Spectroscopy of Hydrothermal Reactions 17. Kinetics of the Surface

first below the critical temperature of water. The zeroth-order kinetics model produced an. Arrhenius activation energy of 32 ( 3 kcal/mol, which is i...
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Spectroscopy of Hydrothermal Reactions 17. Kinetics of the Surface-Catalyzed Water-Gas Shift Reaction with Inadvertent Formation of Ni(CO)4 Davide Miksa and Thomas B. Brill* Department of Chemistry and Biochemistry, University of Delaware, Newark, Delaware 19716

The rate constants and Arrhenius parameters for the reaction of CO in H2O were determined at 230-270 °C and 27.4 MPa by the use of a titanium flow reactor with real-time detection by infrared spectroscopy through sapphire windows. These rate measurements appear to be the first below the critical temperature of water. The zeroth-order kinetics model produced an Arrhenius activation energy of 32 ( 3 kcal/mol, which is in the range of previously reported values at higher temperatures, but the preexponential factor [ln(A, mol kg-1 s-1)] of 20.5 is much larger. The higher overall reaction rate is consistent with heterogeneous catalysis by the reactor surfaces considering (1) the zeroth-order kinetics, (2) the high A factor, (3) the activation energy in the range for the water-catalyzed reactions, and (4) the previously determined dependence of the decomposition rate of the putative formic acid intermediate on the metal used to construct the cell. Extremely toxic Ni(CO)4 was observed to form as a result of extraction of Ni from slightly corroded 316 stainless steel tubes that connected the cell/reactor to the flow control system. Ni(CO)4 formed under somewhat limited conditions, but its occurrence forewarns of the potential hazard of hydrothermal processing when a high CO concentration might be present in a nickelcontaining reaction vessel. Introduction The water-gas shift reaction 1 (hereinafter the WGSR) is counted among the core control reactions of high-temperature processes. For example, the homogeneous WGSR plays a major role in determining several product ratios from high-temperature events, such as explosion,1 combustion,2 and flash pyrolysis.3 As such, the WGSR can be employed in practical processes including the synthesis of raw materials from coke or coal where the temperature can be well controlled. Coal gasification produces H2 and CO in a ratio of approximately 1.

CO + H2O h CO2 + H2

(1)

Reaction 1 enables these ratios to be adjusted as desired by the choice of the gas temperature. H2/CO ratios of 2 and 3 optimize the stoichiometry for the production of methanol and methane, respectively. Heterogeneous catalysts typically used in the industrial processes include those composed of mixtures of CuO, ZnO, and Al2O3.4 Interest has developed recently in the kinetics of the WGSR in the single-phase hydrothermal environment, in which the density of H2O is much higher than that of steam,5-10 and is driven primarily by the use of supercritical water oxidation to denature organic compounds in waste streams. Kinetic measurements of the rate of the uncatalyzed, homogeneous, forward reaction 1 have been reported at temperatures in the range of 380-600 °C, which is above the critical temperature of H2O (374 °C), and at pressures in the range of 2-60 MPa, which includes the critical pressure (22.1 MPa).7,9,10 * Corresponding author. E-mail: [email protected].

The Arrhenius activation energies extracted from these data (23-35 kcal/mol) differ but are near the limit of the quoted experimental errors. The range of the A factors (103.3-107.2) somewhat exceeds the quoted errors. Direct comparisons of these measurements are difficult to make, however, because the data depend to some degree on the reactor. Kinetic data for reaction 1 were previously unavailable for the dense hydrothermal regime below the critical temperature of H2O. This portion of the phase diagram is important for geochemistry, such as with subsurface formation water, and for industrial processing applications,11 such as the preheat zone of a supercritical water oxidation reactor. To obtain kinetics information, the forward direction of reaction 1 was investigated with a flow-cell reactor12,13 operating at 230-270 °C and 27.5 MPa by following the rate of formation of CO2 in real time by FTIR spectroscopy. In addition to the rate measurements, an unexpected and potentially hazardous side process was found to occur, which could present a risk to hydrothermalists encountering high concentrations of aqueous CO in nickelcontaining reactors. CO in the homogeneous aqueous phase extracted nickel in the form of Ni(CO)4 from a somewhat corroded 316 stainless steel inlet tube of the flow reactor. Ni(CO)4 was detected only in a narrow range of operating conditions, but the estimated concentration of this extremely toxic product greatly exceeds the toxic limit of about 50 ppb. Ni(CO)4 would not have been detected and could result in inadvertent operator exposure had not real-time detection been employed for this work. Researchers must be aware of the potential formation of Ni(CO)4 with the use of nickelcontaining reaction vessels operating under similar conditions.

10.1021/ie010076j CCC: $20.00 © 2001 American Chemical Society Published on Web 06/09/2001

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Figure 1. Cross-sectional and end views of the short path length IR flow cell used for these studies.

Experimental Section Bubbling of CO into H2O housed in a stirred autoclave at 6 MPa produced aqueous CO at room temperature in sufficient concentration for detection by IR spectroscopy. After allowing 2 h for saturation to be approached, the solution was transferred under pressure to an Isco syringe pump. The resulting CO concentration was estimated to be 0.011 m based on the Henry’s Law constant of 9.5 × 104 mol/(kg bar).14 Ni(CO)4 was purchased from Powdermet Inc., Sun Valley, CA, so that a comparison of its spectrum in aqueous solution from this work could be made with an authentic sample. The flow controls and spectroscopy cell have been described previously in detail.12,13 Figure 1 shows the cell design. Briefly described, infrared spectral measurements in the transmission mode under hydrothermal conditions were accomplished by compressing a gold foil disk into which a slot had been cut, thus creating a path length of 20 µm between two sapphire windows. The slot enabled the fluid to enter the cell, pass by the windows as a thin sheet, and exit. The cell body was constructed of grade 2 titanium into which holes had been drilled to accept cartridge heaters and thermocouples. The cell was insulated between two Ca2SiO4 blocks. The inlet and exit tubes connected to the cell body were also made of 12-cm lengths of titanium, but these tubes were connected to the flow reactor by more flexible 316 stainless steel (SS) tubes. These steel tubes were determined by a thermocouple measurement to remain at about 25 °C during long experiments. Visual examination of the internal cell body showed no signs of corrosion. On the other hand, visual examination revealed mild corrosion of the stainless steel tubes, which had been used earlier for studies of the decarboxylation kinetics of CF3CO2H.15 This corrosion has consequences, which are discussed later in this article. A Visual Basic computer program enabled the temperature, pressure, and flow rate to be set and continuously controlled as desired to within (1 °C, ( 0.1 MPa, and an absolute error of (0.01 mL/min, respectively. The relative error in the flow rate is about (0.001 mL/ min. Temperatures of 230-270 °C at 27.4 MPa pressure

Figure 2. Temperature dependence of the molar absorptivity of aqueous CO in the range below the occurrence of reaction 1.

were required to obtain a measurable extent of reaction. A single fluid phase having a density based on pure water of 0.84-0.94 g/cm3 existed under these conditions. The flow rates used were in the range of 0.10-1.00 mL/ min, which converts to residence times of 45-4.5 s based on an internal cell volume of 0.0819 cm3. A correction was made for the density change of water with temperature at constant pressure. To determine the temperature dependence of the molar absorptivity of CO, the fastest flow rate (1 mL/min) and temperatures below the reaction temperature were used to ensure that no CO oxidation took place. Figure 2 shows the absorptivity of CO in this same temperature range and reveals a decrease of 12%. The reaction rate was measured in real time by following the change in the intensely absorbing asymmetric stretching mode of CO2 at 2343 cm-1 using a Nicolet Magna 560 FTIR spectrometer. Thirty-two scans were summed at 4 cm-1 resolution at each flow rate. After the water background at the same temperature and pressure had been subtracted, the CO2 absorption profile was fit with a four-parameter Voigt function to obtain the area.16 The rate constants resulting from these time-dependent concentrations were obtained by

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Figure 3. Selected IR spectra of the CO and CO2 stretching modes during the WGSR at 230 °C and 275 bar.

Figure 4. Zeroth-order rate plots for reaction 1 based on CO2 formation.

averaging 4-6 independent measurements at each temperature. Weighted least-squares regression was used to calculate the error in the Arrhenius data, with care taken during conversion to logarithmic space.17 Results and Discussion CO in Water. The small Henry’s Law constant belies the low solubility of CO in H2O, and it was necessary to pressurize CO and H2O together to increase the solubility of CO so that its IR absorption could be detected. Figure 3 contains selected IR spectra showing the forward step of reaction 1 as a function of time at 230 °C and 275 bar. Aqueous CO and CO2 each are characterized by a single broad absorption centered at 2153 and 2343 cm-1, respectively. The P and R rotational branches that appear in their gas-phase spectra are absent in the aqueous-phase spectra because of quenching of rotational motion on the vibrational time scale. The IR molar absorptivity of CO in the gas phase at 25 °C is 0.047 times that of the asymmetric stretching mode of CO2. This difference in the molar absorptivity was found to transfer roughly to the aqueous phase based on the comparison at 200 °C, where the ratio is 0.043. The area of the CO2 band at constant concentration increases as the temperature is increased. For example between 200 and 240 °C the absorptivity increases by 5%.16 In contrast, Figure 2 shows that the absorptivity of CO decreases by 12% over the same temperature range. These opposite trends suggest that interactions of CO and CO2 with H2O have different effects on the dipole oscillator. Such differences might be expected in view of the fact that CO has a permanent dipole moment, whereas the lowest moment of CO2 is a quadrupole. These absorptivity differences were useful for data interpretation but their origins were not investigated further. Also indicated in Figure 3 is the fact that Ni(CO)4 formed in the solution. This occurrence will be discussed later in this article. Kinetics of the WGSR. Despite the effort made to enhance the CO concentration in water, the rate of formation of CO2, rather than the disappearance of CO, provided the more sensitive measurement of the oxidation of CO by water in reaction 1. This is because CO2 has a higher IR absorptivity and a narrower line width, which permit a more accurate integration of the peak and thereby a more accurate measurement of the

Figure 5. Arrhenius plot of the rate constants for reaction 1. Table 1. Zeroth-Order Rate Constants for the Formation of CO2 by Reaction 1 T (°C)

k × 105 (mol/kg)

230 240 250 260 270

1.09 2.64 4.87 9.35 12.21

change in concentration. Figure 3 shows that the IR absorption intensity of CO2 at 230 °C and 27.4 MPa increases at longer residence times as a result of reaction 1. Evidence that a single liquid phase is maintained comes from the fact that no gas-phase P and R rotational branches separate from the single CO2 and CO bands. Conversion of the area of the CO2 peak into concentration16 yields the rate plot shown in Figure 4, which was interpreted to be zeroth-order. Because of the relatively low temperatures used here to study the reaction, the maximum extent of conversion of CO to CO2 for which data were taken was 22%. As a result, it is, in truth, not possible to distinguish the rate plot in Figure 4 from a first-order rate plot. Making the simplifying assumption of zeroth-order kinetics, however, provides the rate constants given in Table 1. As is shown in Figure 5, the Arrhenius plot of the rates yields an activation energy Ea ) 32 ( 3 kcal/mol and ln(A, mol/kg s) ) 20.5. The value of Ea obtained here resembles that given by Rice et al.9 of 34.7 ( 8.6

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kcal/mol for the first-order homogeneous WGSR at 410-520 °C when the H2O concentration is 20 M. Their accompanying preexponential factor, however, is ln(A, s-1) ) 16.8, which is much smaller than that determined in our work when our value of ln A is converted to the same units [ln(A, s-1) ) 25]. Other published experimentally determined Arrhenius values are Ea ) 22.7 kcal/mol and ln(A, s-1) ) 9 at 410-590 °C and 24.5 MPa in Hastelloy C2767 and Ea ) 23.5 kcal/ mol and ln(A, s-1) ) 6.6 at 380-440 °C and 10-30 MPa in 316 SS.10 Although the activation energy is roughly similar to the values determined from previous measurements, the notable feature in the present work is the unusually large A factor, which produces higher reaction rate compared to those determined in previous studies. This is a suggestion that the reactor walls might heterogeneously catalyze reaction 1 in our work. It is difficult to specify the mechanism by which catalysis occurs, but a starting point is the widely acknowledged formate species as an intermediate in reaction 1.2,18-22 Understanding about the molecular basis of homogeneous catalysis of the WGSR rests heavily on the computations of the reaction barrier in which one or more H2O molecules are involved in the transition state. Formic acid thus becomes the intermediate species, although it has not yet been detected experimentally. Ruelle et al.19 obtained an Ea value of 48.7 kcal/mol with the incorporation of one H2O molecule. Melius et al.2 calculated an activation energy of 35.6 kcal/mol for the transition state in which two H2O molecules are involved, whereas Akiya and Savage20 found 47.9 kcal/ mol with two H2O molecules in the transition state. These values of Ea are much smaller than those for the homogeneous uncatalyzed decarboxylation of formic acid in the gas phase, which are typically on the order of 60-70 kcal/mol.23-27 The activation energy of 32 (3 kcal/mol obtained from Figure 5 resembles the range of these calculated values and the previously obtained experimental values more closely than the range found for the uncatalyzed reaction. The similarity is therefore consistent with a catalyzed form of reaction 1. As noted above, the large A factor suggests the possibility of heterogeneous catalysis by the reactor walls. A formate species, which is the putative intermediate of reaction 1, is well-known to interact with metal surfaces,28-34 which heterogeneously catalyze the decomposition to CO2 and H2. Previous kinetic measurements of the hydrothermal decomposition of aqueous HCO2H were made at 280-330 °C and 27.4 MPa using the same cell design as employed here, and the rates were compared for different metals of cell construction,35 namely, 316 SS, grade 2 Ti, and a 90/10 Pt/Ir alloy. In all cases, a gold foil spacer was used to create the flat duct through which the fluid flowed. The decarboxylation rate depended on the material used to construct the cell, which is evidence that at least some heterogeneous catalysis is taking place. Because formic acid is believed to be the intermediate in reaction 1, reaction 1 is, by inference, probably heterogeneously catalyzed in the present study. Formation of Ni(CO)4. An additional piece of evidence that is consistent with the role of heterogeneous surface chemistry involving CO is the observation in Figure 3 of the potentially disastrous consequence that Ni(CO)4 is formed in the flow reactor under certain

Figure 6. Temperature dependence of the formation and decomposition of aqueous Ni(CO)4 at 275 bar. The formation of Ni(CO)4 is also shown and confirmed by comparison with the spectrum of an aqueous solution of authentic Ni(CO)4.

conditions. The only metals in the directly heated portion of the cell are titanium and gold. Titanium tubes of 12-cm length were used to connect the cell to the 316 SS tubes of the flow control apparatus. As noted in the Experimental Section, these connections to the SS tubes do not exceed a temperature of about 25 °C when the cell/reactor is operating at 280 °C. However, because of corrosion of these SS tubes resulting from their prior use in a study of reactions of CF3CO2H, it is likely that the internal surface area was larger and that the surface metals might have become somewhat chemically activated. It should be noted that no Ni(CO)4 was detected when fresh uncorroded 316 SS tubes were used to connect the cell to the flow control system. The aqueous CO was able to extract Ni from the corroded tubes in the form of Ni(CO)4 by the well-known Mond reaction 2.36

Ni + 4CO h Ni(CO)4

(2)

The Ni(CO)4 was then carried into the cell where it was identified by its strongly IR-active CO stretching mode of t2 symmetry located at 2044 cm-1. Figure 6 shows the behavior of Ni(CO)4 as a function of the cell temperature in one of the attempts to obtain kinetics for reaction 1. Because the residence time and temperature are each an adjustable variable, data for a constant residence time with a variable temperature are shown in Figure 6. Aqueous Ni(CO)4 was detected in various experiments with a cell temperature in the range of 200-240 °C and a pressure of 27.4 MPa. Above these temperatures, aqueous Ni(CO)4 is decomposed in the hot CO-H2O fluid mixture at the approximate residence time shown. The temperature range in which Ni(CO)4 is detected, however, is a strong function of the CO pressure, as eq 2 suggests. Without excess CO in the flow reactor, aqueous Ni(CO)4 was found to decompose at about 80 °C at a reasonably fast rate at 27.4 MPa.

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Figure 7. Concentration of aqueous Ni(CO)4 as a function of time at 230 °C and 275 bar.

The approximate concentration of Ni(CO)4 can be estimated using its gas-phase absorption coefficient37 of 4.1 × 104 atm-1 m-1 with the assumption that this value is approximately the same in aqueous solution. Curve resolution and integration of the area of the absorption in Figure 2 produces a concentration estimate in the range of 0-3.5 × 10-4 M or (0-39 ppm), as shown in Figure 7. The linearity of this plot until decomposition begins is consistent with an approximately zeroth-order rate of formation of Ni(CO)4. The concentration of 39 ppm is far in excess of the maximum allowable vapor value of 0.05 ppm and is of concern because Ni(CO)4 is poorly soluble in water and is expected to evaporate when the pressure is relieved. The conditions of formation of Ni(CO)4 in the hydrothermal reactor used here appear to be rather specialized, i.e., a large CO concentration dissolved in water and corroded 316 SS transfer tubes. Nevertheless, these conditions innocently existed in our system and realistically might be encountered elsewhere in laboratories or operations. Unless a sensitive real-time detection method is employed, such as is afforded by IR spectroscopy, the formation of Ni(CO)4 might be unnoticed. Because Ni(CO)4 is one of the most toxic metal carbonyl compounds, the consequences of improper venting or mishandling of the effluent are obvious. Although the formation of Ni(CO)4 from Ni and CO has long been known,36 to our knowledge, it has not been observed before under hydrothermal conditions. It is worthy of note, however, that the formation of other metal carbonyls under hydrothermal conditions has precedence. Recently, the iron carbonyl dimer [Fe2(RS)2(CO)6] was formed under hydrothermal conditions in the presence of FeS, CO, and RSH, where RSH is an alkanethiol.38 Further work is in progress to measure the kinetics of decarbonylation of metal carbonyls under hydrothermal conditions. Acknowledgment We are grateful to the Army Research Office for support of this work through Grant DAAG55-98-0253. Literature Cited (1) Bernecker, R. R.; Smith, L. C. On the Products Formed in the Combustion of Explosives. Freeze-out of the Water-Gas Reaction. J. Phys. Chem. 1967, 71, 2381. (2) Melius, C. F.; Bergan, N. E.; Shepherd, J. E. Effects of Water on Combustion Kinetics at High Pressure. In 23rd Symposium

(International) on Combustion; The Combustion Institute: Pittsburgh, PA, 1990; p 217. (3) Maiella, P. G.; Brill, T. B. Spectroscopy of Hydrothermal Reactions III. The Water Gas Reaction, “Hot Spots”, and Formation of Volatile Salts of NCO- from Aqueous [NH3(CH2)2NH3]NO3 at 720 K and 276 bar by T-Jump/FTIR Spectroscopy. Appl. Spectrosc. 1996, 50, 829. (4) Kochloefl, K. Development of Industrial Solid Catalysts. In Handbook of Heterogeneous Catalysis; Ertl, G., Knozinger, H., Weitkamp, J., Eds.; Wiley-VCH: Weinheim, Germany, 1997; p 1831. (5) Helling, R. K.; Tester, J. W. Oxidation Kinetics of Carbon Monoxide in Supercritical Water. Energy Fuels 1987, 1, 417. (6) Helling, R. K.; Tester, J. W. Oxidation of Simple Compounds and Mixtures in Supercritical Water: Carbon Monoxide, Ammonia, and Ethanol. Environ. Sci. Tech. 1988, 22, 1319. (7) Holgate, R. H.; Webley, P. A.; Tester, J. W.; Helling, R. K. Carbon Monoxide Oxidation in Supercritical Water: The Effects of Heat Transfer and the Water-Gas Shift Reaction on Observed Kinetics. Energy Fuels 1992, 6, 586. (8) Hirth, T.; Franck, E. U. Oxidation and Hydrothermolysis in Supercritical Water at High Pressures. Ber. Bunsen-Ges. Phys. Chem. 1993, 97, 1091. (9) Rice, S. F.; Steeper, R. R.; Aiken, J. D. Water Density Effects on Homogeneous Water-Gas Shift Kinetics. J. Phys. Chem. A 1998, 102, 2673. (10) Sato, T.; Kurosawa, S.; Adschiri, T.; Arai, K. Kinetics of the Water-Gas Shift Reaction in Supercritical Water. Kagaku Kogaku Ronbushu 1999, 25, 993. (11) Brill, T. B. Geothermal Vents and Chemical Processing: The Infrared Spectroscopy of Hydrothermal Reactions. J. Phys. Chem. A 2000, 104, 4343. (12) Kieke, M. L.; Schoppelrei, J. W.; Brill, T. B. Spectroscopy of Hydrothermal Reactions I: The CO2-H2O System and Kinetics of Urea Decomposition in an FTIR Spectroscopy Flow Reactor Cell Operable to 725 K and 335 bar. J. Phys. Chem. 1996, 100, 7455. (13) Schoppelrei, J. W.; Kieke, M. L.; Wang, X.; Klein, M. T.; Brill, T. B. Spectroscopy of Hydrothermal Reactions IV. Kinetics of Urea and Guanidinium Nitrate at 200-300 °C in a Diamond Cell, Infrared Spectroscopy Flow Reactor. J. Phys. Chem. 1996, 100, 14343. (14) Wilhelm, E.; Battino, R.; Wilcock, R. J. Low-Pressure Solubility of Gases in Liquid Water. Chem. Rev. 1977, 77, 219. (15) Brill, T. B.; Miksa, D.; Gunawardena, N. R. Spectroscopy Studies of Hydrothermal Reactions of Organic-Inorganic Mixtures. In Proceedings of the Joint Sixth International Symposium on Hydrothermal Reactions and Fourth International Conference on Solvothermal Reactions; Kochi University: Kochi, Japan, 2000. (16) Maiella, P. G.; Brill, T. B. Spectroscopy of Hydrothermal Reactions 11. Infrared Absorptivity of CO2 and N2O in Water at Elevated Temperature and Pressure. Appl. Spectrosc. 1999, 53, 351. (17) Cvetanovic, R. J.; Singleton, D. L. Comment on the Evaluation of the Arrhenius Parameters by the Least Squares Methodology. Int. J. Chem. Kinet. 1977, 9, 481. (18) Elliott, D. C.; Haller, R. T.; Sealock, L. J., Jr. Aqueous Catalyst Systems for the Water-Gas Shift Reaction. 2. Mechanism of Basic Catalysis. Ind. Eng. Chem. Prod. Res. Dev. 1983, 22, 431. (19) Ruelle, P.; Kesselring, U. W.; Nam-Tran, H. Ab Initio Quantum-Chemical Study of the Unimolecular Pyrolysis of Formic Acid. J. Am. Chem. Soc. 1986, 108, 371. (20) Akiya, N.; Savage, P. E. Role of Water in Formic Acid Decomposition. AIChE J. 1998, 44, 405. (21) Yoneda, K.; Honda, Y.; Momlyana, N.; Abe, R. Water Gas Conversion under Elevated Pressure III. Behavior of Water Gas in Aqueous K2CO3 Solution under Elevated Pressure and Temperature. J. Chem. Soc. Jpn. 1943, 46, 554. (22) Zielke, C. W.; Rosenhoover, W. A.; Gorin, E. Direct Zinc Chloride Hydrocracking of Subbituminous Coal. Regeneration of Spent Smelt. Prepr., Div. Fuel Chem., Am. Chem. Soc. 1976, 21, 163. (23) Blake, P. G.; Hinshelwood, C. Homogeneous Decomposition Reactions of Gaseous Formic Acid. Proc. R. Soc. London A 1960, 255, 444. (24) Blake, P. G.; Davies, H. H.; Jackson, G. E. Dehydration Mechanisms in the Thermal Decomposition of Gaseous Formic Acid. J. Chem. Soc. B 1971, 1923.

Ind. Eng. Chem. Res., Vol. 40, No. 14, 2001 3103 (25) Saito, K.; Kakumoto, T.; Kuroda, H.; Torii, S.; Imamura, A. Thermal Unimolecular Decomposition of Formic Acid. J. Chem. Phys. 1984, 80, 4989. (26) Goddard, J. D.; Yamaguchi, Y.; Schaefer, H. F., III. The Decarboxylation and Dehydration Reactions of Monomeric Formic Acid. J. Chem. Phys. 1992, 96, 1158. (27) Francisco, J. F. A Comprehensive Theoretical Examination of Primary Dissociation Pathways of Formic Acid. J. Chem. Phys. 1992, 96, 1167. (28) Trillo, J. M.; Munuera, G.; Criado, J. M. Catalytic Decomposition of Formic Acid on Metal Oxides. Catal. Rev. 1972, 7, 51. (29) McCarty, J.; Falconer, J.; Madix, R. J. Decomposition of Formic Acid on Ni(110). J. Catal. 1973, 30, 235. (30) Falconer, J.; Madix, R. J. Kinetics and Mechanism of the Autocatalytic Decomposition of HCO2H on Clean Ni(110). Surf. Sci. 1974, 79, 473. (31) Benziger, J. B.; Madix, R. J. The Decomposition of Formic Acid on Ni(100). Surf. Sci. 1979, 79, 394. (32) Barteau, M. A.; Bowker, M.; Madix, R. J. Acid-Base Reactions on Solid Surfaces: The Reactions of HCO2H, H2CO, HCO2CH3 with O2 on Ag(110). Surf. Sci. 1980, 94, 303. (33) Benziger, J. B.; Schoofs, G. R. Influence of Adsorbate Interactions on Heterogeneous Reaction Kinetics. Formic Acid Decomposition on Nickel. J. Phys. Chem. 1984, 88, 4439.

(34) Schustorovich, E.; Bell, A. T. An Analysis of Formic Acid Decomposition on Metal Surfaces by the Bond-Order-Conservation-Morse-Potential Approach. Surf. Sci. 1989, 222, 371. (35) Maiella, P. G.; Brill, T. B. Spectroscopy of Hydrothermal Reactions 10. Evidence of Wall Effects in the Decarboxylation Kinetics of 1.00 m HCO2X (X ) H, Na) at 280-330 °C and 275 Bar. J. Phys. Chem. A 1998, 102, 5886. (36) Mond, L.; Langer, C.; Quincke, F. Action of Carbon Monoxide on Nickel. J. Chem. Soc. London 1890, 57, 749. (37) Tepe, R. K.; Vassallo, D.; Jacksier, T. Synthesis of Iron and Nickel Carbonyl ass Calibration Materials for Spectroscopic Systems. Spectrochim. Acta Part B 2000, 55, 165. (38) Cody, G. D.; Boctor, N. Z.; Filley, T. R.; Hazen, R. M.; Scott, J. H.; Sharma, A.; Yoder, H. S., Jr. Primordial Carbonylated IronSulfur Compounds and the Synthesis of Pyruvate. Science 2000, 289, 1337.

Received for review January 24, 2001 Revised manuscript received April 18, 2001 Accepted May 8, 2001 IE010076J