J. Phys. Chem. C 2008, 112, 3619-3626
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Nanoparticle-Catalyzed Clock Reaction Surojit Pande,‡ Subhra Jana,‡ Soumen Basu,‡ Arun Kumar Sinha,‡ Ayan Datta,† and Tarasankar Pal*,‡ Department of Chemistry, UniVersity of North Texas, Denton, Texas and Department of Chemistry, Indian Institute of Technology, Kharagpur - 721302, India ReceiVed: NoVember 8, 2007; In Final Form: December 19, 2007
Bulk Cu2O or cuprite is the only stable copper(I) compound present in plentiful amount in earth’s crust. It is a challenging job to take bulk Cu2O to a nanoregime and to stabilize it in solution. No wonder that Cu2O in its nanoregime would act as a photocatalyst. We report a new synthetic protocol for the first time to obtain monodispersed, stable, exclusively cubic Cu2O nanoparticles in surfactant-free condition and its catalytic action for methylene blue (MB)-hydrazine reaction in aqueous medium. The blue color of the dye, MB, faded away upon the addition of hydrazine, producing colorless leuco methylene blue (LMB) indicating the progress of the redox reaction. The rate of this redox reaction has been found to be enhanced in the presence of the nanocatalyst, Cu2O. The success of the reaction demonstrates a simple ‘clock reaction’. An oscillation between a blue MB color and colorless solution due to formation of LMB is observed on periodic shaking. This oscillation continues for over 15 cycles. Studies on the effect of bulk Cu2O and nanoparticles of CuO and Cu(0) have not been successful for demonstration of the ‘clock reaction’. Thus, the importance of Cu2O nanoparticles in the clock reaction is established beyond doubt. The Cu2O nanoparticles were characterized by different physical methods. TEM studies authenticate the cube shaped monodispersed particles. The electrochemical studies indicate that nano-Cu2O shows a couple of redox peaks which correspond to the redox Cu(II)/Cu(I) system. Kinetic studies authenticate a first-order reaction mechanism. Further, quantum chemical calculations reveal that the nanoparticles reduce the activation energy by ∼17 kcal/mol, thereby making the reaction 2.4 × 107 times faster compared to the gas phase.
Introduction The ‘clock reaction’, one of the most popular types of chemistry demonstrations, is widely available from innumerable experiments. During the progress of a clock reaction an initial induction period occurs before a significant concentration of one of the chemical species involved is produced. Clock reactions not only provide a crowd-pleasing, visually dramatic reversible color change but also provides an engaging illustration of redox phenomena, reaction kinetics, and the principles of chemical titration. There are a few reports of clock reactions in the field of chemistry. One of the best known of these is the Landlot clock reaction between sulfite and excess iodate.1 Typical examples of such reactions are the iodate-bisulfite reaction in the presence of HgCl2, which is similar to that of the iodate-(aminoimino)methanesulfinic acid reaction, and the periodate-dithionate reactions.2,3 Bromate and bromite undergo clock reactions with iodide, iodine and phenol, sulfite, thiourea, N-acetylcysteine, guanylthiourea, etc.4-8 The other clock reactions involve oxychlorides, including chlorite reactions with iodide, bromide, iodine, thiourea, hydroxymethane-sulfinic acid, and thiocyanate.8-10 Chlorate also undergoes the clock reaction with iodine.11 Another well-known model of such reactions is the iodate-arsenous acid reaction, iodine-bisulfate clock, the formaldehyde clock, and the hydration of carbon dioxide.12-15 Snehalatha et al. demonstrates the clock reaction * Author to whom correspondence should be addressed. E-mail: tpal@ chem.iitkgp.ernet.in. † University of North Texas. ‡ Indian Institute of Technology.
involving methylene blue and L-ascorbic acid.16 Before that report, oxygen-dependent reversible color generation of methylene blue was reported from our laboratory.17 Now we have become successful in demonstrating a clock reaction that involves the same dye, methylene blue (MB), and metal nanoparticles in aqueous solution. MB is a water-soluble cationic dye and has a basic dye skeleton of a thiazine group, and it is used as an oxidationreduction indicator in chemistry and biology.18 It is easily reduced to the colorless hydrogenated molecule, leucomethylene blue (LMB), which can, in turn, be oxidized back to its original form. Since the color-fading reaction typically takes only a few minutes, it provides a window to study the redox system under many different reaction conditions in solution phase. Nanoparticle-catalyzed reactions have now become a fascinating field of research.19-21 We have already reported the catalytic properties of coinage metal nanoparticles for a number of chemical reactions.22,23 Cu2O nanoparticle, a p-type semiconductor, is a promising material with potentials for solar energy conversion, micro- into nanoelectronics, magnetic storage devices, biosensing, and catalysis.24-26 In recent years Cu2O nanomaterials with different structures have been synthesized by various methods.27-32 Synthesis of Cu2O nanocubes33 and nanoflowers34 have been reported from aqueous surfactant solution which deserve special mention. However, there is no report for a simple and facile ‘green chemistry’ method for controlled synthesis of Cu2O nanocubes and their subsequent application. This article reports a kinetic strategy to obtain exclusively stable Cu2O nanocubes and its application for a clock reaction involving MB and hydrazine in aqueous medium. The kinetics
10.1021/jp7106999 CCC: $40.75 © 2008 American Chemical Society Published on Web 02/20/2008
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of the reaction have been examined experimentally, and quantum chemical calculation supports the thermodynamics of the catalytic reaction. 2. Experimental Section 2.1. Reagents and Instruments. All the reagents were of AR grade. Double distilled water was used throughout the course of the reaction. CuSO4, glucose, NaOH, hydrazine hydrate, and methylene blue were purchased from Sisco Research Laboratory. Cation exchange resin, SERALITE-SRC-120, was from Sisco Research Laboratory, India. X-ray diffraction study (XRD) was performed from X-ray powder diffraction using a X-ray diffractometer with Cu KR radiation (λ ) 1.5418 Å). Transmission electron microscopy (TEM) measurements of the metal nanoparticles were performed in a Hitachi H-9000 NAR instrument on samples prepared by placing a drop of fresh metal nanoparticle suspension on copper grids precoated with carbon films, followed by solvent evaporation under vacuum. X-ray photoelectron spectroscopy (XPS) analysis was performed on an ESCA LAB MK II using Mg as the excitation source. The samples were prepared by placing one drop of the prepared nanoparticle suspension onto a clean glass slide and then allowing them to dry in air. All absorption spectra were recorded in a Shimadzu UV-160 spectrophotometer (Kyoto, Japan) from a solution in a 1 cm quartz cuvette. A threeelectrode system composed of a glassy carbon electrode (GCE) as the working electrode, a standard hydrogen electrode (SHE) as reference electrode, and a platinum foil counter electrode were used to examine the electroactivity of the synthesized cubic nanoparticles. 2.2. Synthesis of Cu2O Nanoparticles. Cu2O nanoparticles were synthesized by a two-step mechanism. In the first step, we prepared the resin-bound Cu(II) composites. The watersoluble Cu(II) precursor, CuSO4 (1 mL of 0.1 M), was allowed to exchange with the H+ ion of the cation-exchange resin (R-H+, 0.1 g suspended in 5 mL of water) and stored overnight. The resin beads, on which Cu(II) precursor ions were immobilized, were washed several times with water to drain out the liberated H2SO4 and unexchanged CuSO4. The solid Cu(II)-bound resin beads in water remain stable for months. The Cu(II)-bound blue-colored beads were employed for the synthesis of Cu2O nanoparticles in aqueous phase. In the second step, aqueous solutions of glucose (0.2 gm in 5 mL water) and NaOH (250 µL of 1 M) were introduced in succession into a conical flask containing Cu(II)-coated wet resin beads. The pH of the solution was ∼10-11. The mixture was placed on a water bath (∼70 °C) for ∼6-7 min. Slowly, the Cu(II) ions get reduced to the Cu(I) state and enter the solution phase with a brilliant yellow color showing a λmax at ∼470 nm. The yellow supernatant solution with Cu2O (∼10-3 M) nanoparticles remained stable for ∼12 days, and in turn the resin beads turned iridescent sea-green color. The green color is presumably due to the formation of a Cu(II)-gluconic acid complex which remains adsorbed onto the resin beads. Separation of the beads from the Cu2O nanoparticle-containing solution and their subsequent glucose reduction produced a new batch of Cu2O nanoparticles in solution. Cu2O nanoparticles were obtained in batches so long as the resin beads remained loaded with Cu(II) ions. However, the produced Cu2O nanoparticles successively decreased in amount for subsequent batches. In the solution phase but in nitrogen atmosphere Cu2O nanoparticles remained unchanged for many months. The yellowish solution of Cu2O nanoparticles was centrifuged and washed with distilled water and finally with absolute ethanol to obtain powdered Cu2O
Figure 1. XRD pattern of Cu2O nanoparticles.
nanoparticles. The product was dried under vacuum, which remained stable over 5 months. Gram level synthesis of Cu2O nanoparticles was possible employing this synthetic protocol. The percentage yield of Cu2O nanoparticles was 74.4%. Variable concentrations of glucose (0.1-0.4 g) and NaOH (0.04-0.06 M) were used for the formation of Cu2O nanoparticles. Formation of Cu2O nanoparticles was inhibited below pH ∼ 9. 2.3. Clock Reaction. In a typical reaction, 100 µL (∼10-3 M) of Cu2O nanoparticles was mixed with an aqueous solution of MB (200 µL of 5 × 10-4 M) in a 1 cm quartz cuvette, and the volume of the solution was made up to 3 mL. Next, 100 µL of 2.0 M aqueous hydrazine hydrate solution was added to the reaction mixture, and time-dependent absorption spectra were recorded in the UV-visible spectrophotometer at 30 ( 2 °C. The blue color of MB disappeared in 12 min. The solution regained its original blue shade just after (5 s) shaking in air. The visual dramatic reversible color change goes on for about 15 cycles. This might continue for a week when shaken for once or twice a day. Undisturbed solution remained colorless for days together. 3. Results and Discussion SERALITE-SRC-120 cation exchange resin is a cross-linked polymer containing styrene and a small proportion of divinylbenzene groups as an integral part of the polymer lattice with an equivalent number of cations such as H+ ions.35 Cu(II) ions from the precursor, CuSO4 solution, are effectively exchanged with the H+ ions of the resin through an electrostatic attraction. Upon the addition of alkaline glucose, a yellow suspension of Cu2O nanoparticles was obtained from the Cu(II)-bound resin beads. The synthesized Cu2O nanoparticles were characterized by XRD, TEM, XPS, CV, and UV-visible studies. 3.1. Characterization of Cu2O Nanoparticles. The XRD pattern of the Cu2O nanoparticles is presented in Figure 1. The four diffraction peaks at 30, 36, 42, and 62° are indexed to the (110), (111), (200), and (220) planes, respectively, indicating the formation of perfectly crystalline Cu2O nanoparticles. The peak positions are in good agreement with the JCPDS36 (card no. 78-2076). The size of the Cu2O particles was calculated to be about ∼15 nm according to the Debye-Scherrer equation using half-width of the diffraction peaks. The morphology of the as-prepared Cu2O particles was studied by TEM. The image (Figure 2a) shows that the particles are monodispersed and cubic shaped with average diameter 25 ( 5 nm. Evolution of nanoparticles from a solid composite material via a wet chemical reduction technique using a water bath (∼70-80 °C) makes the copper nanoparticles almost monodispersed, presumably due to digestive ripening.37 The HRTEM image of cubic Cu2O nanoparticles has also been indicated in Figure 2b. Selected area electron diffraction (SAED) pattern on a single particle has been shown in Figure 2c. Figure 2d shows the EDAX analysis, which reveals that a randomly selected particle contains copper and oxygen.
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Figure 3. Wide X-ray photoelectron spectrum of the as prepared Cu2O nanocube (a) and high-resolution XPS spectra of Cu2P3/2 (b) and O1S (c).
Cu2O nanocubes. The peak at 936.63 eV (Figure 3b) corresponds to the binding energy of Cu2p3/2 and is in good agreement with data observed for Cu2O.38 Figure 3c indicates the binding energy of O1s at 536.3 eV and is in good accordance with the reported value.38 The reference peak was corrected to C1s (288.6 eV). In order to resolve the interband or excitonic transitions of Cu2O particles we have carried out the UV-visible absorption measurement of Cu2O nanoparticles. The colloidal solution of the particles of Cu2O shows a yellow color having a broad peak with an absorption maximum at ∼470 nm which is shown in Figure 4a. An estimate of the optical band gap has been given for a semiconductor using the following equation:
REp ) K(Ep - Eg)1/2
Figure 2. TEM image (a), HRTEM image (b), SAED pattern (c), and EDAX (d) of Cu2O nanocubes.
XPS as a surface monitoring technique was employed to further support the formation of Cu2O nanoparticles. Figure 3a shows the wide X-ray photoelectron spectrum of the as-prepared
where R is the absorption coefficient, K is a constant, Ep is the discrete photo energy, and Eg is the band gap energy. A classical Tauc approach is further employed to estimate the Eg value of Cu2O nanoparticles.39 The plot of (REp)2 versus Ep based on the direct transition is shown in Figure 4b. The extrapolated value (the straight lines to the X axis) of Ep at R ) 0 gives absorption edge energies corresponding to Eg ) 2.8 eV. This value is quite greater than the value for bulk Cu2O (2.17 eV). The increase in the band gap of the Cu2O nanoparticles is indicative of quantum confinement effects arising from the cube shape of the particles. Figure 5 shows the cyclic voltammograms from the use of a glassy carbon (GC) electrode with and with out the Cu2O
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Pande et al. SCHEME 1: Schematic Representation for the Preparation of Cu2O Nanocubes from Resin Beads
Figure 4. Absorption spectrum of Cu2O nanocube (a) and plot of (REp)2 vs EP for direct transition of Cu2O nanocubes (b). Condition: [glucose] ) 0.2 g in 5 mL of water and [NaOH] ) 0.05 M.
Figure 5. Superimposed cyclic voltammograms of (a) blank glassy carbon electrode and (b) cubic Cu2O nanoparticles at 5 mV/s scan rate of second cycle in 0.1 M NaOH.
dispersed in 0.1 M NaOH at a scan rate of 5 mV/s. The compound shows an irreversible oxidation peak at 0.41 V vs SHE indicating one electron per copper oxidation which is absent in trace a corresponding to the voltammogram of GC devoid of Cu2O. A comparison of this irreversible peak with available data of Cu complexes in the literature clearly suggests the oxidation of Cu(I) to Cu(II) at this potential.40,41 It is also interesting to see that cubic Cu2O helps oxygen evolution by a large extent, and hence this oxide semiconductor sample can be very good for photoelectrolysis of water. The cathodic part is due to adsorbed oxygen since solvent decomposition (oxygen evolution in this case) has occurred because of a more positive switching potential. 3.2. Mechanism of Cu2O Nanoparticle Formation. Hydrazine straightforwardly reduces bulk CuO, Cu2O, and even CuSO4 to metallic copper from aqueous medium. On the other hand, an alkaline solution of glucose reduces CuSO4 solution to copper nanoparticles.23 However, in the present case resinbound Cu(II) ion is reduced to Cu2O nanoparticles in the presence of an alkaline solution of glucose. The unusual reduction product, Cu2O, is obtained because of the presence of solid resin matrix. This happens via the formation of Cu(OH)2 and kinetic control of the reduction. The formation of Cu(OH)2 is possible due to its insolubility in excess alkali, and hence the controlled release of Cu(II) ion from the resin matrix
is understandable from mass action. Alkaline glucose reduces Cu(OH)2 to metastable Cu(OH) and finally to Cu2O nanoparticles upon heating.42 This heating effect provides a unique tight size distribution to the evolved Cu2O nanoparticles, and the sample color always shows a brilliant yellow shade unlike the change in color of Cu2O nanoparticles from yellow to orange to red due to size variation.34,43 The kinetic control and in situ stabilization of particles make the procedure reproducible. It has been reported that the size of Cu2O nanoparticles increases due to the ready supply of Cu(II) ions during the growth process.43 The above discussion is somewhat similar to the reduction of Cu(II)-bound tartrate by Fehling’s solution where bulk Cu2O is produced without any kinetic control. In the present case, controlled formation of Cu2O nanoparticles and simultaneous complexation of Cu(II) ion by gluconic acid are interesting as previously described. Copper(II)-bound resin beads (before the addition of glucose) oxidize the KI, solution but the reported iridescent sea-green-colored solid resin beads obtained after Cu2O nanoparticle formation could not oxidize KI (Scheme 1). A similar observation was also noted with copper(II)-gluconic acid complex (λmax 660 nm) in solution when prepared deliberately. Electronic spectral measurements authenticate the binding effect of gluconic acid toward Cu(II) ions. Thus, complexation of Cu(II) species and stabilization of Cu2O nanoparticles by the in situ-produced gluconic acid are confirmed. 3.3. Mechanistic Aspects of the Clock Reaction. The oxidized and reduced form of MB exhibit an intense absorption band in the region 200-700 nm.44,45 So, the progress of the reaction can be monitored by a UV-visible spectrophotometer measuring the changes in the specified (λmax ) 664 nm) absorbance maxima. An aqueous solution of MB was mixed with Cu2O nanoparticles and hydrazine under ambient conditions. Then the MB color bleaching starts, and the absorbance of the solution was measured at an interval of 1 min. The pH of the reaction medium was in the alkaline (pH ∼ 8) range. With the progress of the reaction, a steady decrease of the absorbance of the dye was noted, as shown in Figure 6a. The blue color of the dye, MB, faded away producing LMB, indicating the progress of the reduction reaction, and finally a colorless solution results. In the absence of Cu2O nanoparticles no such decrease in absorbance of the dye was observed in the experimental time scale. This indicates the importance of Cu2O nanocube catalyst particles in the reaction. Simple shaking the colorless reaction mixture (after MB reduction) or purging air through it under room-temperature condition leads to regeneration of blue color. On standing, the blue color of the solution is reduced again by the excess hydrazine present in aqueous solution. Thus, an oscillation between a blue color and colorless solution due to regeneration of MB and the formation of LMB is observed periodically upon shaking. This observation advocates the ‘clock reaction’ (Scheme 2). The oscillation continues for a week where a dramatic reversible color change is observed once per day, followed by slow MB degradation without showing any peak in the UV-visible region (Figure 6b). After the degradation of MB, the oscillation happens to continue only
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Figure 6. Absorption spectra for successive decolorization of MB (a), comparative account to show the degradation of MB (b), absorbance vs time plot (c), ln A vs time (d), and rate vs concentration (e). Condition: [MB] ) 3.3 × 10-5 M, [N2H4] ) 0.06 M, and [Cu2O] ) 3.3 × 10-5 M.
upon the addition of a fresh batch of MB and hydrazine into the solution. However, no fresh catalyst was needed in this case. This clearly speaks for the stability of Cu2O catalyst in the reaction mixture. Studies on the effect of bulk Cu2O, nanoparticles of CuO and Cu(0), and other reducing agents such as NaBH4 and glucose could not exhibit the ‘clock reaction’. Thus the importance of Cu2O nanoparticles in the ‘clock reaction’ is established beyond doubt. Variable concentration of catalyst Cu2O (1.6 × 10-5 to 6.6 × 10-5 M), MB (1.6 × 10-5 to 5 × 10-5 M), and the reductant hydrazine (0.03 to 0.2 M) have made the demonstration of the clock reaction crowd pleasing and successful. The plot of absorbance as a function of time (Figure 6c) shows a profile of exponential nature, which speaks for a pseudo-first-order reaction. Moreover, a plot of ln At vs time (Figure 6d) gives a
straight line. This observation further confirms a pseudo-firstorder reaction kinetics, and from the slope of this curve, rate constant value k was obtained. Interestingly enough, a plot of rate vs concentration of hydrazine produces a straight line (Figure 6e). This also supports the pseudo-first-order kinetics. The ‘clock reaction’ was monitored at four different temperatures (10, 27, 45, and 60 °C), and by using the Arrhenious equation, the activation energy (Ea) of the reaction is calculated to be 3.71 kJ mol-1. UV-visible spectral measurement reveals that at low concentration of MB (1.6 × 10-6 M) and hydrazine (7 × 10-3 M) independently or jointly adsorb onto the Cu2O nanoparticle surface. In both cases, the blue-shifted nanoparticle peak (470 nm) authenticates the fact. The blue shifting by 7 and 17 nm were observed for MB and hydrazine, respectively. We have
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Figure 7. Schematic representation of the potential energy profile for the reduction of MB with hydrazine in the gas phase. Selected N‚‚‚H bond lengths are given in angstroms.
SCHEME 2: Schematic Representation of the ‘Clock Reaction’ (i.e., color fading and regeneration of MB).
also studied the effect of a wide variety of anions (∼10-4) introduced into the solution containing Cu2O nanoparticles. It was observed that stronger nucleophiles such as borohydride, cyanide, and hydrazine shift the peak position of Cu2O to the blue region effectively. On the other hand, citrate, nitrate, and iodide did not influence the absorption peak to a meaningful extent. However, chloride, bromide, and thiocyanate played an intermediate role in this respect. In order to critically understand the role of the nanocatalyst in the kinetics of this reaction, we have additionally performed quantum chemical analysis for the reduction of MB by hydrazine in the gas phase using a detailed transition-state analysis. The structures of the reactant and the product (Figure 7) are optimized through the MPW1K/6-31+G(d,p) level of theory46 within the Gaussian 03 suite of programs.47 The MPW1K functional has been well-documented to be very suitable for the hydrogen transfer reaction while the popular B3LYP functional is known to underestimate the barrier height.48 The structures of the reactant, product, and the transition state are further verified using additional frequency calculations with no imaginary frequency for the reactant and product and one imaginary frequency for the transition state corresponding to
the H-atom movement between the hydrazine and the MB units. It is clearly seen from Figure 7 that the redox process is mildly endothermic (∆Hreaction298 ) 4.8 kcal mol-1) and the barrier height is quite large (∆Hactivation298 ) 16.3 kcal mol-1) which is 15.4 kcal mol-1 more than for the reaction in the presence of Cu2O nanoparticles. The dramatic decrease in the activation energy for the reaction in the presence of the nanoparticles is a result of stronger chemisorption of the products in comparison to the reactants which removes mild endothermicity of the gasphase reaction and makes it overall exothermic. Assuming that the Arrhenius model is valid for the gas phase reaction as well (with negligible contribution from hydrogen tunneling through the barrier49) and the preexponential factor is same for the catalyzed as for the gas-phase reaction, we calculate that the rate is enhanced by 2.4 × 107 times in the presence of the nanoparticles. The mechanism of the ‘clock reaction’ stands to be interesting as it involves simple redox chemistry devoid of photoactivation.50,51 Hydrazine supplies electrons to MB via cubic Cu2O nanoparticles, as a result of which colorless LMB is produced in solution. Here Cu2O acts as a nanocatalyst which remains active for many days in solution. Excess hydrazine in turn decreases the dissolved oxygen concentration in water which facilitates the LMB formation. The blue color of MB was regenerated in a few seconds when the solution was shaken in air, i.e., exposed to atmospheric oxygen. UV-visible spectra indicate that no oxidative degradation of MB takes place for about 15 cycles of the reversible reaction. This is authenticated from the decreased absorbance value at the unaltered λmax (664 nm) during LMB formation (Figure 6a). The reversible color generation is not inhibited by Cu2O nanocubes. The electron transfer from hydrazine to MB takes place via Cu2O nanocube surfaces presumably because of the reasonable affinity of both the reductant (hydrazine) and oxidant (MB) to the nanocube surfaces (Scheme 3). This is explained by taking the intermediate redox potential value of Cu2O in relation to the oxidant and reductant.35 We might consider a mechanism for the plausible electron-transfer reaction where hydrazine transfers electrons to Cu(I) to reduce it to Cu(0). Then dissolved oxygen oxidizes Cu(0) back to Cu(I) and the electrons thus facilitate the reduction
Nanoparticle-Catalyzed Clock Reaction SCHEME 3: Scheme To Show the Catalytic Activity of Cu2O Nanocubes for the ‘Clock Reaction’
of MB. In turn, the reversible process, i.e., LMB to MB, may hold responsible for the oxidation of Cu(0). However, this fact could not be substantiated with experimental evidence. Hence, we consider the electron relay via Cu2O nanoparticles from donor N2H4 to acceptor cationic MB. Nonspherical (hexagonal) Cu2O nanoparticles have recently been used for the degradation of methyl orange.52 Presumably, because of the very low concentration of Cu2O nanoparticles in the reaction mixture, selective facet-induced catalysis could not be monitored. However, the electron-transfer process might have some bearing on the biological system of O2 transfer via Cu(II)/Cu(I) because of the involvement of Cu-N interaction. Conclusion In conclusion, we have reported a green chemistry protocol for the generation of exclusively monodispersed, stable Cu2O nanocubes at gram level without using any template. Here, the Cu2O nanoparticle-catalyzed clock reaction has been demonstrated using MB and hydrazine in an aqueous medium. Electrochemical measurement reveals the possible application of the semiconductor Cu2O nanocubes for water decomposition. To the best of our knowledge, it is the first attempt to demonstrate a nanoparticle-catalyzed ‘clock reaction’ substantiated both by experiment and quantum chemical calculations. The thermodynamic and the kinetic parameters of the clock reaction have been investigated. The reported clock reaction involves Cu(I)/hydrazine/O2, so it might have a bearing on biological processes. This reaction has the potential for use to determine the concentration of dissolved oxygen in solution through a simple method. Acknowledgment. The authors are thankful to the UGC, NST, DST, and CSIR, New Delhi, and Indian Institute of Technology, Kharagpur, for financial assistance. We are also thankful to Dr. Vijayamohanan and Mr. Bhaskar R. Sathe of NCL, Pune, for extending help in CV analysis. References and Notes (1) Landolt, H. Ber. Dtsch. Chem. Ges. 1886, 19, 1317. (2) Mambo, E.; Simoyi, R. H. J. Phys. Chem. 1993, 97, 13662. (3) Lente, G.; Fa´bia´n, I. Inorg. Chem. 2004, 43, 4019. (4) Simoyi, R. H.; Masvikeni, P.; Sikosana, A. J. Phys. Chem. 1986, 90, 4126. (5) Chinake, C. R.; Simoyi, R. H. J. Phys. Chem. 1996, 100, 1643. (6) Clarke, J. R. J. Chem. Educ. 1974, 51, 255. (7) Sorum, C. H.; Charlton, F. S.; Neptune, J. A.; Edwards, J. O. J. Am. Chem. Soc. 1952, 74, 219. (8) Darkwa, J.; Olojo, R.; Olagunju, O.; Otoikhian, A.; Simoyi, R. H. J. Phys. Chem. A 2003, 107, 9834. (9) Chikwana, E.; Otoikhian, A.; Simoyi, R. H. J. Phys. Chem. A 2004, 108, 11591. (10) Chinake, C. R.; Simoyi, R. H.; Jonnalagadda, S. B. J. Phys. Chem. 1994, 98, 545. (11) Oliveira, A. P.; Faria, R. B. J. Am. Chem. Soc. 2005, 127, 18022. (12) Harrison, J.; Showalter, K. J. Phys. Chem. 1986, 90, 225.
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The contribution from quantum mechanical tunneling can be qualitatively understood within the Wigner approximation for an Eckart potential, Γ(T) ) 1 + (1/24)[1.44νi/ T]2, where νi is the vibrational frequency representing the TS barrier’s curvature. The imaginary mode of 1998 cm-1 corresponding to the N‚‚‚
3626 J. Phys. Chem. C, Vol. 112, No. 10, 2008 H‚‚‚N vibration when substituted into the equation leads to a tunneling broadening of 7.95 % to the classical transition state theory (TST) rate at 298 K. Thus, even with the inclusion of tunneling corrections, the rate of the gas-phase reaction is ∼107 orders of magnitude lower than in the presence of nanocatalysts.
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