Flocculation-Induced Homolysis of Hydrogen Peroxide in Aqueous

flocculation of titanium dioxide nanocrystallites aqueous suspension upon ... values in the absence of nanoparticles indicates that flocculation induc...
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J. Phys. Chem. B 2006, 110, 6123-6128

6123

Flocculation-Induced Homolysis of Hydrogen Peroxide in Aqueous Colloid Solution of Titanium Dioxide Nanoparticles Yikui Du† and Joseph Rabani* Department of Physical Chemistry, The Hebrew UniVersity of Jerusalem, Jerusalem 91904, Israel ReceiVed: NoVember 20, 2005; In Final Form: January 27, 2006

Thermal generation of oxygen and hydroxylated aromatic compounds by hydrogen peroxide, catalyzed by flocculation of titanium dioxide nanocrystallites aqueous suspension upon addition of hydrogen peroxide, is reported. The oxidation involves catalytic cleavage of a peroxide molecule followed by hydroxyl reaction with the organic solutes. The catalytic hydroxylation is associated with formation of TiO2-H2O2 aggregates, which occurs within a specific range of [TiO2]/[H2O2] ratio. Comparison of the activation energy to literature values in the absence of nanoparticles indicates that flocculation induces an increase of the rate without decreasing the activation energy. This is, to the best of our knowledge, a unique case of nanoparticles catalysis driven by formation of a three-dimensional structure of the suspended particles.

Introduction There is a wide range of applications of confined reactive systems such as nanoporous materials for catalytic industrial processing, advanced filtration systems, and decontamination of gaseous and water pollutants. Despite the importance of such systems and the intensive study of nanoporous materials in recent years, understanding the effects of confinement of reactants in the nanostructures is still limited. The present manuscript concerns the thermal decomposition of hydrogen peroxide to hydroxyl radicals in nanovolumes of aqueous TiO2H2O2 three-dimensional structures. The general mechanism of hydrogen peroxide homolysis presumably involves O-O cleavage according to reaction 1 followed by Haber-Weiss chain propagation (eqs 2 and 3).1

H2O2 f 2OH

(1)

OH + H2O2 f H2O + HO2

(2)

HO2 + H2O2 f O2 + H2O + OH

(3)

The slowness of the overall process is a result of the high dissociation energy of reaction 1. The activation energies measured in the gas phase and supercritical water agree with the O-O bond dissociation energy of 42-47 kcal/mol and are similar to the bond dissociation energy of 50 kcal/mol,2-5 although a considerably lower value has been reported for liquid water.6 Combining reduction potentials7 and hydration energies of H2O2 and OH radicals8 yields a value of 42.3 kcal/mol for hydrogen peroxide dissociation into two hydroxyl radicals in water solution at 298 K. The subsequent reactions 2 and 3 are relatively fast, with reaction rate constants of 2-3 × 107 M-1 s-1 for reaction 29 and 0.210-0.511 M-1 s-1 for reaction 3 at acidic pH. The rates of reactions 2 and 3 depend on pH because * To whom correspondence should be addressed. E-mail: rabani@ vms.huji.ac.il. † Postdoctoral fellow from Peking University, Beijing, P. R. China. Present address: State Key Laboratory of Molecular Reaction Dynamics, Institute of Chemistry, Chinese Academy of Sciences, Beijing 100080, China. E-mail: [email protected].

of the ionic dissociation of the reactants H2O2, OH, and HO2, possessing pKa values of 11.612-11.8,13,14 11.815-11.9,16 and 4.817-4.9.18 Enhanced H2O2 decomposition on solid surfaces of metal oxides has been known for a long time. Thus, catalytic decomposition of hydrogen peroxide vapors over solid metal oxides has been attributed to cyclic transfer of electrons between the oxide (or surface impurities) and H2O2.19-22 Formation of superoxide ions and hydroxyl radicals in water during decomposition of hydrogen peroxide over metal oxides supported on alumina showed a correlation between the catalytic efficiency of the metal oxides and the redox potential of the respective metal cations.23 O2-, OH, and O- radicals have been identified in the H2O2-MgO (insulator) system,24,25 while the semiconductors (e.g., ZrO2) and zeolites showed signals of O2- radical ions.25,26 O2- has been observed upon heterogeneous decomposition of H2O2 over TiO2. Rutile TiO2 (in contrast to anatase) shows, in addition, the existence of a pair of surface O-‚‚‚Opairs.27 Catalytic chain decomposition of hydrogen peroxide on iron oxide, involving redox reactions of Fe(II) and Fe(III) with H2O2, has been discussed.28 This mechanism, however, does not account for H2O2 decomposition on oxide surfaces for which there is no higher oxidation state, such as MgO, SiO2, Al2O3, TiO2, CeO2, and ZrO2.29 In the present manuscript we report the thermal formation of hydroxylation products in aqueous solution containing hydrogen peroxide and TiO2 nanocrystallites under highly specific conditions. These conditions differ from any of the earlier reports as the basic requirement for the observed catalytic hydroxylation is the formation of aggregates with (TiO2)n-H2O2-(TiO2)n three-dimensional structures. Experimental Section Na2CO3 (Baker), NaHCO3 (Frutarom), formic acid (Baker), sodium formate (Riedel-de Haen), 2-phenylethanol (C6H5CH2CH2OH, PEA, Aldrich), and H2O2 (Merck) were used as received. An Orion Ross combination glass electrode was used for pH measurements ((0.01 pH units at 25 °C).

10.1021/jp0567117 CCC: $33.50 © 2006 American Chemical Society Published on Web 03/03/2006

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Figure 1. Linear build-up of hydroxyl products from PEA: 10 mM TiO2, 100 mM H2O2, and 5mM PEA at pH 3.8, 24 °C, with no stirring.

Titanium dioxide colloid solutions, 4.7 nm average diameter, have been prepared by hydrolysis of TiCl4 followed by dialysis against dilute HCl until a pH of 2.7.30 TiO2 (0.11 M; molecular concentration) aqueous suspensions at pH 2.7 were prepared and served as stock solution from which further dilutions were made. The pH was adjusted by HCl (Baker) or NaOH (Frutarom). Product Analysis. PEA and Its Hydroxylated Products. Samples were analyzed by HPLC (Merck-Hitachi L-6200A) using a L-4500 diode-array detector. The products were eluted at a flow rate of 1 mL/min from a LiChroCART 75-4 Superspher 100 RP-18 column (Merck). The mobile phase was 30% methanol in water (v/v). Calibrations were carried out with standard solutions of C6H5CH2CH2OH, o-C6H5CH2CH2OH (2HOPEA), m-HOC6H4CH2CH2OH (3-HOPEA), and p-HOC6H4CH2CH2OH (4-HOPEA). The retention times were 5.5 (4HOPEA), 6.7 (3-HOPEA), 8.8 (2-HOPEA), and 16.5 min (PEA). O2 was determined using a Hewlett-Packard GC model 5890 with He gas carrier. All measurements have been carried out in a small (several milliliters) cuvette wrapped with Al foil to completely avoid light, immersed in a thermostat (25 °C unless otherwise indicated). The suspensions were prepared in a dark room under red light, although illumination at 420 nm (75 W Xe lamp, interference filter) yielded results similar to the dark reaction. Results The key observation is the linkage between the catalytic hydroxylation of PEA, oxygen evolution, and flocculation of H2O2/TiO2, indicating that the three-dimensional structure is responsible for the enhanced rate of homolysis. Thus, the role of TiO2 is not just providing a surface on which the peroxide undergoes homolytic dissociation to yield OH radicals. Removal of the TiO2/H2O2 aggregates by decantation leaves a yellow solution of pertitanic acid,31 which shows no catalytic activity. Hydroxylation of PEA. Hydroxylation of PEA takes place in the TiO2 suspension upon addition of hydrogen peroxide (Figure 1), producing 2-, 3-, and 4-HOPEA, which builds up linearly with time up to at least 18 h and concentrations of 20, 40, and 100 µM, respectively. Under the conditions of Figure 1 the molar ratio between product concentration and unreacted [PEA] is less than 3%. At relatively low initial PEA concentration such as 3.3 × 10-4 M, deviations from linearity are pronounced already in the first hour as the reaction products compete with the PEA for the catalyst.

Du and Rabani

Figure 2. Concentration profiles in PEA solutions: 10 mM TiO2, 100 mM H2O2, and 0.33 mM PEA at pH 3.8 at 30 °C. Note that the differences in time profiles compared to Figure 1 are due to the much lower [PEA].

TABLE 1: Relative Concentrations of Hydroxylated PEA 2-HOPEA 3-HOPEA 4-HOPEA total flocculant homogeneous (radiolysis)

0.13 0.29

0.25 0.31

0.62 0.40

1 1

This is shown in Figure 2, where the maximum of the hydroxylation product’s concentrations vs reaction time is observed at ([2-HOPEA] + [3-HOPEA] + [4-HOPEA])/[PEA] ) 1.1. This reflects the competition between PEA and its oxidation products for the oxidizing species, which are presumably hydroxyl radicals. Since both reactant and immediate stable products share an aromatic ring linked to similar aliphatic residues, similar diffusion-controlled hydroxylation rate constants and consequently observation of a maximum at approximately equal concentrations of products and reactant are expected. The build-up initial rates of the three hydroxylation products do not depend on the PEA concentration when [PEA] > 0.5 mM. The relative concentrations of 2-HOPEA, 3-HOPEA, and 4-HOPEA observed at 10 mM TiO2 and 100 mM H2O2 at pH 3.6 are shown in Table 1 for the heterogeneous catalysis. Comparative results of bulk hydroxylation are also presented in Table 1. The bulk OH radicals were produced by γ-radiolysis of aqueous solution containing PEA in the presence of nitrous oxide. Under such conditions, almost all radiation primary intermediates are converted to hydroxyl radicals. The results of Table 1 show comparable amounts of 2-HOPEA, 3-HOPEA, and 4-HOPEA produced by bulk OH, with a moderate preference of the latter. On the other hand, the thermal process results in the formation of much higher concentrations of the 4-HOPEA, predominantly at the expense of 2-HOPEA. This observation is attributed to a steric effect induced by adsorption of the reactant molecule to the TiO2 nanocrystallites. Figure 3 demonstrates the effect of hydrogen peroxide concentration at constant TiO2 and PEA at pH 2 and 3.6, respectively. At pH 3.6, flocculation is observed at [H2O2]/ [TiO2] > 2.5 as a pale yellow cloud, building up slowly upon addition of H2O2 to the TiO2 suspension. After 2-12 h, depending on the H2O2 concentration, a pale yellow cloudy precipitate is observed. Its formation has no significant effect on the rate of the catalytic oxidation (measured during the first 48 h). The rate of H2O2 homolysis increases linearly with [H2O2] above 0.025 M until a plateau is reached at [H2O2] > 0.2 M. At lower [H2O2], when flocculation is not pronounced, deviation from linearity is observed (see Figure 3, pH 3.6). The correlation between flocculation and the enhanced rate indicates that TiO2-

Flocculation-Induced Homolysis of Hydrogen Peroxide

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Figure 4. Effect of pH on the initial rate of hydroxylation: 100 mM H2O2, 5 mM PEA, 10 mM TiO2 at 24 °C.

Figure 5. Initial formation rates of HO-PEA at constant [H2O2]/ [TiO2]: [H2O2]/[TiO2] ) 10, 5 mM PEA at pH 3.8 and 28 °C. Figure 3. Effect of hydrogen peroxide on initial formation rates: 10 mM TiO2, 5 mM PEA at (A) pH 2 (20 °C) and (B) pH 3.6 (22 °C).

H2O2 aggregates are essential for catalysis. At pH 2 the rate is considerably slower than that at pH 3.6. Flocculation accompanied with a yellow color (darker than at pH 3.6) is observed above 0.06 M. Again, the rate increases linearly with [H2O2] until a plateau is reached above 1 M. The deviation from linearity, clearly observed at pH 3.6 at low H2O2 concentrations, is not observed at pH 2, presumably because of the much higher concentration range used for measurement of the rate. The higher concentration of H2O2 required for reaching a plateau at pH 2, compared to pH 3.6, reflects a higher adsorption coefficient at the higher pH. Further detailed study of the pH effect is shown in Figure 4. The maximum rate is observed at pH ≈ 3.5. At lower pH equilibrium (eq 4) is shifted to the left and thus flocculation is inhibited. At pH > 4 the stability of the -Ti-O2H segment decreases because of the effect of higher pH on decreasing the positive surface charge of the TiO2 nanocrystallite.

(TiO2)nz+ + H2O2 h [(TiO2)n-1-O2-Ti‚‚‚OjOjH](z-1)+ + H+ (4) The higher activity and the faster increase of rate with hydrogen peroxide concentration at pH 3.6 compared to pH 2 are attributed to enhanced chemisorption of H2O2 to TiO2. The effect of TiO2-H2O2 catalyst concentration on the hydroxylation rate is shown in Figure 5 at constant [H2O2]/ [TiO2] ) 10. The rate increases linearly with catalyst concentra-

Figure 6. Effect of [TiO2] on the hydroxylation rate: 100 mM H2O2, 5 mM PEA at pH 3.8, 26 °C.

tion. This is expected as added catalyst provides additional active sites for hydroxylation. The rate at constant H2O2/TiO2 increases linearly with the concentration of the two, suggesting that the nature of the flocculant does not change. The effect of [TiO2] on the hydroxylation rate at constant hydrogen peroxide concentration is shown in Figure 6. The rate increases with [TiO2] at relatively low concentration, apparently because of increasing the TiO2-H2O2 complex concentration. A maximum is observed at [TiO2]/[H2O2] ≈ 0.06. The rate decreases upon further addition of TiO2 as the ratio of [H2O2]/ [TiO2] decreases below the flocculation point. This, again, indicates the importance of flocculation in the catalytic process.

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Du and Rabani

TABLE 2: Catalytic Efficiency of Different TiO2: Yields of HPEAa pH [4-HOPEA], µM/h [3-HOPEA], µM/h [2-HOPEA], µM/h

300B

TiO2 self-made

2.8 7.2 2.9 1.6

3.8 8.3 3.3 1.8

a The reaction system contains 10 mM TiO2, 100 mM H O , and 5 2 2 mM PEA, 26 °C, and no stirring after mixing.

Figure 8. Effect of [PEA] on the initial rate of oxygen evolution. Suspension of 10 mM TiO2 and 100 mM H2O2 (initial concentration), pH 3.6, 28 °C.

Figure 7. Evolution of O2. Suspension containing 10 mM TiO2 and 100 mM H2O2 (initial concentration) at pH 3.6. The oxygen concentration represents the total amount of oxygen divided by the volume of the suspension, although most of the oxygen is actually present in the gas phase.

Comparison between different TiO2 preparations is presented in Table 2. The commercial 300B TiO2 is a basic suspension; there is no observable homolysis in unbuffered suspension of 10 mM 300B TiO2, 100 mM H2O2, and 5 mM PEA if no attempt is made to change the pH of 11. The results in Table 2 showing the similarity of 300B and the self-made TiO2 were obtained after addition of H2SO4 to adjust the pH to 2.8. An exceptionally high rate was observed with the commercial 300A suspension at pH 1.5 under conditions similar to Table 2. The respective rates for 2-HOPEA, 3-HOPEA, and 4-HOPEA were 3.7, 7.6, and 16.7 µM/h. No hydroxylation of PEA was observed under similar experimental conditions when TiO2 powder (such as Degussa’s P-25) was used. Similarly, no observable hydroxylation took place in layers made of 300B or the self-made TiO2. Oxygen Evolution. Oxygen evolution is observed in the absence as well as at moderate concentrations of PEA. Typical results are shown in Figures 7 and 8. The nonlinear formation of O2 becomes pronounced at O2 concentration > ∼13 mM, corresponding to depletion of about 25% of the initially present hydrogen peroxide. Added PEA suppresses the formation of O2, as can be seen from both Figures 7 and 8. Figure 8 shows the dependency of the initial rate of oxygen evolution as a function of [PEA]. This rate decreases from 1.5 mM/h in the absence of PEA to 1 mM/h in the presence of 1 mM PEA. This effect is expected in view of the competition between PEA and H2O2 for the hydroxyl radicals. Figure 7 shows that after 50 h the oxygen evolution amounts to 75% of the initial hydrogen peroxide. This demonstrates the catalytic nature of the reaction, since the decomposed peroxide concentration exceeds by severalfold the molecular TiO2 concentration. Activation Energy. The hydroxylation of PEA and evolution of oxygen show strong temperature dependency, as shown in Figures 9 and 10, where the natural logarithm of the initial rate

Figure 9. Arrhenius plot for the hydroxylation of PEA (4-HOPEA). R is the initial hydroxylation rate, T is the absolute temperature, [TiO2] ) 10 mM (self-made), [PEA] ) 5mM, [H2O2] ) 50 (open circles) or 100 mM (closed circles) at pH 3.6. Note that the temperature range of the measured values is 5-50 °C. Under the conditions of the two deviating points, measured at 40 and 50 °C, respectively, the reaction rates were not the “initial” ones as the relatively rapid accumulation of products may have given rise to secondary reactions.

Figure 10. Arrhenius plot for oxygen build-up. [TiO2] ) 10 mM (selfmade), [H2O2] )100 mM, no PEA, pH 3.6. The slope yields Ea ) 20-25 kcal/mol.

is plotted against 1/T. Although strictly speaking these are not Arrhenius plots, since it is the rate rather than the rate constant, the difference does not influence the observed activation energy if the kinetics is simple first or second order. The linear behavior

Flocculation-Induced Homolysis of Hydrogen Peroxide within a fairly large temperature range suggests that this is indeed the case. The activation energy is the same (23 ( 2 kcal/mol) for both systems containing 50 and 100 mM H2O2. The activation energy does not depend on [H2O2] in the range of 50-100 mM, although the rate changes considerably (Figure 3), indicating no saturation of the TiO2 surface. Note that a similar activation energy is measured for molecular oxygen evolution in the presence as well as absence of PEA (Ea ) 25 ( 5 kcal/mol, Figure 10). This is in agreement with the cleavage of the -O-OH bond in hydrogen peroxide being rate determining in both systems. Discussion General. The thermal PEA hydroxylation reactions show that spontaneous formation of hydroxyl radicals takes place. Only negligible amounts of hydroxylated products are observed in the absence of TiO2, and no hydroxylation whatsoever takes place in the absence of H2O2, even at 50 °C. On the other hand, hydroxylation of PEA proceeds in the absence of hydrogen peroxide upon band gap excitation of the TiO2, which is known to produce hydroxyl radicals. The similar activation energies for PEA-hydroxylation and oxygen evolution (the latter in the absence of PEA) strongly suggest that the rate-determining step is hydrogen peroxide cleavage to OH radicals. The activation energy for the noncatalytic dissociation of H2O2 into two OH radicals in aqueous solution, 17 kcal/mol, has been measured by Takagi and Ishigure,6 who confirmed earlier results of 1832 and 2233 kcal/mol, respectively. These values are lower than the experimental value of 22-23 kcal/ mol measured in the presence of TiO2 nanocrystallites under the conditions of Figures 9 and 10. Thus, enhancement of reaction 1 under conditions leading to flocculation is not due to a lower activation barrier. Recently Hiroki and LaVerne29 reported catalytic decomposition of hydrogen peroxide on various metal oxide surfaces in aqueous suspension. Their reported activation energies of 9-11 kcal/mol at 5 µM H2O2 are considerably lower than that in the absence of oxide. The difference between their work and the results presented here indicate different surface states as we observed only a negligible effect of TiO2 unless flocculation occurs. Li, Chen, and Zhao34 and Wu et. Al.35 reported enhanced decomposition of H2O2 by visible light in the presence of TiO2 (Degussa’s P-25). Apparently the P-25 TiO2 behaves differently than the TiO2 preparations used in the present work. Mechanism Analysis. Upon addition of hydrogen peroxide to the TiO2 colloid solution, at first a layer of hydrogen peroxide is produced on the surface via coordination bonds to surface Ti atoms.36 Part of the bound hydrogen peroxide may become ionized (deprotonation) and thus decrease the surface charge and the repulsive forces between the TiO2 particles. Attractive van der Waals forces between the particles will subsequently produce catalyst agglomeration. Pairs of nanoparticles may become bound by hydrogen bonding and possibly also by -Ti‚‚O-O‚‚Ti- bridges as illustrated in Figure 11. From the level of hydrogen peroxide, which is required to reach a plateau in the rate of PEA hydroxylation vs H2O2 concentration (Figure 3), it is obvious that the overwhelming majority of the H2O2 molecules cannot be bound to the TiO2. The flocculant just acts as a catalyst by confining hydrogen peroxide molecules in the three-dimensional structure. Obviously this system is rather dynamic, exchanging

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Figure 11. TiO2 nanocrystallites bound by sharing H2O2 molecule. The hydrogen peroxide forms coordination bonds with Ti. Two nanocrystallites may share more than a single H2O2 molecule, while this type of binding is not limited to two TiO2 particles.

hydrogen peroxide in the bulk with molecules confined within the three-dimensional structure. Confinement is known to have pronounced effects of thermophysical properties of fluids, increasing the densities in the pores compared to the bulk and changing phase equilibria.37 It was recently reported that chemical reactions, such as monomer-dimer equilibrium, are also influenced by confinement of the reactants to a small volume.38 In a recent theoretical study39 the effect on chemical equilibria has been attributed mainly to the higher density of the reactant in the pores compared to the bulk. The higher density, which drives the equilibrium toward the dimers, is achieved because of the wall-fluid interactions. In the case of molecules confined in a nanocrystallite structure such as in our case, chemical and energetic heterogeneity of the pore walls due to nanocrystallite surface states induce preferential adsorption of the hydrogen peroxide molecules. The confinement to a small volume may affect the vibration energy levels of water and hydrogen peroxide clusters and enhance the HO-OH cleavage. It might be argued, by analogy to the earlier works,37,38 that the confinement is expected to enhance the bimolecular dimerization of the monomer molecules, namely, stabilize the hydrogen peroxide rather than catalyze its homolysis. However, the HO-OH bond cleavage requires H2O2 encounters with other molecules and clusters, transferring to the peroxide their vibration energies. These encounters are enhanced in the confined volume, where the concentration of molecules is higher compared to the bulk, increasing their frequency. Furthermore, the physical properties of hydrogen peroxide as well as water clusters are likely to differ in the confined volume, and it is conceivable that the vibration energy levels and distribution may involve a better match to the HO-OH bond dissociation energy. Although the back reaction of OH radicals is also enhanced, there are two factors which may inhibit this reaction. First, the lifetime of hydroxyl radicals in the confined volume is relatively short because they react with the PEA substrate and hydrogen peroxide. These reactions are also expected to enhance in the confined volume because of the high local concentrations of the reactants. In this model the rate of energy transfer to the -O-O- bond in the confined system is enhanced, while the role of the bound H2O2 is stabilization of the three-dimensional catalytic structure. While the mechanism proposed above applies to all kinds of bound or free hydrogen peroxide molecules in the confined volume, an alternative mechanism involves preferential cleavage of -O-O- bonds in the Ti-O-O-Ti bridges. In such a case the hydroxyl radical pairs are produced on the surface of two different nanocrystallites. This may hinder their recombination, particularly if the atomic distances between Ti and O in TiOH are different than in Ti-O-O-Ti. We are unable, however, to provide an explanation for a preferential cleavage of bridging -O-O- bonds. Acknowledgment. The authors are indebted to Eran Rabani, Tel Aviv University, Israel, for invaluable discussions.

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