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Photocatalytic Degradation of Methylene Blue by Titanium Dioxide: Experimental and Modeling Study Chen Xu,† G. P. Rangaiah,*,† and X. S. Zhao‡ †

Department of Chemical & Biomolecular Engineering, National University of Singapore, 4 Engineering Drive, Singapore, 117585 School of Chemical Engineering, The University of Queensland, St Lucia, Brisbane, QLD 4072, Australia

‡

ABSTRACT: The application of semiconductors in water treatment via photocatalysis of various pollutants has attracted much attention from researchers. In this work, photocatalytic degradation of methylene blue by P25 titanium dioxide was studied experimentally and then via modeling. The effects of lamp choice, concentration of catalyst, and methylene blue were analyzed. Desorption of methylene blue at the start of light radiation was observed, and analyzed in detail for the first time. Both desorption and degradation processes were modeled, and experimental data was fitted to a pseudo-first-order model with sufficient accuracy. The effects of catalyst and initial dye concentration on reaction rate constants were discussed. There has been much research effort to find more effective photocatalysts as well as a more efficient reactor design.6 One possible development for photocatalytic water purification is to utilize sunlight as the light source.7 However, the band gap of TiO2 corresponds to the ultraviolet wavelength, which is only a small fraction of solar radiation. Therefore, semiconductors with smaller band gaps such as ZnS,8 CdS,9 and graphene and its derivatives,10,11 as well as addition of dopants, have been analyzed to utilize light of longer wavelength or enhance photocatalysis performance.12−14 Another approach to improving photocatalyst efficiency is to suppress the recombination of electron−hole pairs.15 Composite materials have also been studied to integrate merits of different materials, such as TiO2− silica,16 CuO/zeolite,17−20 and reduced graphene oxide−CdS− ZnO.21 Since the majority of photocatalytic processes occur close to the surface of the catalyst, the morphology of the catalyst is also crucial to its performance. Various methods have been proposed to synthesize catalysts of higher specific surface areas.22−24 Most photocatalysts are very small in size to obtain high specific surface areas; hence, it is energy intensive to remove them from the reactor suspension in a slurry reactor. Different reactor designs have been proposed to immobilize photocatalysts without compromising the overall performance.25−27 Quite a few contaminants have been analyzed for photocatalytic water treatment. Besides real-life wastewater, common target pollutants include aromatic compounds such as phenol,28,29 toluene,16 and chlorobenzene30 as well as dyes such as Rhodamine B,11 Acid Red 114,31 ethyl violet,32 and methylene blue.33−36 The dyes are of great interest as they are usually hard to degrade by conventional methods. Their remaining concentrations can also be easily determined by measuring the light absorbance of the reaction suspension or filtered samples. Environment acidity would affect the electric

1. INTRODUCTION Heterogeneous photocatalysis by semiconductors such as titanium dioxide (TiO2) is a promising technology for water purification.1 It can degrade water pollutants such as benzene, various dyes, and complex mixtures of water contaminants in industrial and domestic wastewaters.2,3 In addition, it is able to degrade many chemical contaminants and microorganisms completely into carbon dioxide, water, and mineral acid.4 Figure 1 shows the mechanism of photocatalysis. When the electron in the valence band of the semiconductor absorbs a

Figure 1. Mechanism of photocatalytic reaction.

photon with energy greater than the band gap (ΔE) of the semiconductor, the electron becomes excited and jumps to the conduction band, leaving a positively charged hole in the valence band. Besides recombination with the electron, the positively charged hole can oxidize water molecules to form hyper-reactive hydroxyl free radicals (•OH). The resulting hydroxyl radicals are the main agent that attack the chemical pollutant molecules or microorganism cells to purify water.5 The excited electron can react with dissolved oxygen molecule to form oxygen radical, which is also active toward organic pollutants. © 2014 American Chemical Society

Received: Revised: Accepted: Published: 14641

June 12, 2014 August 18, 2014 August 27, 2014 August 27, 2014 dx.doi.org/10.1021/ie502367x | Ind. Eng. Chem. Res. 2014, 53, 14641−14649

Industrial & Engineering Chemistry Research

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Equipment, Nanjing, China) for 30 min before its exposure to light radiation.39 The light absorbance of the reaction mixture reached equilibrium in about 20 min in the dark. Therefore, 30 min of light shielding and stirring are enough to reach adsorption−desorption equilibrium. Figure 2 shows the

charge on both the functional groups of dyes and the photocatalyst. Therefore, dyes with different functional groups would have different affinities to the photocatalyst in different pH environments; hence, their degradation kinetics would also be different.37 During the photocatalytic degradation process, the pH of the mixture is likely to change due to the formation of mineral acid, thus affecting the degradation kinetics. Photocatalytic degradation is a complex process, involving at least three phases (water, solid catalyst, light, and sometimes gas bubbled in to facilitate photocatalysis by dissolving oxygen to form peroxide radicals). The parameters deciding the overall reaction rate include the temperature; the solid catalyst particle size, morphology, and concentration; the target pollutant concentration and its ease of degradation; the water flow velocity and pattern; the light distribution field inside the reactor; the emission power and spectrum of the light source; the pH and the catalyst surface properties under that environment; and the water turbidity. In the literature, only a few of these parameters were studied in each paper while all other variables were held constant. In most cases, the degradation reaction fits into a Langmuir−Hinshelwood (L− H) model: dC kKC −r = = dt 1 + KC

Figure 2. Schematic of the photoreactor.

schematic of the photoreactor. Upon illumination, samples were taken at fixed time intervals and their light absorbance was measured immediately using a Shimadzu UV1601PC UV− visible spectrophotometer. The difference between the light absorbance at 600 nm of the drawn sample and the slurry at respective TiO2 concentrations was recorded. This wavelength was chosen because, during the light absorbance calibration, the change in light absorbance with respect to the methylene blue concentration is the most significant at 600 nm while the TiO2 concentration is held constant. The experimental run was stopped when the change in light absorbance was minimal. Each experiment was repeated three times, and the average values in light absorbance difference were recorded. The variation in light absorbance difference in the three runs at the same time under the same experimental conditions was within 5%, as shown by the small error bars in Figure 3.

(1)

where C is the target pollutant concentration, k is the reaction rate constant, and K is the adsorption equilibrium constant.16 When the concentration of the target pollutant is low, which is normally the case for degradation of dyes (a few dozens parts per million is common), KC ≪ 1, and the reaction rate can be simplified to a pseudo-first-order kinetic model with respect to the target pollutant concentration. This work studies experimentally and then models the photocatalytic degradation process of methylene blue by TiO2. Impacts of different lamps, initial methylene blue concentrations, and TiO2 concentrations on the photocatalytic process were examined. The experimental materials and methods are described in section 2. The adsorption−desorption process of methylene blue on TiO2 during photocatalytic degradation is discussed in section 3.1. This phenomenon is discussed in detail for the first time. The effects of reaction parameters are analyzed in sections 3.2−3.4. Modeling of the photocatalytic degradation of methylene blue is presented in section 4. Findings of this work are summarized in section 5.

2. EXPERIMENTAL MATERIALS AND METHODS Photocatalytic degradation was conducted with different lamps, initial methylene blue concentrations, and TiO2 concentrations. Degussa P25 TiO2 was used as the catalyst. It is polycrystalline and consists of 80 wt % anatase and 20 wt % rutile. Its average particle size is 25 nm, and its specific surface area is about 50 m2/g. It has been the research benchmark for photocatalysis.38 Methylene blue (Aldrich, used as purchased) was used as the target pollutant. It was dissolved in deionized water first to reach a concentration of 500 ppm for later use. The water used in this work was deionized by an Elga Micromeg water deionizer. A given amount of P25 TiO2 was added to deionized water. The mixture was ultrasonicated for 15 min at room temperature to disperse the solid catalyst particles before the addition of an appropriate amount of the 500 ppm methylene blue solution. The mixture was then stirred in darkness in a SGY-II B-Type Multifunctional Photochemistry Reactor (Stonetech Electric

Figure 3. Change in light absorbance over time when the radiation was shielded or not for the 0.2 g/L−20 ppm run; error bars are also shown.

In this work, different runs of photocatalytic degradation experiments were labeled as “TiO2 concentration−initial methylene blue concentration”. For example, the 0.2 g/L−20 ppm run means that the TiO2 concentration is 0.2 g/L while the initial methylene blue concentration is 20 ppm. For this particular run, 0.01 g of TiO2 was added into 48 mL of deionized water before the addition of 2 mL of 500 ppm methylene blue. Four TiO2 concentrations were chosen: 0.2, 0.3, 0.4, and 0.5 g/L. The initial methylene blue concentrations 14642

dx.doi.org/10.1021/ie502367x | Ind. Eng. Chem. Res. 2014, 53, 14641−14649


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a local maximum until all methylene blue molecules are degraded. 3.2. Effects of Different Lamps. Three different lampsa 350 W xenon lamp (labeled as “Xe350”), a 300 W mercury lamp, and a 500 W mercury lamp (labeled as “Hg300” and “Hg500”, respectively)were used as the light source in this work. The Hg lamps have the highest emission power at 365 nm, while the Xe lamp has the highest emission power around 850 nm. Figure 4 shows the degradation performance of these

were 20, 25, 30, and 35 ppm. The following three types of lamps were used: a 350 W xenon lamp, a 300 W mercury lamp, and a 500 W mercury lamp.

3. RESULTS AND DISCUSSION 3.1. Adsorption−Desorption Equilibrium during Degradation. Figure 3 shows the change in the light absorbance of the reaction suspension from the addition of methylene blue (time = −30 min) to the start of light radiation (time = 0 min) and the end of photocatalytic degradation. When the dye solution was added into the mixture, the dye molecules started to adsorb on the surface of the solid catalyst particles. No observable degradation occurred when there was no light radiation. The dye molecules absorb more light when they are free in the solution. Therefore, the adsorption process decreases the light absorbance of the mixture (region I). When the light is turned on, the light absorbance of the mixture at 600 nm increases with time first before decreasing to the value around the same as that of the TiO2−water slurry at the concentration of 0.2 g/L, as the change in light absorbance of the mixture is the most significant at 600 nm. Most work involving photocatalytic degradation of methylene blue did not report this phenomenon, and it has not been analyzed in detail. Therefore, to further study this unusual observation, the sampling time interval for several runs was reduced and the radiation was shielded for several times. Since the lamp requires several minutes to stabilize after it is turned on, the lamp was shielded to prevent light radiation reaching the slurry, instead of turning it off. It can be observed in Figure 3 that, during both increasing and decreasing the process of light absorbance, shielding the suspension from light radiation kept the light absorbance of the mixture at almost a constant level. This indicates that both increasing and decreasing of light absorbance are a result of light radiation. Decreasing light absorbance in the later part of the experiment is obviously due to degradation of methylene blue (region III). Since TiO2 is not able to degrade methylene blue without UV light radiation, shielding off the light stopped the light absorbance from decreasing. Increase in light absorbance at the start of light radiation is primarily due to desorption of methylene blue molecules from the surface of the solid catalyst particles. This desorption is probably due to the pH change in the slurry brought about by degradation products. However, since the methylene blue concentration is very low, the pH change is expected to be very small. Electric charges on both functional groups of methylene blue and the surface of TiO2 were altered, affecting the affinity between methylene blue molecules and the catalyst, shifting the adsorption equilibrium. The concentration of dissolved methylene blue molecules in the solution (i.e., not adsorbed on the surface of the catalyst particles) thus increases, increasing the light absorbance of the mixture. Therefore, when the radiation is shielded off, the light absorbance of the mixture did not fall back to the minimal value at adsorption equilibrium. At the same time, some methylene blue molecules are being degraded by the photocatalytic reaction when the radiation is on, decreasing the light absorbance. The overall light absorbance change is thus a competition between desorption (increase) and degradation (decrease). As the reaction goes on, the number of methylene blue molecules released from the catalyst surface decreases as less methylene blue molecules are now available for desorption. The light absorbance of the mixture thus starts to decrease after reaching

Figure 4. Degradation of methylene blue with different lamps: TiO2 concentration = 0.5 g/L and initial methylene blue concentration = 20 ppm.

lamps with respect to energy consumption. The energy consumption is calculated as the electric energy consumed by the lamps, i.e., the product of the power of the lamps and irradiation time. The TiO2 concentration is 0.5 g/L, and the initial methylene blue concentration is 20 ppm. It can be observed that degradation using the xenon lamp is the least energy efficient among the three lamps tested (Figure 4). This is mainly due to the lower proportion of UV radiation in the xenon lamp emission spectrum than in mercury lamp emission spectra. Therefore, a larger amount of electric energy is “wasted” as light radiation of longer wavelength which is not able to overcome the band gap energy of TiO2. When comparing the Hg300 and Hg500 data, it can be observed that the 300 W mercury lamp is able to degrade more methylene blue with the same amount of energy consumed. It can then be deduced that, for well-mixed batch reactors, a larger number of smaller reactors with lamps of lower power is preferred to a smaller number of larger reactors with lamps of higher power, if all other operating conditions are kept constant. For continuous reactors, lower water flow rate and lower UV lamp power are more favorable while keeping the UV dosage constant (amount of UV radiation received per unit volume of water). Our previous work has also revealed that a lower water flow rate is more favorable in terms of higher UV dosage received for the suspended particles in water.40 Hence, a low water flow rate is preferred in degradation performance. However, a lower water flow rate means a higher residence time and, hence, a higher number of reactors required and larger plant size for a certain daily treatment capacity, increasing the capital investment. Hence, there is a trade-off between capital and operating costs for optimal design. 3.3. TiO2 Concentration. Photocatalytic degradation of methylene blue was carried out with different TiO2 and initial methylene blue concentrations. Figure 5 shows the degradation performance under a 500 W mercury lamp of different TiO2 concentrations when the initial methylene blue concentration was varied in the range from 20 to 35 ppm. At the initial 14643



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Figure 5. Degradation of methylene blue with different TiO2 concentrations, a 500 W mercury lamp, and initial methylene blue concentrations of (a) 20, (b) 25, (c) 30, and (d) 35 ppm.

change is more predominant when the methylene blue concentration is lower. It should also be noted that, for each initial methylene blue concentration, the increase in the light absorbance during the first few minutes of light radiation becomes smaller and even is not observed with increasing TiO2 concentration. For example, in Figure 5a, the light absorbance keeps decreasing with time for the 0.4 g/L−20 ppm and 0.5 g/L−20 ppm runs. The difference between the initial light absorbance and the highest light absorbance is about 0.2 for the 0.2 g/L−20 ppm run and about 0.1 for the 0.3 g/L−20 ppm run. As previously discussed in section 3.1, the increase in light absorbance during the first few minutes is due to desorption of methylene blue molecules from the surface of the solid catalyst particles. This effect is more dominant for lower TiO2 concentration as higher TiO2 concentration offers more sites for adsorption. Hence, the change in adsorption−desorption equilibrium due to light radiation does not effectively increase the number of “free” methylene blue molecules in the solution when the TiO2 concentration is high enough. On the other hand, the methylene blue molecules in the solution were being degraded all the time, lowering the light absorbance. Therefore, this desorption effect is not observable in terms of increase in light absorbance at higher TiO2 concentration, such as the 0.4 g/L− 20 ppm and 0.5 g/L−20 ppm runs. As the initial methylene blue concentration increases, more adsorption sites on the solid catalyst particles are occupied by the methylene blue molecules. When light radiation is on, the amount of desorbed methylene blue molecules is thus higher. Hence, the TiO2 concentration for which such a desorption effect becomes observable increases with increasing initial methylene blue concentration. When the initial methylene blue concentration is 20 ppm (Figure 5a), this effect becomes observable when the TiO2 concentration is 0.3 g/L or lower. In Figure 5b (25 ppm), such a desorption effect is observed when the TiO2 concentration is 0.4 g/L or lower. When the initial methylene blue

methylene blue concentration of 20 ppm, the degradation speed is in this order of TiO2 concentration (g/L): 0.3 > 0.4 > 0.2 > 0.5 (Figure 5a). For other initial methylene blue concentrations, this trend is a bit different. For initial methylene blue concentrations of 25 and 30 ppm, a TiO2 concentration of 0.3 g/L degrades the dye at the highest speed, followed by 0.4 and 0.5 g/L, and 0.2 g/L is the slowest (Figure 5b,c). When the initial methylene blue concentration is 35 ppm, the degradation speed is in this order of TiO2 concentration (g/L): 0.4 > 0.3 > 0.5 > 0.2 (Figure 5d). These observations are because, on one hand, higher TiO2 concentration offers more reaction sites for oxidation of water molecules and production of hydroxyl radicals, thus increasing the reaction rate. On the other hand, TiO2 also increases the light absorbance of the mixture, lowering the average light radiation and the total amount of photons received by the photocatalyst. The production rate of electron−hole pairs is lowered, and thus fewer hydroxyl radicals are produced, decreasing the reaction rate. At lower TiO2 concentration, the increasing effect dominates, while at higher TiO2 concentration the latter decreasing effect plays a more important role. Therefore, the degradation speed would first increase and then decrease with increasing TiO2 concentration. The optimal TiO2 concentration is also different with different initial methylene blue concentrations. At lower initial methylene blue concentration, the optimal TiO2 concentration is lower. The concentration of 0.3 g/L offers the fastest degradation for initial methylene blue concentrations of 20−30 ppm, while for 35 ppm 0.4 g/L is the fastest. This is because, at lower methylene blue concentration, the contribution of light absorbance by TiO2 is relatively higher; i.e., the light absorbance is more sensitive to an increase in TiO 2 concentration. Increasing the TiO2 concentration (for example, from 0.2 to 0.3 g/L) brings more relative increase in light absorbance for a reaction mixture of lower methylene blue concentration. Therefore, the effect of TiO2 concentration 14644



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Figure 6. Degradation of methylene blue with different initial methylene blue concentrations, 500 W mercury lamp, and TiO2 concentrations in different plots: (a) 0.2, (b) 0.3, (c) 0.4, and (d) 0.5 g/L.

such increase is observed in the 0.5 g/L−20 ppm and 0.5 g/L− 25 ppm runs, while the increase is quite obvious in the 0.5 g/ L−30 ppm and 0.5 g/L−35 ppm runs. As previously discussed in section 3.1, the increase in light absorbance is primarily due to desorption of methylene blue from the surface of the solid catalyst particles. Section 3.3 has also discussed this phenomenon at higher TiO2 concentration and lower methylene blue concentration. When the initial methylene blue concentration increases, the change in adsorption− desorption equilibrium would increase the concentration of dissolved methylene blue in the solution. Hence, the light absorbance increases. On the other hand, an increase in initial methylene blue concentration also makes a decrease in the light absorbance faster as the degradation reaction rate increases with increasing methylene blue concentration. Therefore, this partially offsets the desorption effect when compared to the increase in light absorbance during the first few minutes of 0.5 g/L−30 ppm and 0.5 g/L−35 ppm runs. The magnitude of this increase in light absorbance is larger in the 0.5 g/L−30 ppm run (about 0.15) than in the 0.5 g/L−35 ppm run (about 0.11). As the TiO2 concentration decreases, the number of vacant adsorption sites decreases. Hence, when there is a shift of equilibrium to the desorption side, the desorption effect is more observable at lower TiO2 concentration. The initial methylene blue concentration for which such an effect becomes observable thus decreases with decreasing TiO2 concentration. When the TiO2 concentration is 0.4 g/L, this desorption effect is observed when the initial methylene blue concentration is 25 ppm or higher (Figure 6c). When the TiO2 concentration drops to 0.3and 0.2 g/L, this desorption effect is observed for all initial methylene blue concentrations (Figure 6a,b).

concentration is 30 ppm or higher, this desorption effect is observed for all TiO2 concentrations (Figure 5c,d). 3.4. Initial Methylene Blue Concentration. Figure 6 shows the photocatalytic degradation of different initial methylene blue concentrations with different TiO2 concentrations. Figure 6 is a replotting of the results in Figure 5 to group them according to TiO2 concentration, so that the effect of methylene blue concentration on both degradation and desorption can be observed directly. The time difference decreases with increasing TiO2 concentration to completely degrade different initial amounts of methylene blue. It is more than 1 h for a TiO2 concentration of 0.2 g/L (Figure 6a) and around 40 min for a TiO2 concentration of 0.5 g/L (Figure 6d). This is because, at higher TiO2 concentration, more surface adsorption sites are available. The reaction rate is thus less sensitive to the increase in light absorbance brought about by an increase in the initial methylene blue concentration. As the initial methylene blue concentration increases, the slope of the degradation curves becomes less steep when approaching the end of the reaction; i.e., it took longer for reaction runs with a higher initial methylene blue concentration to degrade for the same amount of light absorbance. That is, the degradation reaction is slower for higher methylene blue concentrations. This is because the intermediate products of the degradation of methylene blue, which have lower light absorbance, would compete with methylene blue for reaction with hydroxyl radicals. At higher initial methylene blue concentration, the intermediates’ concentrations are also higher, which lowers the methylene blue degradation rate. The above observation is consistent with the decrease in photocatalytic degradation efficiency with increase in target pollutant concentration as reported in the literature.41 Another notable observation from Figure 6 is that, as the initial methylene blue concentration increases, the increase in light absorbance during the first few minutes is different. This is especially obvious for the 0.5 g/L runs. Figure 6d shows that no

4. MODELING OF DEGRADATION AND DESORPTION As discussed in section 3, the change in light absorbance of the mixture can be attributed to two processes: desorption increases light absorbance and degradation decreases it. 14645



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received by the catalyst, so k1 and k2 are dependent on their concentrations. Assume that the light absorbance (X) is linear with respect to both Ca and Cf:

Therefore, to model the change in light absorbance, both of these have to be considered. The degradation reaction can usually be fitted into eq 1. As the methylene blue concentration is low in this work (35 ppm maximum), it can further be simplified since KC ≪ 1. dC kKC −r = = ≈ kKC dt 1 + KC

X = αCf + βCaCs ⎞ ⎛ αk 2 Cs + βCs⎟Ca0 exp( −k 2t ) =⎜ ⎠ ⎝ k1 − k 2 ⎛ k 2CsCa0 ⎞ + α⎜Cf0 − ⎟ exp( −k1t ) k1 − k 2 ⎠ ⎝

(2)

where C is the methylene blue concentration, k is the degradation reaction rate constant, and K is the adsorption equilibrium constant. The L−H model is based on the assumption that the adsorption equilibrium is established much faster than the reaction. From experimental observation, it was found that the adsorption equilibrium is continuously shifting toward the desorption side when light radiation is on; i.e., the proportion of adsorbed methylene blue is always decreasing (section 3.1). Therefore, the L−H model is no longer applicable here. Hence, for simplicity, it is assumed that the net desorption rate is first order with respect to the concentration of methylene blue adsorbed on the surface of the catalyst; thus

dCa = −k 2Ca dt

(7)

= m exp( −k1t ) + n exp( −k 2t )

where ⎛ k CC ⎞ m = α⎜Cf0 − 2 s a0 ⎟ k1 − k 2 ⎠ ⎝

⎞ ⎛ αk 2 n=⎜ + β ⎟Ca0Cs ⎠ ⎝ k1 − k 2

This assumption about light absorbance and methylene blue concentration can be partially validated as the light absorbance changes nearly linearly with respect to the total methylene concentration, as shown in Figure 7.

(3)

where Ca is the concentration of methylene blue adsorbed on the surface of the catalyst (in mass per unit mass of catalyst), and k2 is the desorption rate constant (in hour−1). Therefore, the concentration of methylene blue adsorbed on the catalyst surface is Ca = Ca0 exp( −k 2t )

(4)

where Ca0 is the concentration of adsorbed methylene blue at the start of light radiation. The change in the concentration of “free” methylene blue in the solution because of both degradation and desorption processes is dCa dCf 1 = −r − = [−k1Cf V + k 2CaCsV ] dt dt V

Figure 7. Variation of light absorbance of 0.5 g/L TiO2 with methylene blue concentration.

(5)

The light absorbance curves in Figure 5 indicate that the free methylene blue concentration first increases and then decreases over time for most of the runs; hence, the desorption rate is higher than the degradation rate under most conditions at the beginning of light radiation. Therefore, k2 > k1, and m is positive. For the same reason, n has to be negative in order to fit the shape of the light absorbance curve with respect to time. Measured light absorbance data from each run of all of the 16 photocatalytic degradation runs using the 500 W mercury lamp were fitted into eq 7. The estimates of m, n, k1, and k2 from this modeling for all runs are summarized in Table 1. Slope and R2 in Table 1 refer to the gradients and the coefficients of determination for the regression line passing through the origin of the plots of fitted data versus actual light absorbance (i.e., parity plots) for each run. All slopes and R2 are close to unity, indicating a good fitting of the experimental data into eq 7. Figure 8 shows the modeling results for selected runs, which were also close to the experimental results. Figures 9 and 10 are, respectively, the plots of k1 and k2 at different TiO2 and initial methylene blue concentrations. It can be observed from Figures 9a and 10a that both the degradation rate (k1) and desorption rate (k2) constants decrease with increasing initial methylene blue concentration. This is because higher methylene blue concentration absorbs more UV

Here, r is the degradation reaction rate, Cf is the concentration of “free” methylene blue in the solution (in grams per liter), k1 is the degradation reaction rate constant (in hour−1), V is the total solution volume (in liters), and Cs is the catalyst concentration (in grams per liter). The second term in eq 5 is the net change in concentration of adsorbed methylene blue, i.e., the rate of adsorption minus the rate of desorption. Since both the adsorption−desorption equilibrium shift and degradation affect the light absorbance at the same time, eq 5 is comprised of two terms: the first term describing the effect of degradation and the second term for adsorption−desorption equilibrium shift. Solving eq 5 k2 CsCa0 exp(−k 2t ) k1 − k 2 ⎛ k CC ⎞ + ⎜Cf0 − 2 s a0 ⎟ exp( −k1t ) k1 − k 2 ⎠ ⎝

Cf =

(6)

where Cf0 is the concentration of free methylene blue in the solution at the start of light radiation. It should be noted that both k1 and k2 are dependent on light radiation received by the catalyst. Both TiO2 and methylene blue affect light radiation 14646



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Table 1. Constants for Photocatalytic Degradation and Desorption run 0.2 0.2 0.2 0.2 0.3 0.3 0.3 0.3 0.4 0.4 0.4 0.4 0.5 0.5 0.5 0.5

g/L−20 g/L−25 g/L−30 g/L−35 g/L−20 g/L−25 g/L−30 g/L−35 g/L−20 g/L−25 g/L−30 g/L−35 g/L−20 g/L−25 g/L−30 g/L−35

ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm ppm

m

n

k1 (h−1)

k2 (h−1)

slope

R2

no. data pts

2.152 3.202 2.743 4.546 2.201 2.959 2.917 4.866 2.347 2.986 2.552 4.926 1.701 2.910 4.793 5.085

−2.046 −2.903 −2.351 −4.033 −1.990 −2.792 −2.571 −4.460 −1.896 −2.606 −2.188 −4.427 −1.134 −2.223 −4.166 −4.176

6.905 4.489 2.826 2.189 6.925 5.619 3.782 2.965 6.941 5.306 3.429 3.168 4.407 4.654 3.664 2.818

10.036 6.499 4.997 3.343 9.211 7.800 6.252 4.267 9.033 7.265 6.018 4.491 6.970 6.455 4.927 3.957

1.0190 0.9969 0.9992 1.0021 0.9929 0.9976 1.0048 0.9939 0.9979 1.0068 1.0074 0.9963 1.0222 1.0097 1.0085 0.9979

0.9844 0.9697 0.9854 0.9894 0.9900 0.9949 0.9906 0.9908 0.9956 0.9976 0.9960 0.9965 0.9903 0.9923 0.9946 0.9946

8 11 17 25 8 11 12 20 9 11 16 17 14 12 15 19

higher TiO2 concentration. This is because higher TiO2 concentration provides more sites for adsorption of water molecules and hence more hydroxyl radicals are produced, increasing the degradation rate. On the other hand, higher TiO2 concentration also absorbs more UV radiation, decreasing the overall photon efficiency and hence decreasing the degradation rate. As discussed in section 3.3, the increasing effect is more dominant at lower TiO2 concentration while the decreasing effect is more dominant at higher TiO2 concentration. The desorption rate constant also decreases with increasing initial methylene blue concentration (Figure 10a). The rate of change of the solid catalyst surface property increases with increasing TiO2 concentration. On the other hand, higher TiO2 concentration provides more sites for adsorption, shifting the equilibrium to the adsorption side, effectively lowering the desorption rate. This decreasing effect is most observable when the initial methylene blue concentration is 20 ppm, due to the fact that majority of the methylene blue molecules are adsorbed on the surface of the catalyst at this low methylene blue concentration. More UV radiation absorbed by higher TiO2 concentration also contributes to lower the desorption rate constant. It can be concluded from experimental data and modeling that both degradation and desorption of methylene blue by TiO2 fit into a pseudo-first-order model well. High methylene blue concentration suppresses both desorption and degradation, and the effect of TiO2 concentration is more complicated.

Figure 8. Light absorbance for selected runs and model predictions.

radiation, thus lowering the overall UV radiation received by the solid catalyst particles, effectively slowing the production rate of hydroxyl radicals and the degradation reaction. This observation is consistent with the research reported in previous work.42,43 Since desorption is also related to the UV radiation, blocking of UV radiation by higher methylene blue concentration decreases the desorption rate constant. Figure 9b shows that the degradation rate constant first increases and then decreases with catalyst concentration, at higher initial methylene blue concentrations (25 ppm and above). On the other hand, when the initial methylene blue concentration is 20 ppm, the degradation rate constant is almost constant when the TiO2 concentration is below 0.4 g/L and then decreases at

Figure 9. Degradation rate constants at different TiO2 and initial methylene blue concentrations. 14647



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Figure 10. Desorption rate constants at different TiO2 and initial methylene blue concentrations.

Increasing the TiO2 concentration enhances both desorption and degradation at lower values. Further increasing it reduces reaction rate constants for both processes. Both of these are mainly due to the change in light radiation received by the catalyst under different experimental conditions.

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5. CONCLUSIONS Photocatalytic degradation of methylene blue by TiO2 was carried out under different conditions. Desorption of methylene blue molecules from the catalyst surface at the start of UV radiation, probably due to a change in the surface property of the solid catalyst, was observed. The lamp with lower power was found to be more energy efficient. Experimental data show that the concentration of TiO2 and the initial methylene blue concentration have a complex impact on the reaction rate. Both net desorption and degradation of methylene blue fit the pseudo-first-order reaction model well. Methylene blue has a negative impact on both degradation and desorption. Increasing the TiO2 concentration first enhances both processes and then suppresses them.

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V = total solution volume (L) X = light absorbance of reaction suspension

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: (65)-6516 2187. Fax: (65)6779 1936. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors acknowledge Dr. Xiong Zhigang of the University of Queensland and Dr. Cai Zhongyu of the University of Pittsburgh for discussions on the experimentation and data analysis reported in this article.

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NOTATION Ca = concentration of methylene blue adsorbed on the surface of the catalyst (g/g of catalyst) Ca0 = concentration of adsorbed methylene blue at the start of light radiation (g/g of catalyst) Cf = concentration of free methylene blue in solution (g/L) Cf0 = concentration of free methylene blue in solution at the start of light radiation (g/L) Cs = catalyst concentration (g/L) k1 = degradation reaction rate constant (h−1) k2 = desorption rate constant (h−1) r = degradation reaction rate (g/L·h) t = time (s) 14648



Article

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