J. Chem. SOC., Furaday Trans. I, 1989, 85(6), 1291-1302 Photocatalytic Oxidation and Adsorption of Methylene Blue on Thin Films of Near-ultraviolet-illuminated TiO, Ralph W. Matthews CSIRO Division of Fuel Technology, Lucas Heights Research Laboratories, Private Mail Bag 7, Menai, NSW 2234, Australia Aqueous solutions of methylene blue are totally mineralized when recirculated over thin films of titanium dioxide illuminated with near-u.v. light. The rate of destruction obeys first-order kinetics with reasonable precision but the apparent first-order rate constant, k', decreases with increasing initial concentration of solute. 1 dm3 of 10 pmol dmP3 methylene blue solution illuminated with a 20 W lamp, is decreased to 5 pmol dm-3 in 1 1.8 min. Sunlight from a 1 m2 parabolic trough is capable of destroying the methylene blue at 6.4 times this rate. The decrease in k' values with increasing concentration is consistent with curves calculated using the integrated form of the Langmuir expression.The adsorption parameter determined in the analysis of the kinetics of the photocatalytic data agreed with the adsorption affinity parameter determined using the classical Langmuir adsorption isotherm for the dark equilibrium data ; this indicated the key role played by adsorption in photocatalytic oxidation with titanium dioxide. The maximum quantum yield for methylene blue destruction at high flow rates with a 10 pmol dm-3 initial concentration was calculated to be 0.0092. Introduction Photocatalytic oxidation is a comparatively recent technique for the destructive removal of organic impurities from water.It is receiving increasing attention as a means of water purification. '-17 Titanium dioxide powder suspensions illuminated with bandgap light have generally been used as the photocatalyst, but supported TiO, may also be used as a stationary phase.'8-20 Since the photo-oxidation reaction takes place at the surface of the photocatalyst, the adsorption characteristics of the solute are expected to be quite important and experimental evidence that supports this view21 has been obtained for a number of compounds. In some cases, compounds that are normally quite resistant to oxidation were oxidized in dilute solutions more rapidly than less-resistant compounds at the same concentrations, apparently because of the strong adsorption of the former on the reactive surface of the catalyst.This was not, however, a definitive explanation since the parameter indicating the degree of adsorption was obtained from an analysis of the data in terms of Langmuir-Hinshelwood kinetics and not from equilibrium studies. An attempt is made to relate Langmuir parameters from equilibrium experiments to parameters obtained from kinetic studies. Experimental Materials Degussa P25 titanium dioxide was used as a photocatalyst. This material is mostly in the anatase form, and has a B.E.T. surface area of 50 m2 g-l and a mean particle size of 12911292 Adsorption of Methylene Blue on Thin Films 30 nm. The photocatalyst was attached in a thin film to the inner wall of borosilicate glass tubing wound in a ~pira1.l~ Methylene blue laboratory reagent used as the solute was obtained from Aldrich.It was selected because it has strong adsorption characteristics on many surfaces, good resistance to light degradation, a well defined optical absorption maximum at 660 nm and is a common dye. Salicylic acid analytical reagent was from BDH. Solutions were prepared using water from a Millipore Waters Milli Q water purification unit. Perchloric acid was used when pH adjustment of solutions was required. Apparatus The borosilicate glass spiral was wound from 7 m of 6 mm 0.d. tubing in 65 turns so that it fitted closely around a 20 W NEC blacklight blue fluorescent tube (TlO), 32.5 mm diameter, 588 mm long. The spiral and fluorescent tube could be mounted in a standard domestic lamp holder.In some experiments, the spiral was mounted at the focus of a small parabolic trough. The inner surface of the spiral was coated with a layer of ca. 75 mg of TiO, as described elsewhere.,O Solutions were pumped through the spiral using a Cole-Parmer peristaltic pump. Unless otherwise stated, solutions were unbuffered and at their natural pH. pH measurements were made with an Orion model 801 meter. Analysis Changes in methylene blue concentration were measured spectrophotometrically using a Cary model 16 instrument. Calibration curves were prepared using solutions of known concentration with an analytical wavelength of 660 nm. For concentrations up to 10 pmoldm-3 the Beer-Lambert law was obeyed with good precision and the molar extinction coefficient was found to be 66 700 & 350 dm3 cm-l mol-l.Changes in salicylic acid concentration were determined by spectrophotofluorimetry using a Perkin-Elmer LS-5 luminescence spectrometer. The excitation and emission wavelengths were 294 and 41 2 nm, respectively. Carbon dioxide analyses were done by conductivity detection using the apparatus and method described elsewhere. ( b ) Procedure Generally, 500 cm3 of solution was pumped from a reservoir (an 800 cm3 beaker open to the atmosphere) through the spiral and back to the reservoir continuously at a flow rate of 100 cm3 min-'. In some experiments different flow rates and volumes were used. The solution in the reservoir was sampled every 2 min and returned to the reservoir after measurement.After 10 min the lamp was switched on. (In the solar experiments a black cloth was removed.) Between runs, water was circulated through the spiral, with the lamp on for ca. 1 h, to oxidize traces of adsorbed decomposition products photocatalytically. Separate experiments were run for the CO, determinations using 40 cm3 of solution at each concentration. The apparatus allowed continuous monitoring of the CO, formation during illumination. Results and Discussion Solute Concentration The effect of initial solute concentration on the rate of disappearance of methylene blue from the reservoir is shown in fig. 1. At all concentrations there was a small, rapid decrease in concentration that occurred in the dark, plateauing to a constant value after ca. 10min. When the lamp was switched on, a further decrease in concentration occurred that followed approximately first-order kinetics after 2 min of illumination.R.W. Matthews 1293 time/min Fig. 1. Disappearance of methylene blue from solutions recirculated through the TiO, photoreactor. Solution volume: 500 cm3, flow rate 100 cm3 min-', 20 W NEC T10 blacklight fluorescent tube. The Dark Reaction The small rapid decrease in methylene blue concentration that occurred on circulating the solution through the spiral without the light on is attributed to adsorption of methylene blue on the surface of the TiO,. When the change in concentration during this 10min period was plotted against time the results in fig. 2 were obtained. The concentrations at 10min were taken to be equilibrium concentrations and the concentration of methylene blue adsorbed on the surface was calculated from the difference between the initial and equilibrium concentrations.The plot of methylene blue adsorbed us. the equilibrium concentration shown in fig. 3 has the shape of a typical Langmuir adsorption isotherm and may be described by the expression : where [MB],,, is the concentration of methylene blue adsorbed, [MB] is the equilibrium concentration of methylene blue in solution, and k , , k , are constants for the given system, k , being related to the adsorption affinity and k, being related to saturation coverage of the surface. Eqn (1) may be rearranged to the standard reciprocal form and reciprocal values of [MB],,, plotted against reciprocal [MB] values (fig.3 insert) or solved by the method of non-linear least-squares to give the Langmuir parameters. The latter method was used to obtain values for k, and k, of 0.0298 0.0064 dm3 pmol-' and 44 FAR I1294 Adsorption of Methylene Blue on Thin Films 0.0 0 2 4 6 8 10 Fig. 2. Adsorption of methylene blue from solutions of different concentration onto the non- illuminated, supported TiO, of the photoreactor. Initial concentrations (pmol dm-3) : 0, 1 ; a, 2; 8, 5 ; (>, 10; 0, 20; 0, 50; 0, 100. timelmin 5.0850.45 pmol dm-3, respectively. The curve shown in fig. 3 was obtained by substitution of these values in eqn (1). The Photocatalytic Reaction The discontinuity between the rate of decrease in concentration during the first 2 min after the lamp was switched on and the rest of the illumination time was presumably due to the lamp warming up to full power.Regression analysis for the linear portion of the results gave the apparent first-order decay constants, k , listed in table 1. The approximate obedience to first-order kinetics, together with a decrease in the apparent rate constant with increasing initial concentration, has been noted for other solutes.11. 1 9 9 2o In one of those reports2' the concentration dependence of the apparent rate constant was explained in terms of the integrated form of the Langmuir adsorption isotherm : t = -ln--+-([S]o-[S]) where t is the time in minutes for the initial concentration of solute, [S]', to decrease to [S] and k,, K are constants for the given system related to the adsorption and reaction properties of the solute.It follows from eqn (2) that at (2) 1 [S]O 1 klK [SI K when [S]/[S]' = 0.5, (3) 0.693 0.5[SI0 0.693 +- K k , K ' - -- '0.5 = ~ k' Therefore, if eqn (2) is a reasonable approximation to the data, a plot of to.5 values against the initial concentration values should yield a straight line whose slope is 0 S / K andR. W. Matthews I I I I I 1295 0 0 20 40 60 80 100 [methylene blue]/mol dm-3 x lo6 Fig. 3. Methylene blue adsorbed versus equilibrium concentrations, from data of fig. 2. Insert : Plot of reciprocal concentrations. The curve was calculated by the method of least squares. Table 1. Apparent first-order decay constants for the photocatalytic destruction of methylene blue at different initial concentrations 1 0.0830 f 0.0005 8.35 2 0.0785 f 0.0005 8.83 5 0.0728 f 0.000 1 9.52 10 0.0700 0.0002 9.90 20 0.0645 f 0.0005 10.7 50 0.0453 f 0.0005 15.3 100 0.0286 f 0.0009 24.2 whose intercept is 0.693/(k1K).This plot, shown in fig. 4, gave the values of 0.0258 & 0.001 3 dm-3 pmol for k , and 3.22 0.09 pmol min-l dmP3 for K . It is noted that the value for k , obtained from these kinetic data agrees, within experimental error, with k, obtained from the equilibrium data, and supports the view that the k,.parameter obtained from the photocatalytic destruction rates is a genuine reflection of the adsorption characteristics of the molecule. The numbers obtained for k, and K were substituted in eqn (2) for different values of the initial methylene blue concentration to calculate the curves passing through the data points of fig.1. The agreement between the k, values in the two analyses begs the question in regard 4421296 Adsorption of Methylene Blue on Thin Films I I I I I ' I 24 - 22 - 20 - 18 - 4 \ $ 16 - 14 - 12 - 10 - [methylene blue] /mol dm-3 x 1 O6 Fig. 4. Half-life of photocatalytic destruction rate us. initial methylene blue concentration from data of fig. 1. to the relationship between k , and K . The k, parameter is related to the maximum coverage at the surface at equilibrium; K depends on the reaction rate of molecules adsorbed at the surface. It is suggested that K = k2q5NT, (4) where q5 is the quantum yield for the destruction of the solute, N is the total number of photons absorbed by the photocatalyst, and is the rate of transport of solute molecules to the surface.Flow Rate It was previously observed that increasing the flow rate through the illuminated spiral markedly increased the photocatalytic oxidation rate of several solutes. 2o Fig. 5 shows the effect of increasing the flow rate with methylene blue as the solute. A similar marked increase in the rate of disappearance with increasing flow rate also occurred. The apparent first-order decay constants obtained by regression analysis of the linear portion of the curves are given in table 2 together with the corresponding values calculated for 500 cm3 of solution. These calculated apparent first-order constants are shown plotted against flow rate in fig. 6 . It is noted that the increased k' values appear to be approaching a limiting value at high flow rates suggesting that the relationship may be described by an expression of the form where k* is the limiting value of k' at high flow rates, FR is the flow rate, and /3 is a constant for the system.The line drawn through the experimental points was obtainedR. W. Matthews 1297 1 I I I I illumination time/min Fig. 5. Disappearance of methylene blue from 500 cm3, 10 pmol dmV3 solutions recirculated through the TiO, photoreactor at different flow rates (cm3 min-I): 0 , 3 0 ; a, 60; @,90; a, 120; 0, 180; 0, 240. (20 W lamp.) Table 2. Effect of flow rate on apparent first-order decay constant of methylene blue. 200 cm3 of 10 pmol dm-3 solution, 20 W lamp flow rate /cm3 min-’ k’lmin-l k’lmin-l t,.,a/min 30 0.090 f 0.0009 0.036 19.2 60 0.152 & 0.0008 0.061 11.4 90 0.173 f0.0006 0.069 10.0 120 0.2 1 9 f 0.00 1 6 0.088 7.9 180 0.236 f 0.0021 0.094 7.4 240 0.275 f 0.0006 0.1 10 6.3 300 0.292 & 0.0004 0.1 17 5.9 a Calculated for 500 cm3 of 10pmol dm-3 solution.It was confirmed experimentally that k’ was directly proportional to the ratio of solution volumes. from eqn ( 5 ) by the method of least squares, with numerical values of 0.0101 +0.0012 min ml-l and 0.154 0.007 min-’ for /3 and k*, respectively. It is evident from eqn (2) and the apparent first-order decay approximation that, at low initial concentrations of PI0 P I solute, In- x k, Kt x k‘t. The approximate equivalence of k’ and k , K has been noted elsewhere.22 Therefore,1298 Adsorption of Methylene Blue on Thin Films 0.12 0 .lo 0.08 - I .s E 0.06 0.04 0.02 0.00 I I I I I I 100 2 00 300 flow rate/cm3 min-I Fig.6. Apparent first-order destruction rate us. flow rate. Data from fig. 5 +data for 300 cm3 min-'. k, K z k' or k* at high flow rates. The limiting value of K for methylene blue at low concentrations and high flow rates is therefore 5.97 f 0.56 pmol dm-3 min-l (0.154/0.0258). One notable difference between the present and previous results with salicylic acid20 is the absence of any downward curvatures in the semilog plots in fig. 5 . A pronounced downward curvature occurred when salicylic acid was the solute, especially at high flow rates. Those curves are well described by an expression of the same form as eqn (2). The reason for the absence of downward curvature with methylene blue is probably related to the adsorbed intermediate decomposition products competing more strongly with methylene blue for the primary oxidizing species than in the case of salicylic acid.Quantum Yield Actinometry on the lamp and spiral was done by circulating 0.006 mol dmP3 potassium ferrioxalate through the uncoated spiral from the reservoir. The analytical method of Hatchard and Parker23 was used to measure the rate of ferrous ion formation, 833 pmol min-l dm-3. According to specifications given by the lamp manufacturer a broad band of radiation is emitted with a maximum at 350 nm; very little radiation is emitted at a wavelength greater than 400 nm (ca. 5%). The quantum yield $(Fe2+) was therefore taken to be 1.2824 and the quantum yield for methylene blue disappearance was calculated to be 0.0092 for the limiting value k*, assuming that the same number of photons were absorbed by the TiO, layer as by the ferrioxalate solution.Solar Illumination The disappearance of methylene blue from 500 cm3 of an initially 10 pmol dm-3 solution with solar illumination in the parabolic trough is shown in fig. 7. The solution wasR. W. Matthews 1299 illumination time/min Fig. 7. Comparison of 20 W fluorescent lamp and solar illumination for methylene blue disappearance and salicylic acid disappearance with 20 W illumination. Flow rate 100 cm3 min-'. 500 cm3 of solution. (>, methylene blue, 20 W lamp, no TiO, in spiral; 0, methylene blue, 20 W lamp; a, methylene blue, 0.25 m2 parabolic trough, sunlight; 0 , salicylic acid, 20 W lamp.circulated through the spiral at 100 cm3 min-'. The k' value from the first-order plot was 0.109+0.002 min-' and the corresponding to.5 value was 6.35 min. For the purposes of comparison, the first-order plot for the disappearance under the same conditions is also shown but using the 20 W U.V. fluorescent tube as the illumination source. In this case the k' value was 0.068 k0.002 min-l and the 20.5 value, 10.2 min. Since the effective area of the parabolic trough illumination was 0.25 m2, it follows that illumination from 1 m2 at the same intensity at the surface of the TiO, would be equivalent to 128 W (electrical) illumination for a lamp having the same radiant output as the NEC lamp (20 YO emitted as radiation of wavelength < 400 nm). That is, the photoactivating power from 1 m2 of direct sunlight is approximately equivalent to 128 W of ' blacklight ' radiation.Assuming that light of a wavelength shorter than the bandgap of anatase (380 nm) photoactivates the catalyst, the result indicates that of the total 900 W m-2,19 ca. 2.8 YO has a wavelength shorter than 380 nm, in reasonable agreement with the data given by Ohta.25 Comparison with Salicylic Acid Also shown in fig. 7 is the destruction of 500 cm3 of 10pmol dm-3 salicylic acid circulated through the spiral at a flow rate of 100 cm3 min-' and illuminated with the 20 W lamp. The k' value was determined to be 0.117 f 0.001 min-', and 5.92 min. That is, under these conditions, salicylic acid was decomposed 1.72 times faster than methylene blue.This experiment also provides the nexus between methylene blue and the1300 Adsorption of Methylene Blue on Thin Films 100 90 80 - 70 5 60 g 50 40 30 20 10 0 E 3 i! P) P) I illumination time/min Fig. 8. Comparison of methylene blue disappearance and CO, appearance from 40 cm3 10 pmol dm-3 solution at pH 3. Flow rate 90 cm3 min-', 20 W lamp. other solutes in ref. (20). Measurements taken with methylene blue as the reference solute in other thin layer TiO, photoreactors could be used to estimate rates of disappearance of these other solutes. Carbon Dioxide Formation The formation of carbon dioxide was studied using solutions at pH 3. The results for 40 cm3 of 10 pmol dm-3 methylene blue circulated through the spiral at ca. 90 cm3 min-' and illuminated with the 20 W lamp are shown in fig.8. Also shown in fig. 8 is the disappearance of methylene blue from a 10 pmol dm-3 solution at pH 3 measured by the change in optical absorbance at 660nm. The solution volume was 500cm3 and the points shown in fig. 8 were the calculated times for the same decrease in concentration to occur in 40 cm3 of solution. It is seen that the rate of CO, formation lags considerably behind the rate of methylene blue disappearance. The to.5 value for CO, formation is ca. 4 min, four times slower than the rate of methylene blue disappearance (to.5 = 1 min). The importance of intermediate decomposition products as major precursors of CO, is thus clearly established. It is noted that total mineralization occurs at ca. 14 min. The observation of 100% conversion to CO, is consistent with reports of total mineralization for other organics photocatalytically oxidized by u.v.-illuminated TiO,.l, 3-9,12,14-17,26,27 Th e present results are consistent with the reaction stoichiometry of eqn (7). C,,H,,N,SCl+ 25.50, --+ 16C0, + 6H,O + 3HN0, + H,SO, + HC1.(7) It can be seen that in a closed system, with no air space, the total oxidation of a 10 pmolR. W. Matthews 1301 dmP3 methylene blue solution would exhaust the ambient oxygen concentration of an initially air-equilibrated solution (ca. 250 pmol dm-3).28 In the present experiments on the destruction of methylene blue, the solution was circulated through a reservoir open to the atmosphere and the oxygen depleted on passage through the reactor replenished with atmospheric oxygen. One of the referees has suggested that the increase in destruction rate may be caused by the possibly higher steady-state concentration of oxygen present in the solution at higher flow rates which would give rise to a higher concentration of adsorbed oxygen.The oxygen concentration is an important parameter in determining the rate of CO, formation in the photocatalytic oxidation of some solutes in free suspensions of Ti0,21(a) and would also be of importance in the present experiments but the extent of the inhibition by oxygen depletion is unknown. The transport of both oxygen and methylene blue to the photocatalyst surface are expected to be important factors in explaining the flow rate dependence in photocatalytic reactors of this type. Conclusions Aqueous solutions of methylene blue can be totally oxidized photocatalytically by contact with a thin layer of titanium dioxide illuminated with near-u.v.light. The rate of disappearance of methylene blue from the solution followed by spectrophotometric measurements at 660 nm obeys first-order kinetics. The decrease in apparent first-order rate constant, k’, with increasing solute concentration can be explained in terms of the integrated form of the Langmuir expression. One of the two parameters extracted from the analysis of the kinetic data agrees well with the adsorption affinity parameter obtained from equilibrium data on the non-illuminated system. The flow rate has a marked effect on k’; a near linear increase in k’ with flow rate occurs at low rates, but a limiting value is approached at high rates.Light from a ‘ blacklight’ fluorescent tube is a good photoactivating source for the titanium dioxide photocatalyst, but sunlight may also be used. The data obtained allow a quantitative comparison between the electrical power requirements using artificial light with direct sunlight. Measurements on the rate of disappearance of salicylic acid in the same photoreactor allow the methylene blue results to be related to other solutes. Hence methylene blue solutions could be used as a reference for other photoreactors of this type. The rate of CO, formation is ca. four times slower than the rate of methylene blue disappearance indicating the importance of intermediate decomposition products as precursors for CO, formation.Thanks are due to Cuno Pacific Pty. Ltd for financial assistance and to Mr Fred J. Fryer for technical assistance. 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