Role of Methanol Sacrificing Reagent in the Photocatalytic Evolution of

Use of D2O/CH3OH produced higher formation rates of HD and D2 than that of H2. The low H2 formation rates indicate that the direct reaction of CH3OH w...
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Role of Methanol Sacrificing Reagent in the Photocatalytic Evolution of Hydrogen Felipe Guzman, Steven S. C. Chuang,* and Cheng Yang† Department of Polymer Science, FirstEnergy Advanced Energy Research Center, The University of Akron, Akron, Ohio 44325-3909, United States ABSTRACT: The contribution of methanol (CH3OH) sacrificing reagent to the photocatalytic evolution of H2 from aqueous solutions has been studied by tracing the reaction of D2O over a Cu/S-TiO2 catalyst under UV illumination. Use of D2O/ CH3OH produced higher formation rates of HD and D2 than that of H2. The low H2 formation rates indicate that the direct reaction of CH3OH with photogenerated holes does not proceed to an appreciable extent in the presence of high concentrations of D2O. The role of CH3OH in accelerating hydrogen formation can be attributed to its ability to produce an electron donor, injecting its electrons to the conduction band.

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INTRODUCTION Production of H2 and O2 from H2O (i.e., water-splitting) by photocatalytic reactions on semiconductor materials such as TiO2 has been extensively studied.1−3 This reaction process holds the promise for generation of low cost hydrogen as energy carrier without CO2 emission. However, significant improvements in both catalyst activities and deactivation resistance are needed to achieve the sustainability potential.4 Thus, fundamental understanding of the photocatalytic watersplitting process is needed to guide further development of water-splitting catalysts. Photocatalytic reactions are initiated on TiO2 by absorption of light, which promotes electrons from the valence band to the conduction band and generates electron/hole pairs.5 Photogenerated holes (h+) can oxidize H2O molecules adsorbed on the catalyst surface, producing protons (H+) and O2. The H+ can further react with electrons (e−) to form H2, as shown in reactions 1−3. TiO2 + hν → e− + h+

(1)

H 2O(l) + 2h+ → 1/2O2(g) + 2H+, (E° = 1.23 V)

(2)

2H+ + 2e− → H 2(g)

(3)

The need for replenishing sacrificing reagents in order to maintain high rates of reaction precludes their potential use for continuous and large scale production of H2. Instead, sacrificing reagents can be used to assist in determining the photocatalyst efficiency of the hydrogen evolution reaction (reaction 3). A fundamental issue that needs to be addressed in the use of hydrogen-containing sacrificing reagents is whether these reagents simply serve as hole scavengers (i.e., species that rapidly react with holes depleting their concentration and thus preventing their recombination with photogenerated electrons) or contribute their hydrogen atoms to produce H2 gas. The objective of this study is to determine the contribution of CH3 OH sacrificing reagent on the overall rate of photocatalytic evolution of H2 from dilute aqueous solutions. The reaction is investigated using H2O/CH3OH and D2O/ CH3OH solutions over a Cu/S-TiO2 catalyst, assuming that H2O and D2O molecules exhibit similar photocatalytic reactivity. The amounts of H2, HD, and D2 produced were determined with a mass spectrometer (MS) and gas chromatograph (GC). The Cu/S-TiO2 catalyst was selected in consideration of the reported band gap narrowing due to sulfur doping of TiO213 and the enhanced absorption of visible light by Cu co-catalysts.14

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EXPERIMENTAL SECTION Preparation and Characterization of Cu/S-TiO2. The STiO2 powder was synthesized by sol−gel method according to previously reported procedures.15 Tetrabutyl titanate (13.6 g, 40 mmol, TCI America) was added dropwise to a mixture of thiourea (12.24 g, 160 mmol) and ethanol (200 mL, 3.42 mol) at 298 K under vigorous agitation, then acetic acid (14.4 g, 160 mmol, 96% Fisher Scientific) was added and allowed to react overnight. The resulted mixture was heated at 423 K for 2 h,

The evolution of stoichiometric amounts of H2 and O2 (i.e., water-splitting) is characterized by reaction rates that are 2 or 3 orders of magnitude smaller than conventional catalytic reactions.6 The low formation rates of H2 and O2 could be due to rapid recombination of holes and electrons and large driving force for the backward reaction producing H2O.7 Addition of electron donor molecules that react irreversibly with holes (i.e., sacrificing reagents) has been shown to be an effective approach for enhancing the rate of H2 formation. Common sacrificing reagents used in photocatalytic reactions include methanol (CH3OH), ethanol (CH3CH2OH), ethylenediaminetetraacetic acid (EDTA), cyanide (CN−), lactic acid (CH3CHOHCOOH), and formaldehyde (HCHO).8−12 Table 1 presents the rates of H2 formation over various photocatalysts in the presence of CH3OH sacrificing reagent, highlighting its prevalent use. © 2012 American Chemical Society

Special Issue: L. T. Fan Festschrift Received: Revised: Accepted: Published: 61

May 6, 2012 June 11, 2012 June 26, 2012 June 26, 2012 dx.doi.org/10.1021/ie301177s | Ind. Eng. Chem. Res. 2013, 52, 61−65

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Table 1. Rates of H2 Formation over Various Photocatalysts in the Presence of CH3OH photocatalyst

co-catalyst (wt %)

light sourcea

TiO2 TiO2 TiO2 TiO2/SiO2 TiO2 TiO2 TiO2 TiO2 TiO2 TiO2 N-TiO2 S-TiO2 TiO2 (NTN)c InVO4 S-TiO2 S-TiO2 S-TiO2 S-TiO2 S-TiO2

Pt(0.1) Pt (1.0) Pt (1.0) Pt (0.5) Cu (1.2) Ni (1.0) Ni (1.0), Ta (2.0) Ni (1.0), Nb (2.0) RuO2 (10), Pt (5) Pd (10) Pt (0.5) Pt (0.5) Ni (4.7) Ni(1.0) Cu (11.0) Cu (11.0) Cu (11.0) Cu (11.0) Cu (11.0)

Hg−P

light intensity (mW/cm2)

Hg−P Xe−P Xe−P Xe−P Xe−P

Xe−Q Xe−Q Vis Hg−Q Hg−Q Hg−Q Vis Hg−Q

CH3OH (vol %)

H2 formation rate (μmol/h g)

0 0 50 50 41 10 10 10 50 50

55 30 3000 86 2700 24.7 0.7 1.0 17334 7334 60 100 880 691 67.1 134 160 11.5 4.28

340 340 30 0 20 25 100 20 0b

30 30 30 30

ref 25 26 27 28 29 30 30 30 31 31 13 13 32 33 this this this this this

work work work work work

a

Hg−Q: combination of a 350−500 W Hg lamp with a quartz cell, Hg−P: combination of a 250−500 W Hg lamp with a Pyrex cell, Xe−Q: combination of a 250−500 W Xe lamp with a Quartz cell, Xe−P: combination of a 250−500 W Xe lamp with a Pyrex cell. Vis: visible light. b Reaction in H2O solution containing 0.4 M Na2S and Na2SO3. cNTN: Ni-intercalated titanate nanotube.

was quantified by a gas chromatograph (GC, SRI 8610C) equipped with a metal packed GC column (80/100 carboxen1004, Supelco) and a helium ionization detector (HID), selected because of its response to volatile inorganics such as CO, CO2, O2, N2, and H2 and the use of helium as carrier and makeup gas in place of the H2 and air streams commonly found with flame ionization detectors (FID). The molar composition of the products was determined by injecting 0.5 cm3 samples into a 40 cm3/min Ar stream entering a mass spectrometer (MS, Balzers QMG 112), and analyzing their mass/charge responses (i.e., m/e = 2, 3, and 4 from H2, HD, and D2, respectively).

dried under vacuum for 3 h, and calcined at 673 K for 3 h to produce the S-TiO2 powder. The atomic content of S atoms on the S-TiO2 powder was determined to be 3.6% by X-ray fluorescence (XRF μRayny 1300 series, Shimadzu). Cu was added to the S-TiO2 powder by the electroless plating approach, which involved the reduction of Cu2+ cations by formaldehyde (HCHO, Alfa Aesar). The S-TiO2 powder was dispersed in a plating solution bath containing copper sulfate (0.04 M CuSO4·5H2O), ethylenediaminetetraacetic acid disodium salt (0.08 M EDTA·Na), formic acid (0.08 M HCOOH), and 5 ppm pyridine as stabilizer. The temperature of the plating bath was kept constant at 343 K and the pH was adjusted to 12.5 under constant agitation. The Cu plated STiO2 catalyst (Cu/S-TiO2) was dried at 383 K for 24 h. The crystalline structure of the Cu/S-TiO2 catalyst was determined by X-ray diffraction (XRD) using a Philips APD 3700 diffractometer equipped with Cu Kα (wavelength 1.5406 Å). Diffuse reflectance ultraviolet−visible (UV−vis) spectra of the catalyst were recorded by a Hitachi U-3010 spectrophotometer equipped with a diffuse reflectance accessory. The spectra were collected at room temperature with BaSO4 as reference. Photocatalytic Studies. The photocatalytic evolution of H2 was carried out in a quartz square tubing reactor (100 cm3), equipped with gastight fittings for admission of gases and sampling of reaction products. During each experiment, a 30 cm3 solution containing 0.1 g of the Cu/S-TiO2 catalyst, CH3OH, and H2O or D2O were dispersed and exposed to visible light (450 W xenon arc lamp, Oriel 6265, 700−1000 nm wavelength range) and UV light (350 W mercury lamp, Oriel 6286). The lamps were positioned at a fixed distance of 5 cm from the reactor, resulting in light intensities in the 50−100 mW/cm2 range, as measured by a thermopile detector (818P001-12, Newport). The reaction was also carried out in the presence of sodium sulfite (Na2SO3, 0.4 M) and sodium sulfide (Na2S, 0.4 M) sacrificing reagents for comparison purposes. The total amount of products evolved (H2, HD, D2 and CO2)

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RESULTS AND DISCUSSION Catalyst Characterization. Figure 1 shows the XRD patterns of the S-TiO2 powder and the Cu/S-TiO2 catalyst before and after the photocatalytic studies. The patterns reveal that the crystalline structure of TiO2 corresponds to the anatase phase16,17 and its bulk structure was not altered during the

Figure 1. XRD patterns of S-TiO2 and Cu/S-TiO2 before and after the photocatalytic reaction. 62

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low potential of photogenerated electrons and (ii) lack of effectiveness of the sacrificing reagents for the hole scavenging reactions. Figure 4 shows the total amount of hydrogen (including H2, HD, and D2) produced as a function of time during UV

photocatalytic reaction. The absence of the rutile and brookite phases in the catalyst structure may be attributed to sulfur doping, which has been suggested to inhibit grain growth of the TiO2 crystallites and limit phase transitions during calcinations.16,17 Cu is present in the form of CuO18 with crystallite size of 15 nm calculated according to the Scherrer equation. The diffuse reflectance UV−vis spectra in Figure 2 reveal that addition of Cu extended the adsorption edge of the S-TiO2

Figure 4. Evolution of hydrogen (including H2, HD, and D2) during UV illumination of Cu/S-TiO2 dispersed in CH3OH, D2O, H2O/ CH3OH (9/1 molar ratio solution), and D2O/CH3OH (11/1 molar ratio solution).

Figure 2. Diffuse reflectance UV−vis spectra of TiO2 (P25, Degussa), S-TiO2, and Cu/S-TiO2.

illumination of the Cu/S-TiO2 catalyst dispersed in CH3OH, H2O/CH3OH (9/1 molar ratio solution), D2O/CH3OH (11/1 molar ratio solution), and D2O. UV illumination of CH3OH resulted in the evolution of significantly more hydrogen than that observed with the other solutions and with the Na-based sacrificing reagents. The evolution of hydrogen from primary alcohols with α-hydrogen such as CH3OH has been proposed to be initiated by the reaction with holes, producing protons (H+) and a hydroxyalkyl radical intermediate (i.e., ·CH2OH), as shown in reaction 4. The ·CH2OH radical intermediate possesses sufficiently negative oxidation potential (i.e, −0.74 V) and could further react to produce H+ and electrons, as shown in reaction 5. These electrons can be injected into the conduction band (doubling current effect).22

toward the visible region, indicating a good contact between the S-TiO2 and Cu grains.14,19 The band gap of the catalyst was estimated from the tangent lines in the plot of the square root of the Kubelka−Munk functions against the photon energy,13,20 as shown in the inset in Figure 2. The band gap energies of the S-TiO2 powder and Cu/S-TiO2 catalyst were 2.39 and 1.54 eV, respectively. The spectrum of TiO2 (p25, Degussa) was included for comparison purposes. Formation of Hydrogen by Photocatalytic Reactions over Cu/S-TiO2. Figure 3 shows the evolution of hydrogen

CH3OH + h+ → ·CH 2OH + H+

(4)

·CH 2OH → CH 2O + H+ + e−

(5)

+

The H produced from reactions 4 and 5 can also react with electrons to form H2 via reaction 3. Transient spectroscopic studies have demonstrated that the reaction of CH3OH and holes (reaction 4) occurs at a higher rate than that of holes and H2O molecules.23 Thus, decreasing the CH3OH concentration by addition of H2O should cause a decline in the rate of H2 evolution, which was observed in Figure 4. Since pure CH3OH is more reactive than pure H2O, the photocatalytic reaction of D2O/CH3OH solutions is expected to produce H2 at a higher rate than D2. MS analysis of the gas products from the reaction of the D2O/CH3OH (11/1 molar ratio solution) is shown in Figure 5. Figure 6 shows the product distribution during the first 10 h of reaction determined by GC and factoring of the relative MS responses presented in Figure 5. The amounts of D2 and HD increased with time while that of H2 leveled off after 7 h of reaction. The formation rate of H2 was 2.8 times lower than that of HD and 4.2 times lower than that of D2 during the first 5 h of illumination. After 10 h of reaction, the rate of H2 formation became 9.3 times lower than that of HD, and 13.2 times lower than that of D2. The low formation rate of

Figure 3. Evolution of hydrogen during UV and visible light illumination of Cu/S-TiO2 dispersed in a H2O solution containing Na2SO3 (0.4 M), and Na2S (0.4 M).

during UV and visible light illumination of the Cu/S-TiO2 catalyst dispersed in an aqueous solution containing Na2SO3 (0.4 M) and Na2S (0.4 M) sacrificing reagents. The formation of H2 increased with respect to time under UV illumination and remained relatively constant under visible light illumination. The rate of H2 formation using these Na sacrificing reagents under UV illumination was found to be in the same order of magnitude as those previously reported.21 The low rate of H2 formation during exposure to visible light could be due to (i) 63

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H+ + D+ + 2e− → HD(g)

(10)

The observation of HD in Figures 5 and 6 clearly confirms the occurrence of reactions 9 and 10, manifesting the function of CH3OH as a sacrificing reagent. However, the role of this sacrificing reagent is not through its direct reaction with h+ but with the reaction of ·OH. The other possible role of CH3OH for enhancing the rate of H2 formation includes inhibiting electron/hole recombination and the backward reaction to form H2O. This latter inhibition role would be consistent with the lack of O2 formation during the photocatalytic experiments. Determination of the inhibition of the electron/hole recombination reaction awaits further verification.

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Figure 5. MS responses obtained after injecting 0.5 cm3 samples of the gas phase species produced during UV illumination over Cu/S-TiO2 dispersed on D2O/CH3OH (11/1 molar ratio solution).

CONCLUSIONS CH3OH sacrificing reagent enhanced the rate of H2 formation from aqueous solutions under UV illumination and contributed less than its stoichiometric ratio of hydrogen to the overall rate of H2 formation. The low rate of H2 formation from D2O/ CH3OH solutions indicates that the reaction of hole with CH3OH does not proceed to an appreciable extent. CH3OH would react with hydroxyl free radical (·OH) produced from the reaction of h+ with H2O. The CH3OH reaction could produce an electron donor, ·CH2OH, injecting electrons to the conduction band, increasing the potential for the evolution of H2.

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AUTHOR INFORMATION

Corresponding Author

*Phone: 1-330-972-6993. Fax: 1-330-972-5856. E-mail: [email protected]. Figure 6. Amounts of H2, HD, D2, and CO2 produced during UV illumination of Cu/S-TiO2 dispersed on D2O/CH3OH (11/1 molar ratio solution). The evolution of gaseous products was determined via GC analysis, by factoring the relative MS responses of Figure 5.

Present Address †

Qingdao Institute of Bioenergy and Bioprocess Technology, Chinese Academy of Sciences, Qingdao, Shandong 266101, P.R. China. Notes

The authors declare no competing financial interest.

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H2 suggests that the reaction of hole with CH3OH (reaction 4) did not occur to an appreciable extent. This observation can be explained by the low concentration of adsorbed CH3OH on the TiO2 surface. A recent IR study has shown that the high content of H2O prevents alcohol from adsorbing as alkoxy (i.e., RCH2O) on the TiO2 surface and allows photogenerated hole to initiate the reaction with H2O as shown below:24 H 2O + h+ → ·OH + H+

(6)

OH− + h+ → · OH, (E° = 2.27V)

(7)

ACKNOWLEDGMENTS This work was partially supported by the FirstEnergy Corp. and the U.S. Department of Energy (Grant DE-FG26-01NT41294). (1) Kudo, A.; Miseki, Y. Heterogeneous photocatalyst materials for water splitting. Chem. Soc. Rev. 2009, 38, 253−278. (2) Esswein, A. J.; Nocera, D. G. Hydrogen Production by Molecular Photocatalysis. Chem. Rev. 2007, 107, 4022−4047. (3) Moon, S.-C.; Matsumura, Y.; Kitano, M.; Matsuoka, M.; Anpo, M. Hydrogen production using semiconducting oxide photocatalysts. Res. Chem. Intermed. 2003, 29, 233−256. (4) Fan, L. T.; Zhang, T.; Liu, J.; Schlup, J. R.; Seib, P. A.; Friedler, F.; Bertok, B. Assessment of Sustainability-Potential: Hierarchical Approach. Ind. Eng. Chem. Res. 2007, 46, 4506−4516. (5) Linsebigler, A. L.; Lu, G.; Yates, J. T., Jr. Photocatalysis on TiO2 Surfaces: Principles, Mechanisms, and Selected Results. Chem. Rev. 1995, 95, 735−758. (6) Fogler, H. S. Elements of Chemical Reaction Engineering, 3rd ed.; Prentice-Hall: Upper Saddle River, NJ, 1999. (7) Ni, M.; Leung, M. K. H.; Leung, D. Y. C.; Sumathy, K. A review and recent developments in photocatalytic water-splitting using TiO2 for hydrogen production. Renewable Sustainable Energy Rev. 2006, 11, 401−425. (8) Bamwenda, G. R.; Tsubota, S.; Nakamura, T.; Haruta, M. Photoassisted hydrogen production from a water-ethanol solution: a

Although the thermodynamic driving force for reaction 2 is higher than that of reaction 7, the former reaction can proceed slowly due to its involvement in a larger number of charge transfers. ·OH produced would react with CH3OH as shown in reaction 8. ·OH + CH3OH → ·CH 2OH + H 2O

(8)

Thus, it is expected that the presence of H2O or D2O would switch the CH3OH reaction pathway from reacting with h+ to reacting with ·OH or ·OD, producing the hydroxyalkyl radical intermediate (·CH2OH) of reaction 8. The ·CH2OH can proceed via reaction 5, providing electrons. These electrons then react with H+ and D+ to produce H2 and HD. D2 O + h+ → ·OD + D+

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