of the alumina substrate. Reaction Mechanism. Examination of Figure 6 shows that when DMMP adsorbs in the absence of ozone, top panel, a certain fraction of the DMMP decomposes, indicated by the evolution of methanol from the reactor. On the basis of earlier studies by us and by others,1-8,24,25 decomposition proceeds via the formation of a surface-bound methyl methylphosphonate fragment and a surface methoxy group, a fraction of which subsequently gives rise to methanol that evolves from the reactor. When ozone is added to this sample, a significant amount of CO2 is produced due to reaction between ozone and the remaining DMMP fragments. When the surface is exposed to ozone first, bottom panel, the ozone is completely decomposed on the surface, giving rise to some active surface oxide species, either an oxide or a peroxide,34 and molecular oxygen that evolves from the reactor. When DMMP is added after flowing the ozone into the reactor for more than 2 h, only a small amount of CO2 is formed, indicating either that the number of active oxide species present is relatively low or their lifetime is short on the time scale of the experiment. Methanol also forms from DMMP decomposing as it would in the absence of ozone. The primary indication that a surface oxide species formed by the action of ozone is the reactive oxidant and not gas-phase ozone is the plateau in fractional conversion observed as a function of ozone concentration shown in Figure 7. For any particular DMMP coverage, only a certain number of sites are available for O3 adsorption. The reaction rate will increase with ozone gas-phase concentration until the remaining surface sites are occupied, after which the observed rate will remain constant with increasing gas-phase ozone concentration. The fact that at higher DMMP exposures (300 µmol) the apparent plateau in the rate occurs at lower ozone concentration, and corresponds to a lower fractional conversion, indicates that fewer adsorption sites are available for ozone adsorption. If the Eley-Rideal mechanism were the appropriate one, with increasing concentration of ozone in the gas phase, the initial reaction rate would continuously increase since there would be no competition for surface sites between the reactants. Another observation was that when the DMMP concentration in the gas phase was reduced by lowering the temperature in the thermostatted enclosure, the fractional conversion remained the same. Given the very high affinity of DMMP for the surface, even though the vapor pressure of DMMP was reduced, the fractional conversion corresponds to approximately the same s DMMP surface coverage, θDMMP , as that used for the higher vapor pressure experiments. Since the ozone concentration was in the “plateau” regime, above 300 ppm, the fractional conversion is expected to be the same. s The fixed DMMP coverage, θDMMP , referred to in the model development does not really refer to a particular coverage as it is used here but rather to a specific exposure. The DMMP surface coverage, θDMMP, that corresponds to a particular fractional conversion is a slowly varying average coverage that changes with DMMP exposure. An initial surface coverage gradient is developed during the induction period, prior to the appearance of CO2 in the infrared cell. The coverage gradient moves through the adsorbent bed changing only slowly with exposure, until the leading edge of the DMMP gradient reaches the end of the adsorbent bed and DMMP breakthrough occurs. Once breakthrough occurs, the entire surface approaches the equilibrium surface coverage, determined by the Langmuir
Mitchell et al. isotherm, as the DMMP exposure continues. As the DMMP surface coverage increases, fewer sites are available for ozone adsorption, and the DMMP fractional conversion decreases. This is manifested in the experimental results by the slowly decreasing fractional conversion of DMMP to CO2 as seen in Figure 5. The reason that the model, eq 6, works for these experiments is that the processes governing DMMP and O3 adsorption on the surface appear to operate in much different fashions and on different timescales. Equation 6 effectively decouples the ozone and DMMP surface coverages in the rate expression and s , to account for introduces a slowly varying parameter, θDMMP the loss of available surface sites due to the increasing DMMP exposure. One possible mechanism for the formation of CO2 is the reaction between methanol formed from the decomposition of DMMP and ozone. However, on the Fe2O3/Al2O3 surface, much more CO2 is produced than could be accounted for simply by a conversion of methanol to CO2 by ozone (see Table 1). Additionally, in Figure 6 it is shown that after dosing the adsorbent with DMMP and purging the reactor with helium to remove the methanol, the addition of ozone results in significant formation of CO2 indicating that the reaction to form CO2 is at least in part due to the reaction of ozone with more strongly adsorbed species, adsorbed methoxy groups, methyl methylphosphonate fragments, or strongly adsorbed DMMP. Enhanced Decomposition. For the supported iron oxide material, the enhanced DMMP decomposition activity observed when ozone is present probably derives from either the ozone effectively making a larger fraction of the adsorption sites active for decomposition or the ozone providing a way for recycling a portion of the adsorption sites. It is quite possible that under the reaction conditions at room temperature without ozone present, a significant fraction of the methoxy groups that result from the decomposition of DMMP are not released as methanol. The standard model for the interaction for DMMP with metal oxides is1,24
DMMP + * f DMMP*
(8)
DMMP* + HO* f MMP* + CH3OH* + O*
(9)
where * represents a surface adsorption site, nominally a surface aluminum site, and MMP is the methyl methylphosphonate fragment. This reaction sequence generates the methanol that eventually evolves from the reactor. However, there is clear evidence for the formation of surface methoxy groups as a result of the adsorption of DMMP on oxide surfaces.6 This reaction is postulated to occur via decomposition of the adsorbed methanol formed in eq 97
CH3OH* + O* f CH3O* + HO*
(10)
Once the surface methoxy group is formed or a gas-phase methanol is reactively adsorbed, there may not be enough energy available to accomplish the reaction
CH3O* + HO* f CH3OH(v) + *O*
(11)
and some of the methoxy groups that result from the decomposition of DMMP remain on the surface (the “up arrow” (v) is used to indicate species that may evolve from the surface). In the presence of ozone, however, another reaction involving the surface methoxy group becomes possible. Xi et al.37 suggest that for the heterogeneous reactions between ozone and organic compounds each ozone molecule can provide one active oxygen,
Room Temperature Reaction of Ozone and DMMP
J. Phys. Chem. C, Vol. 111, No. 26, 2007 9425
which leads to the following possible stoichiometry
is converted to CO2, assuming that each molecule of CO2 indicates the decomposition of one molecule of DMMP. After exposure to 300 µmol DMMP, the limiting fractional conversion drops to 0.06. The surface reaction of ozone with DMMP converts much more of the adsorbed DMMP to gas-phase products, on a per molecule basis, with approximately 2.7 times as much gas-phase carbon liberated in the form of CO2 and CO as was liberated in the form of methanol in the absence of ozone. Part of the increased activity may be due to ozone providing a mechanism for converting adsorbed methoxy groups to CO2, while some of the additional activity may come from the recycling of surface sites. The reaction between ozone and DMMP requires the presence of Lewis acid sites for the activation of adsorbed DMMP. However, the increased reactivity of the alumina-supported iron oxide material compared to alumina appears to have more to do with the oxide’s activity for the decomposition of ozone rather than a particular adsorption interaction between DMMP and the supported iron oxide.
CH3O* + 3O3(ads) f HO* + CO2(v) + H2O(v) + 3O2(v) (12) The reaction shown in eq 12 is energetically more favorable than that shown in eq 11, given the formation of carbon dioxide and water as reaction products, and it provides a way to recycle some of the surface sites. In addition, a Brønsted site is created that can participate in the reaction with other adsorbed methoxy groups via eq 11, or with adsorbed DMMP via eq 9. The adsorption of DMMP on oxide surfaces gives rise to methanol and adsorbed methoxy groups. As mentioned in the results, the heterogeneous bimolecular reaction between ozone and methanol quantitatively converts methanol to CO2 at room temperature. Thus, it seems most likely that the reaction between an adsorbed methanol or a surface methoxy group and an ozonegenerated surface active oxygen species is the most important contributor to the formation of CO2. However, it is possible that adsorbed molecular DMMP may yield CO2 in the presence of the active oxygen species, giving rise to an increased DMMP decomposition yield when ozone is added. The nature of the enhanced reactivity of the aluminasupported iron oxide material compared to the alumina itself may have to do with either the nature of the DMMP/surface interaction or the O3/surface interaction. In the absence of ozone, the alumina-supported iron oxide appears to promote a unique surface decomposition mechanism compared to other oxides,2 although the decomposition activity is approximately equal to that of alumina per unit surface area.3 However, assuming that the fragmentation still yields a surface methoxy fragment, the DMMP/surface aspect of the reaction system probably does not account for the enhanced activity. However, supported iron oxide is an effective ozone decomposition catalyst, and evidence suggests that the heterogeneous oxidation reactions of ozone with organic compounds involve atomic or molecular oxygen that result from the decomposition of O3 on the oxide surface as the primary oxidant.37 Thus it seems probable that it is the enhanced O3 decomposition activity of the aluminasupported iron oxide material, rather than an improved DMMP/ surface interaction, that makes the supported iron oxide material a more effective reactive adsorbent than alumina in the presence of O3. Conclusions The reaction between DMMP and ozone on aluminasupported iron oxide has been examined at room temperature. The reaction yields much larger amounts of gas-phase decomposition products from the surface reaction than experiments carried out using oxygen as the oxidant. The reaction follows the Langmuir-Hinshelwood mechanism and involves a reaction between adsorbed DMMP or methoxy fragments and a surfaceactive oxygen species from ozone, the exact nature of which is not known. The surface reaction appears to saturate when the concentration of ozone exceeds 300-400 ppm. Higher concentrations of ozone have no observable effect on the surface reaction. The gas-phase reaction, however, does show an increased rate with increasing ozone concentration, although the gas-phase reaction is not a particularly effective method for decomposing DMMP, given the low rate constant. During the initial stages of DMMP adsorption and reaction, after exposure to 100 µmol DMMP, the limiting fractional conversion is found to be 0.12, or 12% of the incoming DMMP
Acknowledgment. This research was supported by a grant from the Army Research Office, Grant Number W911NF-041-0377. In addition, the work was partially supported by The WaterCAMPWS, a NSF Science and Technology Center (Center of Advanced Materials for the Purification of Water with Systems) under the National Science Foundation agreement number CTS-0120978, and by NASA through the Clark Atlanta University High Performance Polymers and Composites Center (NCC3-552). References and Notes (1) Mitchell, M. B.; Sheinker, V. N.; Mintz, E. A. J. Phys. Chem. 1997, 101, 11192-11203. (2) Tesfai, T. M.; Sheinker, V. N.; Mitchell, M. B. J. Phys. Chem. 1998, 102, 7299-7302. (3) Sheinker, V. N.; Mitchell, M. B. Chem. Mater. 2002, 14, 12571268. (4) Mitchell, M. B.; Sheinker, V. N.; Tesfamichael, A. G.; Gatimu, E. N.; Nunley, M. J. Phys. Chem. B 2003, 107, 580-586. (5) Mitchell, M. B.; Sheinker, V. N.; Cox, W. W., Jr. New Active Decomposition Materials for Sorbent Decontamination. In Proceedings of the 2002 Scientific Conference on Chemical and Biological Defense Research; Battelle Eastern Science and Technology Center: Aberdeen, MD, 2003. (6) Mitchell, M. B.; Sheinker, V. N.; Cox, W. W., Jr.; Gatimu, E. N.; Tesfamichael, A. B. J. Phys. Chem. B 2004, 108, 1634-1645. (7) Li, Y. X.; Klabunde, K. J. Langmuir 1991, 7, 1388-1393. (8) Segal, S. R.; Suib, S. L.; Tang, X.; Satyapal, S. Chem. Mater. 1999, 11, 1687-1695. (9) Obee, T. N.; Satyapal, S. J. Photochem. Photobiol. A: Chem. 1998, 118, 45-51. (10) Rusu, C. N.; Yates, J. T., Jr. J. Phys. Chem. B 2000, 104, 1229912305. (11) Decker, S. P.; Klabunde, J. S.; Khaleel, A.; Klabunde, K. J. EnVion. Sci. Technol. 2002, 36, 762-768. (12) Jiang, Y.; Decker, S.; Mohs, C.; Klabunde, K. J. J. Catal. 1998, 180, 24-35. (13) Wagner, G. W.; Procell, L. R.; O’Connor, R. J.; Munavalli, S.; Carnes, C. L.; Kapoor, P. N.; Klabunde, K. J. J. Am. Chem. Soc. 2001, 123, 1636-1644. (14) Huling, S. G.; Arnold, R. G.; Sierka, R. A. Contaminant Adsorption and Oxidation via the Fenton Reaction. U.S. Patent No. 20,040,134,857, 2004. (15) Soria, J.; Coronado, J. M.; Conesa, J. C. J. Chem. Soc., Faraday Trans. 1996, 92, 1619-1626. (16) Fernandez-Garcia, M.; Gomez Rebollo, E.; Guerrero Ruiz, A.; Conesa, J. C.; Soria, J. J. Catal. 1997, 172, 146-159. (17) Atkinson, R.; Carter, W. P. Chem. ReV. 1984, 84, 437-470. (18) Stedman, E. H.; Niki, H. EnViron. Lett. 1973, 4, 303-310. (19) Atkinson, R.; Aschmann, S. M.; Winer, A. M.; Pitts, J. N., Jr. Int. J. Chem. Kinet. 1981, 13, 1133-1142. (20) Reed, C.; Lee, Y.-K.; Oyama, S. T. J. Phys. Chem. B 2006, 110, 4207-4216.
9426 J. Phys. Chem. C, Vol. 111, No. 26, 2007 (21) Mitchell, M. B.; Sheinker, V. N.; Cox, W. W., Jr. Unpublished results. (22) Kanan, S. M.; Tripp, C. P. Langmuir 2001, 17, 2213-2218. (23) Kanan, S. M.; Tripp, C. P. Langmuir 2002, 18, 722-728. (24) Templeton, M. K.; Weinberg, W. H. J. Am. Chem. Soc. 1985, 107, 97-108. (25) Templeton, M. K.; Weinberg, W. H. J. Am. Chem. Soc. 1985, 107, 774-779. (26) Aurian-Blajeni, B.; Boucher, M. M. Langmuir 1989, 5, 170-174. (27) Cabrejas Manchado, M.; Guil, J. M.; Pe´rez Masia´, A.; Ruiz Paniego, A.; Trejo Menayo, J. M. Langmuir 1994, 10, 685-691. (28) Rosynek, M. P. J. Phys. Chem. 1975, 79, 1280-1284. (29) Parkyns, N. D. J. Phys. Chem. 1971, 75, 526-531. (30) Atkins, P.; de Paula, J. Physical Chemistry, 8th ed.; W. H. Freeman and Company: New York, 2006. (31) Baxter, R. J.; Hu, P. J. Chem. Phys. 2002, 116, 4379-4381.
Mitchell et al. (32) Lundie, D. T.; McInroy, Al R.; Marshall, R.; Winfield, J. M.; Jones, P.; Dudman, C. C.; Parker, S. F.; Mitchell, C.; Lennon, D. J. Phys. Chem. B 2005, 109, 11592-11601. (33) Thomas, K.; Hoggan, P. E.; Mariey, L.; Lamotte, J.; Lavalley, J. C. Catal. Lett. 1997, 46, 77-82. (34) Sullivan, R. C.; Thornberry, T.; Abbatt, J. P. D. Atmos. Chem. Phys. 2004, 4, 1301-1310. (35) Roscoe, J. M.; Abbatt, J. P. D. J. Phys. Chem. A 2005, 109, 90289034. (36) Kwamena, N.-O. A.; Earp, M. E.; Young, C. J.; Abbatt, J. P. D. J. Phys. Chem. A 2006, 110, 3638-3646. (37) Kwamena, N.-O. A.; Thornton, J. A.; Abbatt, J. P. D. J. Phys. Chem. A 2004, 108, 11626-11634. (38) Xi, Y.; Reed, C.; Lee, Y.-K.; Oyama, S. T. J. Phys. Chem. B 2005, 109, 17587-17596. (39) Rethwisch, D. G.; Dumesic, J. A. Langmuir 1986, 2, 73-79.
