Ind. Eng. Chem. Res. 2008, 47, 8119–8125
8119
Kinetics and Mechanism of Decomposition of Hydrogen Peroxide over Pd/SiO2 Catalyst Yury Voloshin,* James Manganaro, and Adeniyi Lawal Department of Chemical, Biomedical, and Materials Engineering, SteVens Institute of Technology, New Jersey Center for Microchemical Systems, 1 Castle Point, Hoboken, New Jersey 07030
Kinetics of hydrogen peroxide decomposition over Pd/SiO2 catalyst was investigated as part of a larger project to determine the overall kinetics of hydrogen peroxide formation by Pd/SiO2 catalyzed direct combination of hydrogen and oxygen. A Pd oxidative/reductive cycle mechanism for the decomposition of H2O2 over Pd/ SiO2 was proposed. A Rideal-Eley rate expression based on this mechanism was shown to accurately correlate experimental data over a wide range of H2O2 concentration and temperature. This rate equation was shown to also apply in the presence of stabilizing concentrations of H2SO4. 1. Introduction In recent years, synthesis of hydrogen peroxide by direct combination of hydrogen and oxygen over a heterogeneous catalyst has been generating an increasing amount of interest among researchers. Direct combination (DC) is a highly desirable alternative to the currently used anthraquinone autoxidation method of producing hydrogen peroxide, which has been the dominant method of producing hydrogen peroxide since the middle of the 20th century.1 DC is the most direct method of producing H2O2 and has sometimes been called the “dream process”.2 A large number of variations of this process have been investigated in either laboratory scale or pilot-plant scale since 1914;3-12 however, commercial production of hydrogen peroxide by the DC method has not been attempted until very recently.13 Two main difficulties must be overcome before producing hydrogen peroxide by DC in marketable amounts. First, safety of the process must be ensured. A gaseous mixture of hydrogen and air is flammable when the concentration of hydrogen is between about 5 and 76%. This is why most of the patented DC processes which use macroreactors are designed to operate with hydrogen concentrations below the explosive limit. The low hydrogen concentrations make it necessary to use high total pressures (in the order of 10-20 MPa), which makes the process expensive to operate. The DC process can be made inherently safe by using one of the two relevant processes described in the literature: a membrane reactor and a microreactor. In a membrane reactor,14 hydrogen diffuses through a membrane into an oxygen-saturated liquid solvent. In a microreactor,15,16 the width of microchannels is smaller than the quenching distance of hydrogen and oxygen radicals, such that an explosive chain reaction is practically impossible. The publications dealing with inherent safety of microreactors15,16 focus on direct water formation; however, the same conclusions apply to hydrogen peroxide formation. Microreactors enable the use of high hydrogen concentrations above the explosive limit, thus eliminating the need to operate the DC process at high pressures and making the process significantly less expensive to operate. Moreover, microreactors can be used for production of hydrogen peroxide at the end user’s site, which would eliminate most of the transportation expenses. The use of different types of microreactors for production of up to 1.2% H2O2 (w/w) has
been demonstrated by the authors of this work and by other researchers.17-19 The second difficulty of the DC process is the low selectivity for H2O2. The DC process involves three reactions that decrease selectivity, as shown in Scheme 1. The reaction of hydrogen and oxygen over a catalyst can yield water as well as hydrogen peroxide. Once H2O2 is formed, it can be consumed by either reduction or decomposition. Thus, the rates of reactions 2-4 must be minimized to increase the selectivity for hydrogen peroxide. Starting in 1961,20 a large amount of research2,21,22 showed certain halides and acids to be strong inhibitors of reactions 2-4. Halide anions, particularly Br- and Cl-, act as a general catalyst poison to slow down the rates of all reactions in Scheme 1. However, reaction 1 is less affected by the halides than reactions 2-4, so that the halides significantly increase the selectivity for hydrogen peroxide.2 This effect applies to halides dissolved in the reaction medium, as well as to those incorporated into the catalyst support.23 H2O2 decomposition (reaction 4) is also inhibited by the presence of acids in the reaction medium, particularly H2SO4, H3PO4, and HCl.22 The inhibition effect of acids is attributed to the formation of a bond between protons and H2O2, which would give H2O2 a partial positive charge and prevent it from adsorbing onto the catalyst.2 Acids and halides promote the formation of H2O2 when either is present in the reaction medium. However, when both an acid and a halide are present, the promoting effect is greater than the individual effect of either compound.2 This is due to the greater efficacy of halides in suppressing the decomposition of H2O2 in the presence of acid: the suppressing effect of halides was greater by approximately a factor of 3 in acidic solution than in pure water.22 It has been proposed2 that the protons enhance the effect of halides by decreasing the pH below the Scheme 1. Reactions Involved in the Formation of Hydrogen Peroxide
* To whom correspondence should be addressed. Tel.: (917) 3498135. E-mail:
[email protected]. 10.1021/ie8000452 CCC: $40.75 2008 American Chemical Society Published on Web 10/08/2008
8120 Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008
Figure 1. Experimental setup for decomposition of hydrogen peroxide.
isoelectric point of the catalyst support, so that the catalyst surface becomes positively charged and the halide anions are much more likely to be adsorbed onto the catalyst, where they inhibit the H2O2 consumption reactions. Thus, the promoting effect of acid on H2O2 productivity has been well established. However, hydrogen peroxide in an acidic solution is not desirable for a number of applications where H2O2 is currently used, such as soil remediation and cloth bleaching.1 In such cases, the DC process must be operated using a solvent that is either acid-free or with a very low acidity. Any attempt to design or model a microreactor for production of hydrogen peroxide requires knowledge of reactions 1-4 in Scheme 1. In previous publications,18,19 we determined the kinetics of reactions 1 and 3 by running experiments under conditions where the decomposition reaction, r4, was completely suppressed by 1% H2SO4 (w/w) and 10 ppm NaBr in the reaction medium. The synthesis reaction18 was isolated by running the process in a differential reactor, where the conversion of reactants is very small, so that the selectivity for H2O2 was 100%. The reduction reaction, r3, was isolated19 by using a mixture of hydrogen and nitrogen as the gas phase and a solution of hydrogen peroxide in the liquid phase. In the present work, our goal was to propose a plausible mechanism for the decomposition of H2O2 over Pd/SiO2 catalyst and from that obtain a broadly applicable rate expression for this reaction. This provides us with an expression for r4, so that the overall rate of formation of H2O2 may be calculated from eq 1: rH2O2,ov ) r1 - r3 - r4
(1)
Direct water formation, r2, is neglected because the selectivity for H2O2 at low conversions of reactants was 100%, which shows that H2O2 is most likely the sole initial product. The possibility of direct water formation at high conversions of reactants is beyond the scope of this article and will be discussed in a future publication on the overall reactor model for the DC process. Decomposition of hydrogen peroxide is a widely studied reaction, and a number of studies were done on decomposition catalyzed by palladium on a variety of supports, for example.24-28 In the literature, the ranges of H2O2 concentration covered were typically less than 1 mol/L and the heterogeneously catalyzed decomposition reaction was always found to follow first-order kinetics with respect to H2O2. A first-order rate expression would not adequately describe the data at higher concentrations of H2O2 because mechanisms of heterogeneously catalyzed reactions are typical of Langmuir-Hinshelwood type.29 At low concentrations, Langmuir-Hinshelwood rate expressions can usually be
simplified to first-order rate expressions. Since our goal was to model a microreactor for production of up to 1.5 mol/L H2O2 (5% w/w), we performed kinetic experiments at sufficiently high concentrations of H2O2 to find a Langmuir-Hinshelwood rate expression for this reaction. In addition, we aim to quantify the effect of sulfuric acid on the rate of decomposition by comparing the Langmuir-Hinshelwood kinetic constants at two different acid concentrations. A study by Choudhary and Gaikwad28 looked at first-order kinetics for H2O2 decomposition at different concentrations of H2SO4 on Pd/C catalyst. The researchers found that the activation energy decreased from 17.65 to 17 kJ/mol and the pre-exponential factor decreased from 107.2 to 61.7 min-1 as the acid concentration was varied from 0 to 10 mol/L (approximately 98% w/w). However, no halides were used in the reaction medium, which makes their result not comparable to the present work. As described above, the combined inhibiting effect of sulfuric acid and halide anions is much greater than the individual effect of each. In this work, all decomposition experiments were carried out in the presence of 10 ppm NaBr to obtain rate expressions that are compatible with the previously obtained kinetics of synthesis and reduction of hydrogen peroxide.18,19 2. Experimental Methods Materials. The catalyst used for all experiments was 2 wt % Pd on SiO2, prepared by the sol-gel method with PdCl2 as the source of palladium. Surface area of the catalyst was 603 m2/g (multipoint BET technique, using Quantochrome Instruments Autosorb-1) and a dispersion of 18.1% (FEG-TEM, model CM20, Philips). The catalyst was ground and sieved to obtain particles with a diameter in the range of 75-150 µm. This is the particle size range used in all experiments in this work, unless specified otherwise. Additional details on catalyst preparation and testing are given in earlier publications,18,30 and XRD and SEM characterization of the catalyst is described in ref 31. The liquid phase consisted of deionized water with a specified concentration of H2O2 (stabilizer-free, ACS reagent grade, Fisher Scientific, 30%), H2SO4 at a specified concentration (ACS reagent grade, 95-98%, Aldrich), and 10 ppm NaBr (Aldrich). The gas phase consisted of air (Extra Dry, Praxair). Experimental Setup. The experimental setup shown in Figure 1 is nearly identical to that used for kinetic experiments on H2O2 synthesis and reduction.18,19 Even though the reaction takes place in the liquid phase, a gas phase was added to simulate the conditions present during H2O2 formation. The main effect of the gas phase was to influence the residence time in the reactor. When the decomposition reaction was run without
Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008 8121 Table 1. Mass Balance of Oxygen and Hydrogen Peroxide liquid H2O2 consumed, O2 predicted, O2 measured, N2 flow rate, flow rate, mol/min mol/min mol/min mol/min mL/min 9.38 × 10-4 1.83 × 10-3
0.05 0.1
5.28 × 10-5 6.35 × 10-5
2.65 × 10-5 3.18 × 10-5
2.90 × 10-5 3.73 × 10-5
any gas phase, the reaction rate was proportional to the increased residence time. Flow rate of the liquid was controlled by an HPLC pump (Laboratory Alliance Series III), and the solution was combined with air in a PTFE micromixer (Upchurch). The flow rate of gas was controlled by mass flow controllers (Porter model 201). The microreactor consisted of 1/16′′ 316 L stainless steel tubing with 765 µm i.d., which was packed with catalyst. Filters of polypropylene wool were placed at the ends of the reactor to retain the catalyst. Concentration of H2O2 in the liquid phase was measured by titration with potassium permanganate, and composition of the gas phase was measured by a Shimadzu GC-14B gas chromatograph with a Mole Sieve 5A column at 35 °C and argon as carrier gas. Pressure drop across the reactor varied between approximately 2 and 7% of the inlet pressure (depending on the total flow rate), while the average of pressures at the inlet and outlet of the reactor was maintained at 2.16 MPa for all experiments. The gas/liquid ratio was held constant at 440 (v/v, at standard conditions). The observed reaction rates were calculated as
(
)
(CH2O2,in - CH2O2,out)Fliq 34 g of H2O2 (2) W mol The reactor was operated without any catalyst at 50 °C before all the experiments reported here to ensure that any H2O2 decomposition activity was due to the catalyst alone. If the decomposition rate without catalyst was significant, then the experimental system was treated with 35% HNO3 for about 30 min to reduce the catalytic activity of the material of construction of the reactor and the supporting tubing. All kinetic experiments were done in a differential reactor mode, so that the conversion of hydrogen peroxide was kept below 10 wt %. The reported reaction rates were measured after the reactor performance had been stabilized for about two hours. When the conversion was higher than 10%, then the average reaction rate was reported as space-time yield (STY). RH2O2 )
3. Results and Discussion Mass Balance. Before carrying out the kinetic experiments for H2O2 decomposition, where air was used as the gas phase, several runs were done with nitrogen as the gas phase. This was done to be able to compare the amount of oxygen evolved to the amount of peroxide decomposed to confirm that the mass balance was satisfied. The results are shown in Table 1. The measured amount of oxygen evolved was determined using the concentration of oxygen in the gas phase at the exit of the reactor and the flow rate of the gas phase. The predicted amount of oxygen was calculated on the basis of the amount of decomposed H2O2. The errors, or the differences between the measured and predicted amounts of oxygen for the two runs, were 9 and 17%. The larger error can be attributed to the difficulty in accurately measuring small concentrations of oxygen in the gas phase. Flow Regime. A description of the flow regime in the packed microreactor is important to understand the mass transfer process during this reaction. The flow regime was observed during flow visualization experiments, which were carried out by replacing the reactor with a transparent glass capillary with a square cross
Figure 2. Effect of flow velocity on the rate of H2O2 decomposition. Reaction conditions: 300 psig, 50 °C; gas phase: air; liquid phase: deionized water with 2% H2O2 and 10 ppm NaBr, 2-6 mg of catalyst.
section measuring 500 × 500 µm2.18 In empty capillaries, and at the flow rates used in the work, the observed flow regime is usually Taylor flow.32 This was confirmed for our system. When the tube was packed with glass beads of the same size range as the catalyst, the flow still consisted of distinct gas and liquid slugs, but it was no longer typical Taylor flow because of the continuous breaking of liquid slug boundaries by the glass beads and long mistlike tails trailing behind liquid slugs. This flow regime was also observed by Wada and co-workers in a post microreactor,33 who referred to it as slug flow. Breaking up of the liquid slugs made the flow more chaotic and facilitated mass transfer between the liquid phase and the catalyst particles.33 External Mass Transfer Limitations. In the case of H2O2 decomposition, external mass transfer steps involve both gas-liquid and liquid-solid mass transfer. The liquid-solid mass transfer refers to the diffusion of H2O2, and the diffusion of water, to and from the catalyst surface, respectively. The gas-liquid mass transfer takes place during the diffusion of generated O2 from the catalyst surface to the liquid phase. Since both the gas-liquid and the liquid-solid mass transfer rates are affected by flow velocity,32 we measured the rate of decomposition at various flow velocities, while keeping the residence time constant. Figure 2 shows that the flow velocity has no effect on the rate of decomposition. Thus, the decomposition reaction is not limited by external mass transfer. Internal Mass Transfer Limitations. The significance of an internal mass transfer limitation was estimated by calculating a first-order Thiele modulus according to eq 3:34 φexp )
[ ]
dp FprH2O2 ′ 6 DeCH2O2
0.5
(3)
where Fp is the density of catalyst particles, rH2O2′ is the reaction rate, and De is the effective diffusivity, which was estimated using the relationship De ) Dε/τ, where D is the binary diffusivity of H2O2 in water at 40 °C (2 × 10-5 cm2/s)35 and where the typical values of porosity (0.4-0.6), tortuosity (2-8), and the constriction factor (0.7-0.8)34 were used. The calculation was done for conditions that give the highest reaction rate achieved in these experiments (170 g/g of Pd/h). The worstcase Thiele modulus was estimated to be 0.15, which corresponds to an effectiveness factor of unity.34 Since the effectiveness factor is a ratio of the observed reaction rate to the reaction rate at the external surface of the catalyst particle, this implies that the reaction is not limited by internal mass transfer. The data shown in Figure 3, in the presence of 10 ppm of NaBr, confirm this conclusion. In this case, the rate of decomposition did not depend on the size of catalyst particles, which shows experimentally that the rate of H2O2 decomposition is not limited by internal diffusion. However, when the experiments were run without NaBr in the reaction medium (Figure 3), the rate of
8122 Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008
Figure 3. Effect of catalyst particle size on the rate of H2O2 decomposition. Reaction conditions: 300 psig, 50 °C; gas phase: air; liquid phase: deionized water with 6% (w/w) H2O2 and NaBr at specified concentration, 6 mg of catalyst.
decomposition increased with decreasing size of the catalyst particles. This shows that, when the rates of decomposition were increased by running the experiments without NaBr, the reaction becomes limitied by internal diffusion. This finding, however, is not directly relevant to the experiments in this work and is shown here only for completeness, because all the experiments reported here were done with 10 ppm NaBr in the reaction medium. In addition to the lack of an internal mass transfer limitation, the results in Figure 3 show that there was no channeling, or large gaps between particles, in the catalyst bed. Channeling becomes likely when drxr/dp < 8.36 In our case, drxr/dp is between 10 and 5. The lack of a particle size effect in the presence of 10 ppm NaBr in Figure 3 indicates an absence of channeling in the catalyst bed. Mechanism of H2O2 Decomposition. In this section, we discuss possible mechanisms and associated rate expressions for the catalytic decomposition of H2O2. In selecting an acceptable rate expression, the following criteria were used: 1. The concentration of oxygen must not be a term in the rate expression because our data showed no significant difference between the rates of decomposition with either air or nitrogen as the gas phase. 2. Reaction rate must approach zero as concentration of hydrogen peroxide approaches zero. 3. Values of rate and equilibrium constants must be positive. 4. R2 > 0.95.
5. Confidence intervals of the kinetic parameters must be smaller than the values of the parameters. Mechanisms of palladium-catalyzed H2O2 decomposition found in the literature proceed either by a radical pathway or through a complex of peroxide with the palladium compound.23 The only mechanism found in the literature for H2O2 decomposition on Pd/SiO2 catalyst was proposed by Choudhary and Samanta,2 and it involves a dissociation of H2O2 into two OH radicals and a subsequent surface reaction of the radicals to form water. A modified version of this mechanism is shown in Table 2. The mechanism was modified by adding step 3, which is necessary to account for the formation of oxygen and which was missing in Choudhary and Samanta.2 Each step is shown in Table 2 together with the corresponding rate equation, which was derived assuming that the given step is rate-determining and the other steps are in pseudo-equilibrium. The cases where steps 1 and 2 are rate-determining were rejected because, in those cases, the reaction rate depends on the concentration of oxygen. The remaining case where step 3 is rate-determining is a potential rate expression for this reaction. In addition to the radical mechanism described above, we propose a second mechanism that consists of cyclic oxidation and reduction of palladium by H2O2, as shown in eqs 4 and 5: Pd + H2O2 f PdO + H2O PdO + H2O2 f Pd + H2O + O2
∆G0 ) -202 kJ/mol
(4)
∆G0 ) -31.8 kJ/mol (5)
This mechanism is classified as Rideal-Eley since H2O2 molecules from the bulk solution directly react with surface molecules. The negative values of ∆G0 show eqs 4 and 5 to be plausible. This mechanism is supported by the finding by Samanta and Choudhary,21,22 where the rates of H2O2 decomposition were higher on Pd0 than on PdO by factors of 3-10, depending on reaction conditions. This is consistent with eqs 4 and 5 because converting Pd0 to PdO stops eq 4 and the remaining eq 5 would most likely be slower because of the smaller change of Gibbs free energy. Table 2 shows two possible rate expressions derived using eqs 4 and 5 corresponding to each of the mechanism steps as the rate-determining step. The rate expression where step 1 is rate-determining was rejected because one of the equation parameters is the concentration of oxygen. This leaves the rate expression based on step 2 as the RDS as the second potential rate expression. Kinetic Experiments in Acid-Free Reaction Medium. The rates of H2O2 decomposition over Pd/SiO2 were measured at different H2O2 concentrations at 30, 40, and 50 °C in differential reactor mode, where the conversion was kept below 10%. Figure 4 shows the experimental kinetic curves. Kinetic parameters were obtained by nonlinear regression of the data in Figure 4
Table 2. Proposed Mechanisms of H2O2 Decomposition with Rate Expressions that Correspond to Each Mechanism Step as Rate-Limitinga Radical Mechanism of Choudhary and Samanta2 step
limiting rate expression
rate expression number
1) 2H2O2 + 4* f 4*OH 2) 4*OH f 2H2O + 2*O + 2* 3) 2*O f O2 + 2*
r ) k1CH2O22/{1 + (K3CO2)1/2 + CH2O1/2[(K3/K2)CO2]1/4}4 r ) k2/[(1/K11/4CH2O21/2) + (1/K11/4)(CO2/K3CH2O2)1/2 + 1]4 r ) k3/{1 + [CH2O/(K1K2)1/2CH2O2] + [CH2O/K11/4K21/2CH2O21/2]}2
1 2 3
Cyclic Pd Oxidation/Reduction Mechanism
a
site.
step
limiting rate expression
1) * + H2O2 f *O + H2O 2) *O + H2O2 f * + H2O + O2
r ) k1CH2O2/[1 + (CH2OCO2/K2CH2O2)] r ) k2CH2O2/[1 + (CH2O/K1CH2O2)]
rate expression number 4 5
Parameters ki and Ki are the rate constants and equilibrium coefficients (respectively) that correspond to step i of a mechanism. * ) Pd catalytic
Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008 8123
of rate expression 5 (Table 3). Thus, rate expression 3 is rejected and rate expression 5 remains as the most likely rate expression for this reaction. The reaction rates predicted by rate expression 5 of Table 2 are shown in Figure 4 together with the experimental data. Arrhenius parameters corresponding to this rate expression were calculated by substituting the isothermal kinetic constants in rate expression 5 with the corresponding Arrhenius expressions (eqs 6 and 7) and regressing the data at all three temperatures simultaneously.
( )
k2 ) Ak2 exp Figure 4. Effect of H2O2 concentration on rate of H2O2 decomposition. Reaction conditions: 300 psig, 5-6 mg of catalyst; gas phase: air; liquid phase: deionized water with 10 ppm NaBr. Dotted lines are 95% confidence intervals. Table 3. Isothermal Kinetic Constants for the Two Potential Rate Expressionsa
k
K1 K2
T, °C
rate expression 3
rate expression 5
30 40 50 30 40 50 30 40 50
39.71 ( 3.91 91.49 ( 3.72 265.50 ( 29.23 709.47 ( 3860.08 294.40 ( 70.72 461.40 ( 616.50 399.48 ( 1390.27 592.08 ( 94.27 290.24 ( 252.35
37.72 ( 0.76 74.51 ( 0.82 178.93 ( 0.62 66.41 ( 2.45 84.99 ( 2.05 101.99 ( 0.84
a Units of parameters: k, g/g of Pd/h; K1, (L/mol)2 for rate expression 3, L/mol for rate expression 5; K2, (L/mol)2.
Figure 5. Comparison of rate expression 5 predictions with integral reactor data for H2O2 decomposition in acid-free reaction medium.
using Polymath software and Levenberg-Marquardt algorithm. Isothermal values of the kinetic parameters for the two potential rate expressions, expressions 3 and 5 of Table 2, were calculated by regressing the data for each temperature separately and are shown in Table 3. Both rate expressions have similar values of R2, which is a measure of how well the model predicts variations in experimental data. However, the kinetic parameters of rate expression 3 have much larger confidence intervals than those
(
K1 ) AK1 exp -
Ea RT
∆HK1
(6)
)
(7) RT The results are shown in Table 4. In another set of experiments, the kinetics were confirmed by comparing predictions with experimental integral reactor data at high conversions in Figure 5, which are acceptably close. Kinetic Experiments in 0.05% (w/w) H2SO4. Kinetic experiments of the same type as described above were repeated with 0.05% H2SO4 in the reaction medium. Figure 6 shows the experimental points together with the predicted reaction rates. The kinetic parameters with Arrhenius constants can be seen in Tables 5 and 6. A comparison of the Arrhenius constants at each acid concentration shows a strong effect of the acid on pre-exponential factors and a weaker effect on the activation energy and the reaction enthalpy of the pseudo-equilibrium step. The pre-exponential factor for the kinetic parameter k increases by a factor of 10 in the acidic solution, while that of the equilibrium constant K decreases by a factor of about 20. The isothermal kinetic constants k (Table 6) decrease by a factor of approximately seven in acidic solution, despite the increase in the pre-exponential factor, because of the increase in activation energy. As expected, a comparison with the previously described results of Choudhary and Gaikwad28 shows that the inhibiting effect of acid is much stronger in the presence of NaBr. Effect of Acid Concentration on Kinetic Parameters. The effect of acid concentration on the rate of decomposition can be seen in Figure 7. After obtaining the kinetics of decomposition at two different acid concentrations, we extrapolated the kinetic constants to the range of H2SO4 concentrations shown in Figure 7. This was possible by the fact that the decomposition reaction is completely suppressed in 1% H2SO4 at temperatures up to 50 °C.18,19 Under these conditions, the value of k must be nearly zero. A logarithmic plot of the isothermal k values against H2SO4 concentration can be constructed using the three points at 0, 0.05, and 1% H2SO4 (Figure 8). At 1% H2SO4, 10-12 was used as the value of k, which is nearly zero. The result is a correlation that enables us to predict k at any acid concentration between 0 and 1% and at any temperature between 30 and 50 °C. The value of k at 1% H2SO4 was chosen as 10-12 because this is the maximum value that allowed us to predict the k values at 0 and 0.05% H2SO4 within 90% of the experimental values. Similar curves with R2 > 0.95 were also obtained for 30 and 50 °C. As far as the pseudo-equilibrium constant K is concerned,
Table 4. Arrhenius Constants of Rate Expression 5 in Acid-Free Solution AK1 (mol/L)
Ak2 (g of H2O2/g of Pd/h) 4.73 × 10
13
( 2.42 × 10
13
1.01 × 10 ( 1.33 × 10 4
4
Ea (kJ/mol)
∆HK1 (kJ/mol)
70.9 ( 1.36
11.8 ( 0.35
Table 5. Arrhenius Constants of Rate Expression 5 with 0.05% H2SO4 in the Reaction Medium Ak2 (g of H2O2/g of Pd/h) 3.03 × 10
14
( 3.31 × 10
13
AK1 (mol/L)
Ea (kJ/mol)
∆HK1 (kJ/mol)
3942.85 ( 686.40
80.5 ( 0.29
12.7 ( 0.46
8124 Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008
the experimental rates of decomposition at various acid concentrations to the predicted rates. The results in Figure 7 show that the predicted results are acceptably close to the experimental data. This figure also shows that a concentration of approximately 0.1% H2SO4 is sufficient to significantly suppress decomposition. The exponential relationship, however, is still tentative because of the small number of H2SO4 concentration points used, and further research is needed to confirm its validity. 4. Conclusions
Figure 6. Effect of H2O2 concentration on rate of H2O2 decomposition. Reaction conditions: 300 psig, 15-20 mg of catalyst; gas phase: air; liquid phase: deionized water with 0.05% H2SO4, 10 ppm NaBr, and a specified concentration of H2O2.
Figure 7. Effect of sulfuric acid concentration on the kinetic constant k of rate expression 5 at 42 °C.
Figure 8. Comparison of experimental and predicted decomposition rates at different concentrations of sulfuric acid. Reaction conditions: 300 psig, 50 °C, 10-20 mg of catalyst; gas phase: air; liquid phase: deionized water with 0.32 mol/L H2O2, 10 ppm NaBr, and a specified concentration of H2SO4. Table 6. Isothermal Kinetic Constants of Rate Expression 5 at Different Acid Concentrations
k2 (g of H2O2/g of Pd/h) K1 (mol/L)
T, °C
acid-free
0.05% H2SO4
30 40 50 30 40 50
37.72 74.51 178.93 66.41 84.99 102.00
5.42 10.03 29.00 29.98 33.69 34.51
we do not have a value of K at a third acid concentration. A preliminary assumption was made of isothermal K’s that follow the same logarithmic relationship as k. This assumption enabled us to obtain a correlation for isothermal K’s at different acid concentrations by constructing a curve similar to Figure 8 with only two K values at 0 and 0.05% H2SO4. Similar curves for isothermal K’s with high R2 values were also obtained for 30 and 50 °C. The exponential relationship between the H2SO4 concentration and kinetic constants was tested by comparing
Knowledge of the kinetics of hydrogen peroxide formation by direct combination of hydrogen and oxygen is necessary for rational design, simulation, and optimization of a microreactor for the DC process. Since the DC process involves parallel and series reactions that produce and consume H2O2, the rate expressions of all relevant reactions are required to be able to predict the overall rate of hydrogen peroxide formation. The three relevant reactions are synthesis, reduction by hydrogen, and decomposition of hydrogen peroxide. In this work, we focused on catalytic decomposition of H2O2. We proposed a cyclic Pd oxidation/reduction mechanism for the decomposition reaction (eqs 4 and 5) and derived a rate expression based on this mechanism (rate expression 5 of Table 2). Values for the kinetic parameters of rate expression 5 were obtained by nonlinear regression of experimental data (Table 3). The proposed mechanism is supported by a close agreement between the experimental and predicted reaction rates in both acid-free and acidic solutions, as well as the observations of Choudhary et al.21,22 described above. An unambiguous proof of the mechanism, however, requires detailed surface studies of the catalyst, which may be done as part of future research. The rates of decomposition in 0.05% (w/w) H2SO4 were lower than those in acid-free solution by approximately 1 order of magnitude, when both sets of experiments were done with 10 ppm NaBr in the reaction medium. This difference was much greater than the difference between rates of decomposition in acid-free and acidic halide-free solutions obtained from the literature. The difference was attributed to the presence of NaBr. Isothermal kinetic parameters were found to decrease in an inverse exponential relationship with respect to the concentration of H2SO4. This relationship was verified by comparing the experimental and predicted rates of decomposition between 0 and 0.2 wt % H2SO4. However, further research is needed to confirm this relationship because it was determined using only three H2SO4 concentration points. The inverse exponential relationship is purely empirical, and further study is needed to relate it to the mechanism of interactions between the acid, bromide, hydrogen peroxide, and catalyst. It was found that a H2SO4 concentration of about 0.1% was sufficient to significantly suppress H2O2 decomposition over Pd/SiO2 catalyst. Acknowledgment We thank the U.S. Department of Energy (Industrial Technologies Program) for financial support under Contract DEFC36-02ID14427 and FMC Corporation for their technical contribution to the program. Y.V. thanks the Robert C. Stanley Graduate Fellowship Program. In addition, we are grateful to Dr. Dalbir Sethi of FMC for his help with the mechanisms of H2O2 decomposition. Nomenclature CA ) concentration of species A in liquid, mol/L dp ) diameter of catalyst particles, m
Ind. Eng. Chem. Res., Vol. 47, No. 21, 2008 8125 drxr ) internal diameter of reactor, m D ) binary diffusivity, cm2/s De ) effective molecular diffusivity, cm2/s Ea ) activation energy, kJ/mol Fliq ) volumetric flow rate of the liquid phase, L/h k ) kinetic rate constant, g of H2O2/g of Pd/h K ) equilibrium constant L ) length of reactor, m RH2O2 ) observed reaction rate, g of H2O2 decomposed/g of Pd/h rH2O2 ) intrinsic reaction rate, g of H2O2 decomposed/g of Pd/h rH2O2′ ) intrinsic reaction rate, mol H2O2 decomposed/g of catalyst/ min W ) mass of active catalytic metal in the reactor, g of Pd w ) catalyst mass per unit volume of reactor, g/cm3 Greek Symbols Fp ) density of catalyst particles, g/m3, Fp ) w/ε ∆H ) enthalpy of reaction, kJ/mol φexp ) experimental Thiele modulus, φexp ) [dp/6][FprH2O2′/ DeCH2O2]0.5 ε ) porosity, dimensionless τ ) tortuosity, dimensionless δ ) constriction factor, dimensionless
Literature Cited (1) Hydrogen Peroxide. In Kirk-Othmer Encyclopedia of Chemical Technology, 5th ed.; Seidel, A., Ed.; Wiley & Sons: Hoboken, NJ, 2005; Vol. 13. (2) Choudhary, V. R.; Samanta, C. Role of chloride or bromide anions and protons for promoting the selective oxidation of H2 by O2 to H2O2 over supported Pd catalysts in an aqueous medium. J. Catal. 2006, 238, 28. (3) Hooper, G. W. Catalytic production of hydrogen peroxide from its elements. U.S. Patent 3,336,112, 1967. (4) Izumi, Y.; Miyazaki, H.; Kawahara, S. Process for preparing hydrogen peroxide. U.S. Patent 4,009,252, 1977. (5) Dalton, A. I.; Skinner, R. W. Synthesis of hydrogen peroxide. U.S. Patent 4,336,239, 1982. (6) Kawakami, M.; Ishiuchi, Y.; Nagashima, H.; Tomita, T.; Hiramatsu, Y. Process for producing hydrogen peroxide. U.S. Patent 5,399,334, 1995. (7) Gosser, L. W. Catalytic process for making H2O2 from hydrogen and oxygen. U.S. Patent 4,681,751, 1988. (8) Fischer, M.; Butz, T.; Masonne, K. Preparation of hydrogen peroxide from hydrogen and oxygen. U.S. Patent 6,872,377, 2005. (9) Paoli. M. Process for the production of hydrogen peroxide from hydrogen and oxygen. U.S. Patent 5,194,242, 1993. (10) Germain, A.; Pirard, J.; Delattre, V.; Van Weynbergh J.; Vogels, C. Process for the manufacturing of hydrogen peroxide by direct synthesis from hydrogen and oxygen. U.S. Patent 5,500,202, 1996. (11) Chuang, K. T.; Zhou, B. Production of hydrogen peroxide. U.S. Patent 5,338,531, 1994. (12) Zhou, B.; Lee, L. Supported noble metal, phase controlled catalyst and methods for making and using the catalyst. U.S. Patent 6,919,065 B2, 2005. (13) Degussa/Headwaters Incorporated Building Demonstration Plant for Direct Synthesis of Hydrogen Peroxide. Business Wire. October 12, 2005. Bnet Business Network. http://findarticles.com/p/articles/mi_m0EIN/ is_2005_Oct_11/ai_n15683135 (accessed Sept. 24, 2008). (14) Melada, S.; Pinna, F.; Strukul, G.; Perathoner, S.; Centi, G. Direct synthesis of H2O2 on bemetallic and monometallic catalytic membranes using methanol as reaction medium. J. Catal. 2006, 237, 213. (15) Janicke, M. T.; Kestenbaum, H.; Hagendorf, U.; Schuth, F.; Fichtner, M.; Schubert, K. The controlled oxidation of hydrogen from an
explosive mixture of gases using a microstructured reactor/heat exchanger and Pt/Al2O3 catalyst. J. Catal. 2000, 191, 282. (16) Veser, G. Experimental and theoretical investigation of H2 oxidation in a high-temperature catalytic microreactor. Chem. Eng. Sci. 2001, 56, 1265. (17) Wang, X.; Nie, Y.; Lee, J. L. C.; Jaenicke, S. Evaluation of multiphase microreactors for the direct formation of hydrogen peroxide. Appl. Catal., A 2007, 317, 258. (18) Voloshin, Y.; Halder, R.; Lawal, A. Kinetics of hydrogen peroxide synthesis by direct combination of H2 and O2 in a microreactor. Catal. Today 2007, 125, 40. (19) Voloshin, Y.; Lawal, A. Kinetics of hydrogen peroxide reduction by hydrogen in a microreactor. Appl. Catal., A, submitted for publication, 2008. (20) Pospelova, T. A.; Kobozev, N. I. Palladium catalyzed synthesis of hydrogen peroxide from its elements. II. Active centres of palladium. Russ. J. Phys. Chem. 1961, 35, 262. (21) Samanta, C.; Choudhary, V. R. Direct formation of H2O2 from H2 and O2 and decomposition/hydrogenation of H2O2 in aqueous acidic reaction medium over halide-containing Pd/SiO2 catalytic system. Catal. Commun. 2007, 8, 2221. (22) Choudhary, V. R.; Samanta, C.; Choudhary, T. V. Factors influencing decomposition of H2O2 over supported Pd catalyst in aqueous medium. J. Mol. Catal. A: Chem. 2006, 260, 115. (23) Choudhary, V. R.; Samanta, C. Direct synthesis of H2O2 from H2 and O2 over Pd/H-beta catalyst in an aqueous acidic medium: Influence of halide ions present in the catalyst or reaction medium on H2O2 formation. Catal. Commun. 2007, 8, 73. (24) Beena, B.; Shivanekar, A.; Chudasama, U. Catalytic activity of Pd(II) sorbed on an inorganic ion exchanger zirconium molybdate. J. Mol. Catal. A: Chem. 1996, 107, 347. (25) Barteri, M.; Farinella, M.; Pispisa, B. Dual- vs single-active site mechanism in the catalytic decomposition of hydrogen peroxide. J. Inorg. Nucl. Chem. 1978, 40, 277. (26) Potekhin, V. V.; Solov’eva, S. N.; Potekhin, V. M. Kinetics and mechanism of hydrogen peroxide decomposition in the presence of the tetraaquapalladium (II) complex. Russ. Chem. Bull. 2005, 54, 1105. (27) Salem, I. A.; El-Maazawi, M.; Zaki, A. B. Kinetics and mechanisms of decomposition reaction of hydrogen peroxide in presence of metal complexes. Int. J. Chem. Kinet. 2000, 32, 643. (28) Choudhary, V. R.; Gaikwad, A. Kinetics of hydrogen peroxide decomposition in aqueous sulfuric acid over palladium/carbon: Effect of acid concentration. React. Kinet. Catal. Lett. 2003, 80, 27. (29) Davis, M. E.; Davis, R. J. Fundamentals of Chemical Reaction Engineering; McGraw-Hill: New York, 2003. (30) Lawal, A.; Dada, E.; Lee, W.; Pfeffer, H. Process for the production of hydrogen peroxide from hydrogen and oxygen. U.S. Pat. Appl., 0233695 A1, 2006. (31) Voloshin, Y. Microchannel reactor system for synthesis of H2O2 by direct combination of H2 and O2. Ph.D. Dissertation, Stevens Institute of Technology, Hoboken, NJ, 2008. (32) Kreutzer, M. T.; Du, P.; Heiszwolf, J. J.; Kapteijn, F.; Moulijn, J. A. Mass transfer characteristics of three-phase monolith reactors. Chem. Eng. Sci. 2001, 56, 6015. (33) Wada, Y.; Schmidt, M. A.; Jensen, K. F. Flow distribution and ozonolysis in gas-liquid multichannel microreactors. Ind. Eng. Chem. Res. 2006, 45, 8036. (34) Fogler, H. S. Elements of Chemical Reaction Engineering; PrenticeHall: Englewood Cliffs, NJ, 1992. (35) US Peroxide Home Page. http://www.h2o2.com/intro/properties/ physical.html (Accessed May 21, 2008). (36) Perry’s Chemical Engineers’ Handbook, 7th ed.; Perry, R. H., Green, D. W., Eds.; McGraw-Hill: New York, 1998; Section 14-48.
ReceiVed for reView January 11, 2008 ReVised manuscript receiVed August 18, 2008 Accepted August 19, 2008 IE8000452
