Methylene Blue–Cu2O Reaction Made Easy in Acidic Medium - The

Nov 20, 2012 - MB•+, i.e., the oxidized form, has a peak at 520 nm.(20) Keeping this redox possibility in mind, we started our present work. First, ...
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Methylene Blue−Cu2O Reaction Made Easy in Acidic Medium Mrinmoyee Basu,† Arun Kumar Sinha,† Mukul Pradhan,† Sougata Sarkar,† Anjali Pal,‡ Chanchal Mondal,† and Tarasankar Pal*,† †

Department of Chemistry and ‡Department of Civil Engineering, Indian Institute of Technology, Kharagpur-721302, India S Supporting Information *

ABSTRACT: Truncated Cu2O cubes with well-defined morphology have been successfully synthesized using a stable Cu(II)−EDTA precursor and glucose as the reducing agent in alkaline conditions under 5 min of microwave irradiation. The truncated cubic Cu2O particles as a solid powder were characterized by different physical methods. Curiously enough, upon the addition of Cu2O particles, the blue-colored methylene blue (MB) in aqueous acidic (pH ≈ 1.0) solution successively bleaches to colorless leuco-MB (LMB). The redox reaction generates a colorless solution which easily reverts back to blue in air. The reaction has been quantitatively monitored by UV−vis spectrophotometry. The reversible color change, i.e., oscillation between a blue MB solution and colorless LMB solution happens to be a periodic phenomenon for more than 50 cycles and is reproducibly demonstrated as a simple “clock reaction”. Acidic (pH ≈ 3.0) conditions favor zeroth-order kinetics for the underlying redox phenomena. Dilute H2SO4 has been proven to be the best choice to provide a passive reaction medium, and the undisturbed reaction mixture showed oscillatory behavior even after one month.

1. INTRODUCTION A most crowd -pleasing and visually dramatic reaction among the most popular types of chemistry demonstrations which are widely available from innumerable experiments is the “clock reaction”. Clock reactions are a special class of chemical processes characterized by an abrupt change in the concentration of a chemical species after a time lag. The most well-known and popular clock reaction is the Landolt clock reaction between sulfite and excess iodate.1 There are various reports of clock reactions which include iodate−bisulfite reaction in the presence of HgCl2; bromate also undergoes clock reactions in the presence of halides involving some organic molecules.2−4 The clock reaction involving dye molecules such as methylene blue (MB) serves the purpose well. The cationic thiazine dye MB having λmax at 663, 614, and 292 nm is blue in color and water-soluble, which is why it has been widely used since its synthesis in 1876. Methylene blue can be doubly reduced to the colorless hydrogenated molecule leucomethylene blue (LMB) by various reducing agents. LMB is stable in only deaerated conditions and can, in turn, be oxidized back to the blue-colored MB. Many researchers have focused their research on this doubly reduced form of MB. This redox reduction of MB to LMB and vice versa can be properly monitored using appropriate reducing and oxidizing agents. MB also undergoes a clock reaction with Lascorbic acid.5 Previously, we published a report on a clock reaction which occurs in strongly alkaline conditions and in which hydrazine was used as the reducing agent.6 In that paper, we proved that without a reducing agent the clock reaction is insignificant. Our continued search in this direction did not find a single report of a clock reaction in acidic conditions without the addition of any reducing agent but involving MB. With the successful compilation and involvement of MB, we thought the © 2012 American Chemical Society

present text would be academically encouraging and important from a crowd-pleasing demonstration point of view. In recent years, controlled synthesis of inorganic micro- and nanostructures having a uniform shape and size has attracted wide-ranging research interest because of their unique physical properties.7 Cuprous oxide (Cu2O) is an important p-type semiconductor having a direct band gap value of 2.17 eV.8 Cu2O is a promising material having potential applications in solar energy conversion, gas sensors, CO oxidation, electrode materials, photocatalysis, and decomposition of water to H2 and O2 under visible light.9 Because of such wide and potential applicability of Cu2O, there always exists a need-based interest for the synthesis of various Cu2O nanostructures having uniformity in size. A variety of Cu2O nanostructures have already been synthesized which include nanowires, nanoboxes, nanocubes, nanotruncated cubes, nanooctahedra, nanocages, nanomultipods, nanospheres, and various hollow structures.10 Therefore, different methods have been developed for the synthesis of different architectures of Cu2O which include the hydrothermal method, microemulsion method, surfactantassisted route, wet chemical method, etc.11 From the previous reports it is ascertained that various Cu2O nanostructure have been synthesized using simple Cu(II) salts and various reducing agents such as glucose, hydrazine, ascorbic acid, etc.6,12 Huang et al. have synthesized different shapes of Cu2O nanostructures starting from cubes to hexapods from a simple wet chemical method using hydroxylamine as the reducing agent.13 Zhang et al. have synthesized a double-tower-tip-like morphology in the Received: October 11, 2011 Revised: October 9, 2012 Published: November 20, 2012 25741

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oxide were performed on a Hitachi H-9000 NAR instrument on samples prepared by placing a drop of fresh metal oxide sol on copper grids precoated with carbon films, followed by solvent evaporation under vacuum. 2.2. Synthesis of Cu2O Truncated Cubes. Well-defined Cu2O nanostructures were synthesized in the solution phase from a simple wet chemical method at high pH from an aqueous solution using the Cu(II)−EDTA complex as the precursor complex and glucose as the reducing agent under microwave irradiation. At first, the Cu(II)−EDTA complex is prepared from CuSO4 and Na2−EDTA. This complex formation is authenticated from UV−vis study as the Cu(II)−EDTA complex has a strong absorption maximum at 730 nm. After that, 5 mL of 0.1 M glucose and 5 mL of 0.1 M NaOH were introduced to this solution and irradiated under microwave for 5 min (Scheme 2). The solution became red in a few minutes,

water/oil microemulsion method in the presence of cetyltrimethylammonium bromide (CTAB).14 Here, we have successfully synthesized Cu2O truncated cubes using the Cu(II)−EDTA complex as the precursor compound and glucose as the reducing agent in alkaline conditions under microwave irradiation. The as-synthesized truncated Cu2O cubes were characterized by UV−vis, FTIR, X-ray powder diffraction (XRD), X-ray photoelectron spectroscopy (XPS), field emission scanning electron microscopy (FESEM), and transmission electron microscopy (TEM) studies. The synthesized Cu2O truncated cubes are a potential candidate for an interesting clock reaction where MB has been taken in strongly acidic conditions (pH ≈ 1.0). The reaction progressed as a result of oxidation of Cu(I) to Cu(II) and the reduction of MB to LMB at pH ≈ 1.0 (Scheme 1). In this reaction the consumption of Cu(I) in each Scheme 1. Schematic Representation of the Synthesis of Cu2O and Its Use for the MB−LMB Clock Reaction in Acidic Solution

Scheme 2. Schematic Representation for the Synthesis of Cu2O from the Cu(II)−EDTA Complex under Microwave Irradiation

cycle occurs. Therefore, addition of Cu2O substrates accelerates the reaction. This clock reaction continues up to several cycles depending on the amount of added Cu2O truncated cubes.

and then a red precipitate was obtained. After 5 min, red-colored particles were collected, washed thoroughly first with water and then with ethanol, and reserved for further characterization. Gram-level synthesis of truncated Cu2O cubes was possible from this synthetic protocol. The percentage yield of Cu2O was 90%. The percentage yield was calculated on the basis of the amount of Cu(II)−EDTA precursor complex. 2.3. Procedure of the Clock Reaction. The as-synthesized truncated Cu2O cubes were used for the clock reaction. For this 30 mL of 3 × 10−5 M MB was taken, and using dilute H2SO4 solution, pH ≈ 3.0 was maintained. To this solution was introduced 0.01 g of solid Cu2O. Just after the addition of Cu2O, MB color bleaching started and the colorless solution regained its original blue color when the reaction mixture was allowed to stand unstirred.

2. EXPERIMENTAL PROCEDURE 2.1. Materials and Analytical Instruments. All the reagents used are of AR grade and have been used without further purification. Doubly distilled water was used throughout the experiment. Copper(II) sulfate was purchased from S. D. Fine-Chem Ltd. Na2−EDTA and NaOH were purchased from Hi Media Laboratory Pvt. Ltd. Glucose was received in standard grade from Sisco Research Laboratory Pvt. Ltd. Methylene blue was purchased from LOBA Chemie. All absorption spectra for the performance of the clock reaction were recorded on a Shimadzu UV-160 spectrophotometer (Kyoto, Japan) by taking the solutions in a 1 cm quartz cuvette. The electronic absorption spectrum of solid Cu2O using DRS (diffuse reflectance spectra) mode was recorded with a Cary model 5000 UV−vis−NIR spectrophotometer. FTIR spectral characteristics of the samples were collected in reflectance mode with a Nexus 870 Thermo-Nicolet instrument coupled with a Thermo-Nicolet Continuum FTIR microscope. The phase and purity of the product were determined by XRD using an X-ray diffractometer with Cu Kα radiation (λ = 1.5418 Å). Scans were collected on dry nanoproducts in the range of 10−80°. Measurements were performed at room temperature. The XPS spectrum was recorded for the as-synthesized truncated Cu2O cubes using a Perkin-Elmer model 1257 with a nonmonochromatized Mg Kα line at 1253.6 eV. The particle size, shape, and morphology of the Cu2O truncated cubes were observed with a field emission scanning electron microscope (Supra 40, Carl ZEISS Pvt. Ltd.), and an energy-dispersive spectrometry (EDS) machine (Oxford link and ISIS 300) attached to the instrument was used to obtain the composition of the product. TEM and high-resolution TEM (HRTEM) measurements of the metal

3. RESULTS AND DISCUSSION The as-synthesized truncated Cu2O cubes were characterized using different physical methods. 3.1. UV−Vis Spectroscopy. The optical property of the freshly prepared Cu2O sample was ascertained with the help of UV−vis spectroscopy to resolve the excitonic or interband transitions. Red-colored Cu2O was taken for UV−vis study in solid form. The electronic spectral information clearly indicates that there is a sharp absorbance band at 550 nm for Cu2O, which is in good agreement with the reported data (see Figure S1, Supporting Information).15 Therefore, UV−vis spectrophotometry demonstrates the formation of the Cu2O nanostructure. 3.2. FTIR Analysis. The quality and composition of the assynthesized truncated cube-shaped Cu2O nanostructure was characterized from an FTIR study. The literature reports that Cu2O has only one IR-active mode, which appears at 610 cm−1.11d,16 Our Cu2O sample shows only one peak at 624 cm−1, which is attributed to the Cu−O vibration of Cu2O crystals (Figure S2, Supporting Information). Here an observable shift of 25742

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the prepared sample. The peaks in the spectra are asymmetric in nature, which indicates that the truncated cubes are composed of more than one chemical state of copper. The peak fit of the Cu 2p3/2 peak reveals a main peak at 932.5 eV and is accompanied by a series of satellite peaks on the high-binding-energy side, 933.5, 940.3, and 943.1 eV. The main peak is known as a characteristic of Cu+, and the shakeup satellite peaks correspond to the Cu2+ state.18 Therefore, one thing has been unveiled: though XRD study affirmed that the prepared cubes are of pure Cu2O, the XPS spectra ascertained the presence of a thin layer of CuO on the Cu2O surface. 3.5. Structural Characterizations of Cu2O: FESEM Analysis and EDS Analysis. The morphology and size of the as-prepared Cu2O were revealed from FESEM. FESEM images of the sample were obtained from the reduction of the Cu(II)− EDTA complex by glucose in alkaline conditions under microwave irradiation and indicate that the sample is composed of highly uniform truncated cubic crystals. From the lowmagnification image of the prepared sample, it is clear that a huge amount of Cu2O was produced under these reaction conditions (Figure 3a,b). From the high-magnification view of the cubic Cu2O crystals, it can be seen that all are well-defined truncated cubes having regular shapes and smooth faces (Figure 3c−f). The morphology of the truncated cubic structure shows a high degree of symmetry. The high-magnification FESEM image also indicates that the truncated edges have lengths and heights in the 1.5 μm range and widths in the 350 nm region (see Figure S3a, Supporting Information). The full array of one truncated cube is in the micrometer range. For detailed structural characterization, TEM and HRTEM were used. The TEM image at low magnification indicates that all the truncated cubes are of of uniform shape and size (see Figure S3b, Supporting Information). The TEM image of a single Cu2O crystal shows the truncated cubic structure, which is in good agreement with the FESEM image (Figure 4a). The truncated nature is also observed from the dark-field TEM image (Figure 4b) of Cu2O. From the HRTEM image of the as-prepared Cu2O truncated cubes, lattice distances were calculated which are 0.30 and 0.21 nm corresponding to (110) and (200) lattice plane spacing (Figure 4c).19 EDS analysis confirmed the presence of the elements Cu and O in truncated Cu2O cubes (see Figure S4, Supporting Information).

the absorption band in the FTIR spectrum of Cu2O was noticed. The finite size of the nanoparticles is basically due to the breaking of a large number of bonds for surface atoms, resulting in the rearrangement of unlocalized electrons on the particle surface. As a result, when particles are reduced to nanoscale dimensions, the absorption bands of the FTIR spectrum shift to higher wavenumber.16 There is no peak beyond 600 cm−1, which is an indication of the absence of CuO as an impurity. As the FTIR spectrum contains only one peak, the synthesized product is pure Cu2O. 3.3. XRD Analysis. The composition and phase purity of the product were examined by X-ray diffraction analysis. Figure 1

Figure 1. XRD patterns of truncated Cu2O cubes.

shows the XRD spectrum of the as-prepared Cu2O nanostructure. All of the diffraction peaks are indexed according to the standard cubic structure of Cu2O (space group Pn3m, JCPDS file no. 05-0667).17 The XRD spectrum shows the diffraction peaks at 2θ values of 29.4°, 36.3°, 42.3°, 61.3°, 73.5°, and 77.2°, which correspond to the crystal planes of (110), (111), (200), (220), (311), and (222), respectively. Peaks due to impurities of Cu(0), CuO, and Cu(OH)2 were not been detected from the XRD pattern. Therefore, from the XRD spectrum we can conclusively claim that the as-synthesized truncated Cu2O cubes are not only highly crystalline but also phase pure. 3.4. XPS Analysis. The surface property of the as-synthesized Cu2O truncated cubes was studied with the help of XPS analysis. XPS is a very powerful technique to study the elemental composition and electronic state of an oxide. Figure 2 shows the wide-range and high-resolution X-ray photoelectron spectra of

Figure 2. XPS spectra of truncated Cu2O cubes: (a) wide-range and (b) high-resolution spectra of Cu 2p3/2. 25743

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Figure 3. FESEM images of truncated Cu2O cubes with a symmetrical structure, a tight size distribution, and smooth faces: (a, b) low magnification, (c, d) medium magnification, (e, f) high magnification.

Figure 4. (a) TEM image at high magnification. (b) Dark-field TEM image. (c) HRTEM image of truncated Cu2O cubes.

4. GROWTH MECHANISM OF CU2O The copper(II)−citrate complex can be reduced by glucose to form Cu2O precipitates, which are widely used in the analytical determination of saccharides. This idea has been extended for the synthesis of uniform Cu2O microcrystals from the solution-phase reaction between the Cu(II)−EDTA complex and glucose in NaOH medium under microwave irradiation:The above equation is the overall chemical reaction for the synthesis of

the Cu2O microstructure. In the intermediate stage Cu(OH)2 is formed and reduced further in steps by glucose to give Cu2O. The high stability constant value of the Cu(II)−EDTA chelate plays a major role in the slow hydrolysis of the complex, leaving the smooth faces for Cu2O truncated cubes. 25744

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5. CLOCK REACTION The pure as-synthesized uniform truncated Cu2O cubes are employed successfully as a novel reagent in the clock reaction in acidic pH, i.e., pH ≈ 3.0, employing the water-soluble bluecolored (λmax = 663 nm) cationic thiazine dye MB (Scheme 3).

oxidizing agent. Here, we tried Cu2O as an oxidizing agent because Cu2O is a well-known reagent in organic synthesis, but surprisingly, upon the addition of Cu2O to the acidic MB solution, a sudden color bleaching took place for MB. After that, the colorless solution turned slowly to blue. The blue coloration of the aqueous phase started to appear at the upper meniscus of the reaction vessel; i.e., the solution which remained in contact with the oxygen/air was oxidized first. In a few minutes the solution regained the original blue color. After that, if we shook the reaction mixture, again the blue solution was discharged. This cycle went on for several rounds depending on the amount of the added Cu2O in the presence of air and went on reversibly. This oscillation between blue and colorless solution is a definite and well-known redox cycle of the MB−LMB reaction. The blue color is the oxidized form, and the solution becomes colorless due to the doubly reduced form of MB, i.e., LMB. This is the first demonstration of the clock reaction in acidic pH in the presence of an oxide nanostructure and also in the absence of any other extra reducing agent. From our earlier report we know that the reduction process, i.e., MB to LMB, takes place after the addition of catalyst and a reducing agent. In our present case, the clock reaction monitored using spectrophotometry becomes very difficult because of the promptness of the reversible reaction. The progress of the forward reaction, i.e., MB to LMB, cannot be monitored so easily by UV−vis spectroscopy because of the very fast color bleaching, but the process becomes crowd-pleasing, simple, and interestingly demonstratable because of the comparatively slower backward reaction. However, in practice at the time of recording the backward reaction, a problem arises due to interference of the onset of the forward reaction that starts simultaneously. Actually for the backward reaction there is no role of Cu2O which is present in the reaction vessel. At pH ≈ 1.0 the clock reaction is pretty fast (with the prescribed amount of catalyst), so it is not easy to monitor the redox process as mentioned already. Hence, the optimum pH was chosen to be somewhat higher (∼3.0) for monitoring the reaction and to record the reversibility of the reaction comfortably and reproducibly. Deliberate use of still higher pH conditions (>3.0) slows the reduction of MB. However, for demonstration and crowd-pleasing purposes, pH ≈ 1.0 seems to be ideal. MB and its reduced form LMB have their characteristic and strong absorbance bands in the 200−700 nm range, which was used to monitor the clock reaction. Progress of the forward reaction, i.e., MB to LMB, can be monitored spectrophoto-

Scheme 3. Schematic Representation of the Clock Reaction Involving MB and Cu2O in Acidic pH

From the literature we know that blue MB can be doubly reduced to colorless LMB.20 There are only a few reports of a reversible reaction involving MB in the acidic range,5 so we became interested in the oxidized form, which is quite stable and can be easily reduced back to MB in acidic pH. MB•+, i.e., the oxidized form, has a peak at 520 nm.20 Keeping this redox possibility in mind, we started our present work. First, 40 mL of 1 × 10−5 M MB solution was taken, and the pH of the solution was adjusted to ∼3.0 using dilute H2SO4 solution. At this low pH no color change of the MB solution as such occurs. This is confirmed by the absorption profile of MB in the visible region, while the aqueous solution goes from the neutral to the acidic range. This marks the stability of MB in the acidic pH range, and no peak appears around 520 nm. Thus, MB•+ formation does not take place in the reaction condition. As we failed to detect the MB•+ in acidic pH, we tempted to get MB•+ species with another

Figure 5. (a, left) Absorbance (A) vs wavelength (nm) and (b, right) absorbance (A) vs time (T). Conditions: [MB] = 10−5 M, 40 mL, pH ≈ 3 maintained by H2SO4, 0.02 g of Cu2O. 25745

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Figure 6. (a) Comparative account to show the degradation of MB. (b) Absorption spectra for successive oxidation of LMB by air.

Figure 7. (a) Absorbance (A) vs time (T) with a variable amount of Cu2O. (b) Rate constant vs dose of Cu2O. reduction

metrically following the decrease in the absorption maximum of MB at 663 nm. First, 40 mL of 1 × 10−5 M MB was acidified with dilute H2SO4, and the solution pH ≈ 3.0 was maintained. To it was introduced 0.02 g of solid Cu2O, and the solution was mixed well. Then the MB color bleaching started readily, and the absorbance of the solution was measured at an interval of 1 min with utmost care. With time the decrease in the absorbance maximum at 663 nm of MB is shown in Figure 5a. The plot of absorbance vs time shows a straight line indicating the zerothorder reaction pathway (Figure 5b), with the kinetics expressed by the equation

MB HoooooooooooI LMB oxidation Thus, a periodic oscillation of blue to colorless and again the regeneration of the blue color occurs. This demonstrates the clock reaction. This clock reaction continues over 50 times and remains active for several weeks upon allowing the reaction mixture to stand in the laboratory conditions. The reaction takes longer even for one cycle using bulk Cu2O, which does not play the same role as the as-prepared truncated Cu2O cubes. Again, CuO and Cu(0) do not participate in the redox reaction to exhibit the proposed clock reaction. The clock reaction has been carried out by taking 30 mL of 3 × 10−5 M MB and with a variable amount of Cu2O (0.005−0.02 g), keeping the solution pH ≈ 3.0 using dilute H2SO4, and the spectral changes are shown in Figure S5 (Supporting Information). It has been observed that for all these cases the kinetics follow the zeroth-order pathway (Figure 7a). The rate constant of each set is compared with the variable amount of added Cu2O sample (Figure 7b). Here it has been observed that, with an increased amount of Cu2O (truncated cubes), the rate constant increases (Figure 7b), but if the amount of Cu2O is decreased below 0.005 g, keeping all other parameters fixed, the clock reaction does not take place because there is an overlap of the forward and backward reactions due to the delayed forward reaction. Therefore, for this case the limiting dose of truncated Cu2O cubes is 0.005 g. Similarly, the effect of the concentration of MB was also observed just by varying the concentration of MB from 1× 10−5 to 4 × 10−5 M and keeping the amount of Cu2O and the concentration of H2SO4 unchanged. In all four cases the

[A]t = −kt + [A]0

where [A]t is absorbance at time t, [A]0 is initial absorbance of MB, and k denotes the rate constant. Using this equation, from the slope of the plot of A vs t, the rate constant “k” was calculated, which is 0.55163 mol L−1 s−1. After 3 min, the solution color almost faded out. Slight shaking of the reaction mixture quantitatively bleached the blue MB color, indicating the formation of LMB in the solution. The color bleaching is due to the formation of LMB, not the degradation of MB. This has been claimed from the UV−vis spectral profile (Figure 6a). The back reaction, i.e., LMB to MB, is stopped if the reaction is carried out in an inert atmosphere without any trace of O2. In the absence of Cu2O, there is no noticeable change in the absorbance of the MB solution. On standing, the faded solution regained its color with time and became blue again, indicating the transformation of LMB to MB by air. This visual observation can also be monitored spectrophotometrically (Figure 6b). 25746

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reaction follows the zeroth-order pathway (see Figure S6 Supporting Information). Actually in the clock reaction Cu2O acts as a redox-active substrate. Cu2O does not react with MB solution at normal pH, i.e., pH ≈ 7. The literature reports that, at 25 °C, the E° value of the MB/LMB couple is 0.011 V, whereas the E° value of the Cu2+/ Cu+ system is 0.15 V at pH ≈ 7.20 Therefore, thermodynamically the electrode potential does not allow the redox reaction to occur between Cu+ and MB at neutral pH, but when the pH of the solution is lowered to 3.0 in H2SO4, the E° value of the MB/LMB system is increased significantly, and at pH ≈ 1.0 this value becomes 0.532 V. Therefore, at pH 3.0 the redox potential of MB/LMB certainly becomes lower than 0.532 V and must be higher than that of the Cu2+/Cu+ system (0.15 V). Under this situation, the redox reaction between Cu+ and MB becomes thermodynamically allowed and color bleaching of MB is observed. MB is reduced to LMB, and simultaneously Cu+ is oxidized to Cu2+, which has been confirmed from FTIR spectroscopy (Figure S7, Supporting Information). Dilute H2SO4 does not react with Cu2O (it remains as uniform truncated Cu2O cubes in dilute H2SO4) under the experimental conditions, which is also confirmed from FTIR spectroscopy and TEM measurement, so here H2SO4 simply maintains the pH and prevents CuO formation. Thus, the necessary condition for the redox reaction is maintained. This clock reaction progressed in the presence of other acids such as HCl, acetic acid, etc. though HCl reacts with Cu2O (Figure S8, Supporting Information). Dilute H2SO4, being a redox-passive acid in the proposed reaction medium, becomes the best choice.

6. CONCLUSION In conclusion, we have reported a simple wet chemical method for the synthesis of uniform truncated Cu2O cubes at the gram level from microwave irradiation. A simple redox-active, crowdpleasing, demonstratable clock reaction has been put forward in dilute H2SO4 solution with a variable amount (5 × 10−3 to 2 × 10−2 g) of catalyst, keeping the MB concentration in the ∼10−5 M range without the need for any reducing agent whatsoever.

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ASSOCIATED CONTENT

S Supporting Information *

UV−vis spectra, FTIR spectra, FESEM images, EDS data, kinetic data changing the amount of Cu2O, kinetic data changing the amount of MB, FTIR spectrum of Cu2O at the time of progress of the clock reaction in the presence of H2SO4, and FTIR spectrum of pure CuCl2 and after the clock reaction using HCl. This material is available free of charge via the Internet at http:// pubs.acs.org/.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We are thankful to the University Grants Commission, Department of Science and Technology, North South Technologies, and Council of Science and Industrial Research, New Delhi, and the Indian Institute of Technology Kharagpur.

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REFERENCES

(1) Landolt, H. Ber. Dtsch. Chem. Ges. 1886, 19, 1317−1365. 25747

dx.doi.org/10.1021/jp308095h | J. Phys. Chem. C 2012, 116, 25741−25747