Photochemical production of hydrogen with zinc sulfide suspensions

J. Phys. Chem. 1984,88, 5903-5913

5903

Photochemical Production of Hydrogen with Zinc Sulfide Suspensions Jean-Franqois Reber* and Kurt Meier Ciba- Geigy, Central Research Laboratories, 4002 Basel. Switzerland (Received: May 30, 1984)

An efficient hydrogen production can be achieved by irradiating suspensions of ZnS in various electrolyte solutions (Sz-, SO:-, S20:-, H2P0,). The deposition of a metal such as platinum on the surface of ZnS microcrystals does not significantly improve the charge transfer reaction. In solutions containing SO?- ions, hydrogen generation occurs concomitantly with the oxidation of S032-ions to sulfate and dithionate. The rate of this reaction strongly depends on the quality of the ZnS sample. With the most active ZnS photocatalyst, a quantum yield of 0.90 has been obtained. Sixteen liters of hydrogen were produced with 1.0 g of ZnS in 35 h without any observable deactivation of the photocatalyst. The rate of hydrogen formation in solutions containing S2-ions is decreased due to the formation of disulfide ions which compete with the proton reduction. Addition of hypophosphite or sulfite ions, which efficiently suppress the disulfide formation, allows hydrogen to evolve at a higher rate. In mixtures of S2- and hypophosphite ions, phosphite and phosphate ions are the oxidation products. ions, formation of thiosulfate ions occurs concomitantly with hydrogen generation. In solutions containing both S2-and After the formation of thiosulfate ions, hydrogen generation continues with a smaller rate due to further oxidation of thiosulfate ions to tetrathionate. The disproportionation of tetrathionate leads to the formation of SO:- ions which can be further oxidized to sulfate. The overall process Sz- + SO;- + 5H20$ ! 5Hz + 2SOd2-is achievable with ZnS. This semiconductor can produce 5 times more hydrogen than cadmium sulfide which is not able to further oxidize thiosulfate. However, the rate of hydrogen formation is not very high in this process due to the reduction of thiosulfate ions which competes with the proton reduction. The properties of active ZnS photocatalysts,the reaction parameters, and the reaction products have been investigated in detail.

Introduction During the past few years, photoelectrochemical processes at semiconductor-electrolyte interfaces found a new interest because of their possible application in the conversion of solar energy into electrical or chemical energy. As only that light which is absorbed in a thin layer close to the semiconductor-electrolyte interface can initiate chemical reactions, large surface areas as those provided by powders are desirable. Various endothermic reactions with semiconductor powders were first achieved with TiOz suspension~.'-~ These early results have stimulated many studies of the photochemical production of hydrogen with suspensions of various semiconductors like TiOZ,k'3 SrTi03,5,14-16 CdS,'7-23 etc. Con-

(1) Freund, T.; Gomes, W. P. In "Catalysis Reviews''; Heinemann, H., Ed.; Marcel Dekker: New York, 1969; Vol. 3, pp 22-26. (2) Krasnovsky, A. A.; Brin, G. P.; Nikandrov, V. V. Dokl. Akad. Nauk SSSR 1975, 220, 1214; 1976, 229, 990. (3) Frank, S . N.; Bard, A. J. J. Phys. Chem. 1977,81, 1484-1488. (4) Bard, A. J. J. Photochem. 1979, 10, 59-75. (5) Lehn, J.-M.; Sauvage, J.-P.; Ziessel, R. N o w . J. Chim. 1981, 5, 291-295. (6) Pichat, P.; Herrmann, J.-M.; Disdier, J.; Courbon, H.; Mozzanega, M.-N. N o w . J. Chim. 1981,5,627-636. (7) Sakata, T.; Kawai, T. Chem. Phys. Lett. 1981, 80, 341-344. (8) Oosawa, Y. J. Chem. SOC.,Chem. Commun. 1982, 221-222. (9) Pichat, P.; Mozzanega, M.-N.; Disdier, J.; Herrmann, J.-M. N o w . J. Chim. 1982,6, 559-564. (10) Sakata, T.; Kawai, T.; Hashimoto, K. Chem. Phys. Lett. 1982, 88, 50-54. (1 1) Mills, A. J. Chem. SOC.,Chem. Commun. 1982, 367-368. (12) Oosawa, Y . Chem. Lett. 1983, 577-580. (13) Borgarello, E.; Pelizzetti, E. Chim. Ind. (Milan) 1983,65, 474-478. (14) Lehn, J.-M.; Sauvage, J.-P.; Ziessel, R. Nouu. J. Chim. 1980, 4, 623-627. (15) Lehn, J.-M.; Sauvage, J.-P.; Ziessel, R.; Hilaire, L. Isr. J . Chem. 1982, 22, 168-172. (16) Domen, K.; Naito, S.; Onishi, T.; Tamaru, K. Chem. Phys. Lett. 1982, 92, 433-434. (17) Darwent, J. R.; Porter, G. J. Chem. SOC.,Chem. Commun. 1981, 145-146. (18) Darwent, J. R. J. Chem. SOC.,Faraday Tram. 2 1981,77,1703-1709. (19) Harbour, J. R.; Wolkow, R.; Hair, M. L. J. Phys. Chem. 1981,85, 4026-4029. (20) Kalyanasundaram, K.; Borgarello, E.; Duonghong, D.; Gratzel, M. Angew. Chem. 1981, 93, 1012-1013.

firming the original discovery of Bard et al.,24,25all these investigations have shown that the combination of the semiconductors with electron transfer or hole transfer catalysts such as Pt, Pd, Rh, Ni, Ru02, WC, etc. is indispensable for an efficient charge transfer reaction and for the protection of the n-type chalcogenide-based semiconductors against photocorrosion. 17-20 As a matter of fact, it is well-known that n-type semiconductors such as cadmium sulfide26-28or zinc s ~ l f i d eare ~ ~not . ~stable ~ in aqueous solutions in which they undergo anodic photocorrosion leading to the formation of sulfur and/or sulfate ions. However, it has been reported that aqueous solutions of reducing agents acting as hole scavengers, among them S2-, S032-,or S2O3'-, stabilize cadmium sulfide e f f i ~ i e n t l y . ~ ' - ~ ~ Exploiting the stabilizing properties of these ions, we have achieved an efficient hydrogen production by irradiating suspensions of metalized CdS powders in solutions containing Szor S032-ions, as well as mixtures of Sz- and or S2- and hypophosphite ions.22~35*36 Because of the instability to light and the wide bandgap of zinc sulfide, photochemical reactions with this semiconductor have received little attention. Stephens et aL3' have shown that zinc sulfide is able to photochemically reduce oxygen to H202in the (21) Borgarello, E.; Kalyanasundaram, K.; Gratzel, M.; Pelizetti, E. Helv. Chim. Acta 1982, 65, 243-248. (22) Biihler, N.; Reber, J.-F.; Meier, K.; Rusek, M. European Patent Application, Publication No. 58136, Aug 1982. (23) Reber, J.-F.; Meier, K.; Buhler, N. "Book of Abstracts", 4th International Conference on Photochemical Conversion and Storage of Solar Energy, Jerusabm, Israel, Aug 1982; Weizmann Science Press: Jerusalem, Israel, 1982-1983; pp 252-254. (24) Kraeutler, B.; Bard, A. J. J . Am. Chem. SOC.1978, ZOO, 5985-5992. (25) Bard, A. J.; Dunn, W. W.; Kraeutler, B. US Patent 4264421, 1981. (26) Lal, P.; Ganguly, P. B. J. Indian Chem. SOC.1929, 6, 547-556. (27) Williams, R. J . Chem. Phys. 1960,32, 1505-1514. (28) Gerischer, M.; Meyer, E. Z . Phys. Chem. (Wiesbaden) 1971, 74, 302-318. (29) Platz, H.; Schenk, W. Angew. Chem. 1936, 49, 822-826. (30) Gloor, K. Helv. Chim. Acta 1937, 20, 853-877. (31) Inoue, T.; Watanabe, T.; Fujishima, A.; Honda, K.; Kohayakawa, K. J . Electrochem. SOC.1977, 124, 719-722. (32) Ellis, A. B.; Kaiser, S . W.; Wrighton, M. S. J . Am. Chem. SOC.1976, 98, 6855-6866. (33) Minoura, H.; Oki, T.; Tsuiki, M. Chem. Lett. 1976, 1279-1282. (34) Minoura, H.; Tsuiki, M. Electrochim. Acta 1978, 23, 1377-1382. (35) Reber, J.-F.; Biihler, N.; Meier, K.; Rusek, M. European Patent Application, Publication No. 100,299, Feb 1984. (36) Biihler, N.; Meier, K.; Reber, J.-F. J. Phys. Chem 1984, 88, 326 1-3268. (37) Stephens, R. E.; Ke, B.; Trivich, D. J. Phys. Chem. 1955, 59, 966-969.

0022-3654/84/2088-5903%01 .SO10 , 0 1984 American Chemical Society I

,

5904 The Journal of Physical Chemistry, Vola88, No. 24, 1984

Reber and Meier

TABLE I: Specific Surface Area, Crystallographic Structure, Crystallinity, and Impurity Content of the ZnS Samples Used in This Work specific surface crystalsample material area,’ m2/g structure linity impurities, ppm I Koch-Light, pure grade 6.1 f 0.1 hexagonal wurzite (a) medium 4.5% BaSOa, 2300 Al, 1700 CI. (Lot 78441‘) 530 Sr, 290 Na, 213 Mg, 120 Ca, 75 Fe I1 Alfa Ventron, el. grade 30.2 f 0.6 cubic sphalerite (@) low 275 Ca, 16 Mg, 15 Cd, 15 Na, 13 Fe (Lot 080581) not visible (cubic?); very low 32100 Na, 1300 Si, 190 Ag, 82 Ba, 29 Ca, 27 Mg I11 ZnS precipitated in an 119 f 2 excess of S2- ions presence of some ZnO IV ZnS precipitated in an 162 3 very low 4100 Na, 1100 Si, 180 Al, 27 Ca, 17 Ni cubic sphalerite (b) excess of Zn2+ions V Alfa Ventron ZnS thermally 1.09 0.05 cubic sphalerite (p) medium 180 Ag, 175 Al, 50 Na, 19 Mg, 10 Cd treated at 850 OC during to high 60 rnin in a stream of H2S VI ZnS of sample I11 thermally 2.33 f 0.05 hexagonal wurzite (a)+ medium 34700 Na, 1200 Si, 125 Ba, 52 Ca, 31 Mg cubic sphalerite (8) to high treated as sample V

*

*

Reference 40. presence of phenol. The photochemical production of hydrogen with suspensions of zinc sulfide has been first achieved by Yanagida et al.38in the presence of various organic sacrificial electron donors, such as methanol, ethanol, tetrahydrofuran, etc. These authors have reported that the in situ precipitation of ZnS leads to a satisfactory efficiency without the use of any redox catalyst. Kisch et al.39have reported that zinc sulfide resulting from the decomposition of a Zn dithiolene is more active than colloidal ZnS for the generation of hydrogen in the presence of sacrificial electron donors like 2,5-dihydrofuran. We have found35that hydrogen could be generated very efficiently by irradiation of suspensions of metalized ZnS powders in the presence of hole scavengers such as Sz- or SO3” ions. In this paper, the production of hydrogen with powders of zinc sulfide free from any charge transfer catalyst and dispersed in solutions containing S2-or SO$ ions, as well as mixtures of Sz- and SO?or S2-and hypophosphite ions, is reported. The properties of active ZnS photocatalysts, the reaction parameters, and the formation of the reaction products have been investigated in detail. Experimental Section Materials. The zinc sulfide powder mainly used in this work (sample I) was supplied by Koch-Light Laboratories Ltd. (pure grade, Lot 78441). Its specific surface area@was 6.1 f 0.1 m2/g. X-ray diffraction measurements have shown that the microcrystals have the hexagonal wurzite structure (a-ZnS) and possess a medium crystallinity. A small concentration of BaS04, revealed by X-ray diffraction measurements, was confirmed by analysis showing the presence of 4.5 wt% BaSO.,. The elemental analysis also revealed the presence of 2300 ppm Al, 1700 ppm C1,530 ppm Sr, 290 ppm Na, 213 ppm Mg, 120 ppm Ca, and 75 ppm Fe. The photocatalytic activity of this ZnS powder was compared with that of other ZnS samples having different physicochemical properties. The specific surface, crystallographic structure, crystallinity, and impurity content of the ZnS samples used in this work are listed in Table I. Both ZnS samples I11 and IV were precipitated from water solutions of zinc acetate and sodium sulfide according to the “single-jet” method, well-known in photographic science. This simple method allows to achieve nucleation and growth of all the microcrystals, either in a sulfide or zinc excess. Sample I11 was obtained by adding dropwise 500 mL of a 1.O M solution of Zn(CH3C00)2.2Hz0 (Merck, analytical grade) to 500 mL of a well-stirred 1.2 M solution of NazS-8.5H20 (Merck, analytical grade) over a period of 15 min. After another 20 min of stirring a t room temperature, the zinc sulfide was collected by filtration. The white precipitate was redispersed in 500 mL of water, stirred for 30 min, filtrated again, and dried at 100 OC (12000 Pa). Sample IV was precipitated according to the same scheme, except that 500 mL of a 1.0 M Na2S.8.5H20 (38) Yanagida, S.;Azuma, T.; Sakurai, H.Chem. Letr. 1982, 1069-1070. (39) Bticheler, J.; Zeug, N.; Kisch, H. Angew. Chem. 1982,94,792-793. (40) The specific surface area of the ZnS samples was determined by the volumetric BET method using nitrogen as adsorbent.

solution was added to an equal volume of a 1.2 M Zn(CH3CO0)z.2H20 solution. For the platinum reduction on the ZnS particles, hexachloroplatinic acid, H2PtC16.xH20, supplied by Aldrich, was used. The photolytical decomposition of zinc sulfide was determined with l,l’-dimethyl-4,4’-bipyridiniumchloride purum supplied by Fluka. Other products used in this study were all analytical grade reagents. Platinum Deposition on the ZnS Microcrystals. Although the ZnS powders were used without further treatment in most of the experiments, some samples have been surface platinized for comparison according to the following methods: Method A . ZnS microcrystals were platinized by photoreduction of R(1V) ions with UV light in the presence of acetic acid. A 10.0-g sample of ZnS powder was added to 90 mL of a 0.25 M solution of acetic acid, buffered at pH 4.5 and containing various amounts of H2PtC16.xHz0. After 60 s of ultrasonic stirring, the obtained suspension was introduced into a 150-mL quartz photoreactor equipped with a water-cooled mercury immersion lamp (Philips HPK 125). At a constant temperature of 60 f 0.5 “C,the vigorously stirred suspension was deaerated with argon for 30 rnin and subsequently illuminated for 30 min. The photocatalyst was then collected by filtration, washed, and dried a t 60 OC (12000 Pa). Method B. A 10.0-g sample of ZnS powder was stirred at room temperature for 10 rnin in 50 mL of water containing 0.4 g of H2PtCl6.xH20. After evaporation of the solvent under reduced pressure, the powder was dried at 60 “ C (12000 Pa). The platinum reduction was carried out in a hydrogen stream for 2 h at 200 OC. Method C. A 10.0-g sample of platinized ZnS powder obtained from method B was suspended in 90 mL of a 0.25 M solution of acetic acid buffered at pH 4.5, deaerated, illuminated, and isolated as described in method A. Determination of the Photoactivity. The photoactivity was determined by measuring the volume of hydrogen produced during the irradiation of suspensions of zinc sulfide in electrolyte solutions. Two different photoreactors were used. (1) Standard experiments were carried out in the quartz photoreactor described in the platinization method A. In each run, 400 mg of zinc sulfide powder, dispersed in 90 mL of electrolyte solution at 60 0.5 OC, was exposed to UV light. Hydrogen evolution41was measured with a 50-mL gas buret. Sulfate ions in the presence of S032-and S2062-ions were analyzed by precipitation with BaClZsolution. The excess of BaClZwas then titrated with EDTA solution.42 The concentration of dithionate ions was determined by oxidation with a hot solution of potassium dichromate after separating the SO?- ions by flushing N2through the acidified solution? Analysis of thiosulfate ions in the presence

*

(41) Samples of gas evolved under illumination have been analyzed by gas chromatography using a thermal conductivity detector, a molecular sieve column, and N2 as carrier gas. (42) Soffer, N. Analyst (London) 1961,86, 843. (43) Friessner, A. Z . Elektrochem. Angew. Phys. Chem. 1904, 10, 265-289.

Photochemical Production of Hydrogen with ZnS Suspensions The Journal of Physical Chemistry, Vol. 88, No. 24, 1984 5905 Results 1 . Influence of the Anodic Reaction on the Photochemical Production of Hydrogen. Figure 1 shows the amount of hydrogen produced by illuminating unplatinized zinc sulfide suspended in solutions containing sulfide or sulfite ions, as well as mixtures of S2-with sulfite or hypophosphite ions. Curve 1 (Figure 1) represents the oxidation of sulfite with the concomitant formation of hydrogen according to H2 2 0 H cathode: 2 H 2 0 2e-

/3

-

+

+

The oxidation of S032-ions follows two different routes, forming sulfate and dithionate ions according to anode:

-

Irradiation time [hours] Figure 1. Amount of hydrogen evolved (mL) vs. illumination time (h) for sample I (400 mg of unplatinized ZnS) dispersed in 90 mL of a solution containing the following: curve 1, 0.35 M N a 2 S 0 3 (pH 9.5); curve 2,0.24 M Na2S (pH 13.2); curve 3,0.24 M Na2S, 0.35 M H3P02, and 1.10 M NaOH (pH 13.8); curve 4, 0.24 M Na2S and 0.35 M N a 2 S 0 3 (pH 13.2). Reaction temperature: 60 O C .

of S2-and S032-ions was carried out iodometrically after separating S2-ions by precipitation with zinc acetate in an alkaline medium and masking S032-with formaldehyde. (2) The quantum yield was determined by illuminating a suspension of 1.O g of zinc sulfide in 90 mL of electrolyte solution thermostatized at 60 f 0.5 "C. The suspension was irradiated in a quartz reactor with light emitted by a high-pressure mercury lamp (Spindler & Hoyer HBO 200 W/2) and selected by a 313-nm (f7 nm) Oriel interference filter (No. 5651). The hydrogen evolution was determined with the detector described by Griiniger et a1.& The electrode had been calibrated with known concentrations of hydrogen. The incident flux of photons used for the quantum yield determination was 9.4 X einstein/h at 313 nm, as measured with a ferrioxalate solution in the same device. Before exposure to light the dispersions were purged with nitrogen for 30 min to prevent oxygen reduction. Photolytical Reduction of Zinc Sulfide. The degree of photolytical reduction was determined according to the method described by Gutitrrez and H e r ~ g l e i n . ~The ~ irradiated ZnS suspension was reacted with methylviologen (1,lt-dirnethyl-4,4'-bipyridinium chloride), and the resulting concentration of the stable radical cation of methylviologen was measured spectrophotometrically. A 20-mg sample of zinc sulfide thoroughly dispersed by ultrasonic stirring in 3.0 mL of a 0.35 M Na2S03solution was placed in a photometer cell equipped with two stop clocks and a small reagent reservoir containing 0.1 mL of a 0.05 M solution of methylviologen and 0.1 mL of a 0.07 M solution of zinc acetate to prevent the adsorption of the methylviologen radical cation (MV+) on the ZnS surface. The whole device was deaerated with purified argon for 40 min, while the temperature of the magnetically stirred suspension was maintained at 55 0.5 OC. After irradiation with a high-pressure Hg lamp (Spindler & Hoyer HBO 200 W/2), the ZnS suspension was mixed with the reagent solution. The absorbance of the blue solution was determined at 600 5 nm (e = 1.1 X L/(mol/~rn)).~~,~~

*

*

(44) Griiniger, H. R.; Sulzberger, B.; Calzaferri, G. Helv. Chim. Acta

2h+ t

1

SO,"

+ 20"

2s032-

-+

+ H,O

-+SO,'-

S,O,~'

where h+ denotes a positive hole produced in the valence band of ZnS. After 290 mL of H2(1 1.8 mmol) was produced, 8.2 mmol of S042-and 4.1 mmol of S202-were found in the solution. This means that 50% of S 0 3 2 -ions were oxidized to sulfate and 50% to dithionate. The photochemical formation of a mixture of sulfate and dithionate ions observed with zinc sulfide is similar to the results obtained by irradiating suspensions of p l a t i n i ~ e dor ~~ colloidal45cadmium sulfide in the presence of sulfite ions. The volume of hydrogen evolved is a linear function of the illumination time as long as the SO?- concentration is sufficiently large. The slope indicates that the rate of hydrogen formation is equal to 39 mL of H2/h for 400 mg of photocatalyst. No fatigue becomes apparent after several hours. Curve 2 (Figure 1) shows the hydrogen formation during the simultaneous oxidation of sulfide to disulfide ions according to anode: 2S2- + 2h+ S2-

-

In an alkaline medium, only disulfide ions are formed and no sulfur precipitation is observed. At the beginning of the reaction, the rate of hydrogen formation-higher in the presence of S2-ions than in a solution containing S 0 3 2 -ions-is equal to 166 mL of H2/h.48 However, it decreases during the course of the reaction, due to the increasing concentration of disulfide ions which compete with the reduction of protons according to 2e-

--

+ 2 H 2 0 H2 + 2 0 H 2e- + S22- 2S2-

Furthermore, the yellow disulfides act as an optical filter, reducing the light absorption of ZnS. Nevertheless, the decrease of efficiency is much less pronounced for ZnS than for platinized cadmium sulfide.36 Moreover, the addition of 2.2 mmol of sulfur to the 22 mmol of sulfide present at the beginning of the reaction only decreases the rate of hydrogen formation from 166 to 124 mL of H2/h. The formation of disulfide ions which competes with the reduction of protons can be prevented by adding a reducing agent to the solution. Curves 3 and 4 (Figure 1) describe the hydrogen production obtained by adding hypophosphite or sulfite ions to the S2-solution. Curve 3 depicts the hydrogen production observed in a solution containing Sz-and hypophosphite ions simultaneously. In this system, previously described for CdS,36an accumulation of disulfide ions is prevented by the reducing action of the H2P02-ions according to anode:

2S2-

- +

+ 2h+

Sz2-+ H2P02-+ 3 0 H -

S22-

2S2-

HP03- + 2H20

In such a medium, the hypophosphite ions are irreversibly oxidized and the Sz- ions are recycled. The rate of hydrogen formation, fast at the beginning of the reaction (295 mL of H2/h), slowly decreases due to the decreasing concentration of H2P02- ions.

1978, 61, 2375-2380.

(45) GutiCrrez, M . ; Henglein, A. Ber. Bunsenges. Phys. Chem. 1983,87, 474-41 8. (46) Yuen, S . H.; Bagness, J. E.; Myles, D. Analyst (London) 1967, 92, 375-381.

(47) Trudinger, P. A. Anal. Biochem. 1970, 36, 222-225. (48) As the rate of hydrogen formation in NazS and NazS/Na2S03solutions decreases with irradiation time, the values given are measured between 5- and 15-min illumination.

5906 The Journal of Physical Chemistry, Vol. 88, No. 24, 1984

Reber and Meier

After all these ions have been consumed, the hydrogen generation suddenly drops and continues with a rate equal to 91 mL of H2/h due to the further oxidation of phosphite ions into phosphate according to Sz2- + HP03- + 3 0 H -

+

2Sz-

+ P043- + 2 H 2 0

At this point, the quantity of hydrogen produced, equal to about 840 mL of H2(34.3 mmol), is larger than the expected value which would result from the oxidation of 31.5 mmol of hypophosphite. This indicates that the further oxidation of phosphite ions into phosphate occurs before the complete consumption of hypophosphite ions. The evolution of hydrogen from a solution containing phosphite (0.35 M) and sulfide (0.24 M) ions is slower (1 12 mL of H2/h), and the accumulation of disulfide ions is not completely prevented. In the absence of Ss ions (alkaline medium) the rates of hydrogen formation obtained in solutions containing hypphosphite or phosphite ions (0.35 M) are much lower, 53 and 20 mL of H2/h, respectively. I”? OO 005 010 015 Hydrogen generation with the concomitant oxidation of hypophosphite ions can be catalyzed by Raney Ni in the absence Concentratidn of s20$ ions [ m o k / l ] of However, in contrast to the photochemical reaction, Figure 2. Rate of hydrogen formation (mL/h) at a given time as a Raney Ni is not able to further oxidize phosphite to phosphate function of the concentration of the photochemically formed thiosulfate ions. 60 “C. Sample I (400 mg of unplaAs previously observed for CdS by several a ~ t h o r s , ~ ions ~ ~(M). ~ Reaction ~ ~ ~ temperature: ~ ~ ~ tinized ZnS) is dispersed in 90 mL of solution containing 0.24 M Na2S sulfide oxidation in the presence of S032-ions (curve 4, Figure and 0.35 M Na2S03. 1) yields thiosulfate ions also in the case of ZnS, according to

-

anode:

2S2-

+ 2h+

SZ2-+ S032

+

-+

S22-

Sz032-+ S2-

The formation of disulfide ions is strongly suppressed by the addition of S03z-ions. After the production of 370 mL (15.1 mmol) of hydrogen, 14.7 mmol of Sz032-was detected in the solution. The oxidation product thiosulfate absorbs less UV light than the yellow disulfides and the sulfide ions. The rate of hydrogen formation, equal to 185 mL of H2/h48at the beginning of the reaction, decreases during the course of the reaction. This significant decrease of efficiency, not observed for cadmium s ~ l f i d e , ’cannot ~ be explained by optical reasons or by the concentration decrease of the sacrificial donors, S2- and SO?- ions. It is due to the increasing concentration of thiosulfate ions which are reduced in competition with protons according to

+ 2H20 2e- + S20322e-

-

-+

H2 + 2 0 H S2- + S032-

Since hydrogen and thiosulfate are formed in an approximately stoichiometric ratio in the concentration range studied, the direct oxidation of sulfite to sulfate ions and the anodic oxidation of thiosulfate to tetrathionate ions (see next section) can occur only to a small extent. Consequently, when both the concentrations of S2-and SO3” ions are not too small, the major valence band process is the thiosulfate formation, and the three other possible processes, disulfide, sulfate, and tetrathionate formation, are nonexistent or very low. Figure 2 shows that the rate of hydrogen formation at a given time decreases linearly with the increasing concentration of thiosulfate. As the proton and the thiosulfate reduction are the two competitive conduction-band processes, the rate of thiosulfate reduction increases with the concentration of S203’- ions. 2. Hydrogen Evolution with ZnS Suspensions in Thiosulfate. Irradiation of a suspension of zinc sulfide in a mixture of S2-and (49) Schroter, R. “Newer Methods of Preparative Organic Chemistry”; Interscience: New York, 1948; p 70. (50) See ref 23, Part B, p 32. (51) Thewissen, D. H. M. W.; Tinnemans, A. H. A.; Eeuwhorst-Reinten, M.; Timmer, K.; Mackor, A. Nouu. J . Chim.1983, 7, 191-194. (52) Borgarello, E.; Erbs, W.; Gratzel, M. Noun J . Chim. 1983, 7, 195-198. (53) Thewissen, D. H. M. W.; Tinnemans, A. H. A.; Eeuwhorst-Reinten, M.; Timmer, K.; Mackor, A. Sol. Energy R&D Eur. Community, Ser. D 1983, 2, 58-63.

-

Irradiation time [hours] Figure 3. Main reaction products (mmol of H2,SO3-*,and S2-) resulting from the irradiation of a zinc sulfide suspension in Na2S203vs. the irradiation time (h). Sample I (400 mg of unplatinized ZnS) is dispersed in 90 mL of a solution containing 0.1 M Na2S203and 1.0 M NaOH.

Reaction temperature:

6 0 OC.

S032-ions yields thiosulfate ions. It has been observed in the previous section that the increasing concentration of S2032- ions induces a decrease of the rate of hydrogen formation attributed to their reduction by conduction-band electrons in competition with protons. The aim of this section is to study the redox reactions occurring during the irradiation of a zinc sulfide suspension in thiosulfate. A 400-mg sample of unplatinized zinc sulfide (sample I), dispersed in 90 mL of a solution containing 0.1 M Na2S203and 1.0 M NaOH, was irradiated during various periods of time. The hydrogen evolution was measured during the whole exposure to light. After filtration of the powder, the different solutions were analyzed as described in the Experimental Section. Illumination of alkaline ZnS suspensions in the presence of thiosulfate ions leads to the consumption of S2O3’- and to the and Sz062-ions. formation of hydrogen, S2-, S032-,Sod2-, In Figure 3, the amount of the products (H2, SO-,; and S2-) formed during the photoreaction is expressed as a function of the irradiation time. Valence-band holes produced by irradiating ZnS suspensions in thiosulfate are able to oxidize S 2 0 3 2 - ions to tetrathionate according to

+

2S2032- 2h+

+

S4062-

Photochemical Production of Hydrogen with ZnS Suspensions The Journal of Physical Chemistry, Vol. 88, No. 24, 1984 5907

(hy

, ,/'

- I I

- I

B

3

*

31

I

I

I

,

,,

2

H2+2A H2O

/

\@

s20:-

?4 H,O

)c, S,O:-+

I

,

73

s2-+so:-

SO',-+ 3 H'

0

SO:-+ H20

2 OH-+SO',(2 so:-,

[ 52026- 1

Figure 5. Schematic illustration of the reactions resulting from the irradiation of zinc sulfide in a thiosulfate solution.

Figure 4. Hydrogen production in the thiosulfate system (corresponding to the experimental data of Figure 3).

In an alkaline medium, tetrathionate ions disproportionate quantitatively into sulfite and thiosulfate according to the net reaction being

+

0.5S2032- 1.5H20

+ 2h+

-

S032-+ 3H+

If the proton reduction would be the only conduction-band process, the amount of hydrogen evolved would be stoichiometrically equal to the amount of sulfite formed in the valence-band process. The excess of sulfite ions formed during the first 12 h and also the presence of sulfide ions indicate that part of the thiosulfate ions are reduced by conduction-band electrons to S2- and SO3" ions, competitively with protons to hydrogen. The rate of thiosulfate reduction cannot be determined while the resulting Sz- and S O P ions react reversibly with photoholes to produce thiosulfate. The concentration of sulfide ions approaches a plateau, indicating that the two reactions-formation and reduction of thiosulfate-tend to equilibrate themselves. Consequently, the cyclic formation and reduction of thiosulfate ions consume a lot of charges which are lost for the most valuable process: hydrogen production. As the number of electrons reacting with adsorbed electron acceptors is equal to the number of holes reacting with electron donors, one can write nlH2

+ x S O ~ (cathode) ~=

(anode)

+ 2S2032-

and xSO~~ (cathode) = xS2- (cathode)

2S2032-= - z S O ~ ~(consumed) = - z S 2 (consumed) The amounts of sulfite and sulfide detected in the solution are n2S032- = x S O ~ ~(cathode) -

+

(anode) - z S O ~ (consumed) ~-

n3S2- = xS2- (cathode) - zS2- (consumed) where n, denotes determined amounts of substances and x , y , and z are not measurable parameters. After resolution of these equations, one finds that the amount of hydrogen evolved should be stoichiometrically equal to nlH2 = n2S032- - 2n3S2The expected amount of hydrogen, calculated according to this equation from the amounts of S032-and S2- determined analytically, is represented by the dashed line in Figure 4. The accordance between measured and calculated amounts of hydrogen is good at low conversion. At higher conversion, the amount of hydrogen is larger than expected, as also shown in Figure 3. The

-

Irradiation time [hours] Figure 6. Amount of hydrogen produced (mL) as a function of the irradiation time (h). Sample IV (400 mg of unplatinized ZnS) is dispersed in 90 mL of a 1.0 M NaOH solution containing 9.0 mmol of N a 2 S 0 3and 7.2 mmol of Na2S. Reaction temperature: 60 OC.

above description does not take into account the direct oxidation of sulfite to sulfate and/or dithionate which would produce an additional amount of hydrogen. The excess of hydrogen produced after 22 h of irradiation, equal to 0.65 mmol, should correspond to 0.32 mmol of S042-. The amount of sulfate detected at this point, in fact equal to 0.38 mmol, indicates that this process occurs at a significant rate. All reactions which occur at the interface zinc sulfide-S2-/ S032-/S2032system are summarized in Figure 5 . The previous results have shown that thiosulfate ions are photochemically formed from a mixture of S2-and Sol-ions with the concomitant generation of 1 equiv of hydrogen. The oxidation of thiosulfate to sulfite and further to sulfate would produce an additional 4 equiv of hydrogen. Thus, the overall process

s2- so32-

5H20

5H2

2s042-

would produce 5 equiv of hydrogen. In order to check the feasibility of such a process, 400 mg of ZnS of the most active sample IV (see section 6), dispersed in 90 mL of solution of 1.0 M N a O H containing 9.0 mmol of N a 2 S 0 3and 7.2 mmol of Na2S, was irradiated during a longer period of time. The amount of hydrogen produced during the irradiation is shown in Figure 6. The rate of hydrogen formation, about 230 mL of H2/h at the beginning of the reaction, slowly decreases and then becomes constant (73 mL of H,/h) for several hours. After 9.8 h of irradiation, 772 mL of hydrogen (31.5 mmol) has been produced, which is approximately 4.4times the amount

5908

The Journal of Physical Chemistry, Vol. 88. No. 24, 1984

Reber and Meier

t

140

t

/

120.

F \ d

E

I

a,

+

c .-0

c d

E

u0

s 0 01 M

01

003

01

03

10

I

-

ELectrolyte concentration [mole/ll Figure 7. Rate of hydrogen formation (mL/h) at 60 OC as a function of the Na2S03or NazS concentration (M) expressed in a logarithmic scale. Sample I (400 mg of unplatinized ZnS) is suspended in 90 mL

of solution.

of sulfide ions present at the start (7.2 mmol). At this moment, 0.007 m o l of S2-, 0.59 mmol of SO?-,0.22 mmol of S20?-, 5.10 mmol of S042-, and 4.95 mmol of SzOs2-were detected in the solution. The amount of hydrogen calculated from these values is 31.3 f 0.1 mmol. 3. Influence of the Electrolyte Concentration. The influence of the electrolyte concentration on the rate of hydrogen formation and S2-ions. was measured for the oxidation reaction of S032As shown in Figure 7, the rate of hydrogen production in both media increases with increasing electrolyte concentration to a constant value. Electrolyte concentrations higher than 1.O M do not further improve the yield of hydrogen formation. All the electrolyte solutions used here absorb part of the light emitted by a mercury lamp. Since the absorption of light by an electrolyte is concentration dependent, the amount of photons absorbed by the ZnS particles decreases with increasing electrolyte concentration." In order to exclude this undesirable optical effect, irradiation of a ZnS suspension in Na2S03 has been carried out by using a Pyrex filter (A > 300 nm). Although sodium sulfite does not absorb light above 300 nm, an identical concentration dependence was observed. Consequently, the saturation observed for the highest electrolyte concentrations does not result from an optical effect. Therefore, the shape of the curves is related to the electron-hole recombination process. Concentrations of hole scavengers lower than 1.0 M are not sufficient to compete efficiently with the electron-hole recombination. A similar loss of activity at high Na2S03concentrations has been observed by Aruga et al.55 in the case of suspensions of cadmium sulfide. In the presence of sulfide ions, hydrogen is not evolved at a constant rate because the disulfide ions resulting from the oxidation of Sz- ions by photoholes compete efficiently with the reduction of protons. Therefore, the rate of hydrogen formation depends not only on the sulfide concentration but also on the concentration of disulfide ions. The evolution of hydrogen resulting from the irradiation of ZnS suspensions in NazS solutions of various concentrations is represented in Figure 8. For each point of the curves which corresponds to a specific period of irradiation, the rate of hydrogen formation, the corresponding concentration of disulfide ions (as So species), and the residual concentration of S2-ions [S2-]

Irradiation time [min.]

-

Figure 8. Influence of the sulfide concentration on the hydrogen production as a function of the irradiation time. Sample I (400 mg of unplatinized ZnS) is dispersed in a 90-mL solution of Na2S. Reaction temperature: 60 "C. I""

-

-

10 M Na& 024M 010 M 003M

E E

Y

c W

Figure 9. Rate of hydrogen formation (mL/h) at a given time in sulfide solutions as a function of the composition of the solution. The ratio ([S2-] - [So])/[So], determined from the curves of Figure 8, is expressed in a logarithmic scale.

- [So] can be determined. Figure 9 shows that the rate of hydrogen formation at a given time follows in good approximation a logarithmic function with the ratio of the residual concentration of S2-ions to the So concentration ([$-I - [ S o ] ) / [ S o ]as , long as the concentration of So species is not to low.56 Hydrogen is generated with a decreasing rate which depends on the respective concentrations of sulfide and disulfide ions according to the following experimental equation r ( H 2 )(mL/h) = K(0.17

+ log ([Sz-]/[So]

- 1))

where K = 108 mL/h. The constant K depends on the experimental conditions (light intensity, temperature, etc.). 4. Influence of the Solution p H . The influence of the pH of the electrolyte solution on the rate of hydrogen formation was

(54) The absorbance of 1.0 M solutions of Na2Sand Na2S03is larger than 1.0 cm-' at wavelengths shorter than 284 and 264 nm, respectively.

( 5 5 ) Aruga, T.; Domen, K.; Naito, S.;Onishi, T.; Tamaru, K.Chem. Lett. 1983, 1037-1040.

(56) The concentration of disulfide ions generated after 10-min irradiation is sufficient.

Photochemical Production of Hydrogen with ZnS Suspensions

PH

The Journal of Physical Chemistry, Vol. 88, No. 24, 1984 5909

-

Figure 10. Rate of hydrogen formation (individual measurements) and

concentration of unprotonated sulfite ions (curve) vs. pH of the solution. Reaction temperature: 60 O C . Sample I (1.0 g of unplatinized ZnS) is dispersed in 450 mL of a 0.35 M solution of Na2S03. O ’ 30

io

50

60

-

70

Temperature [“I Figure 12. Temperature dependence of the rate of hydrogen formation in 0.35 M Na2S0, and 0.24M Na2S. Sample I (400 mg of unplatinized ZnS) is dispersed in 90 mL of solution.

150-

t E

-1:

is 6.91 at 18 the actual concentration of ions is the determining factor and not the pH. Glasstone and Hickling5* have shown that the electrolytic oxidation of potassium sulfite at a Pt anode produces a mixture of sulfate and dithionate ions and that the formation of dithionate ions is highest a t a p H of about 8. In the case described here, the shoulder observed between pH 7.5 and 9.0 may be due to a change of the ratio between the two oxidation reactions. However, this question was not further examined in this work. In Figure 11, the rate of hydrogen formation of ZnS dispersed in solutions of sulfide ions and the actual concentrations of hy-

drosulfide and unprotonated sulfide ions are plotted as a function of the pH of the solution. In this medium, the rate of hydrogen formation is completely independent of the pH and of the degree of dissociation of the sulfide ions. This surprising result is quite different from the one obtained with suspensions of CdS activated with Pt or RuO? In the latter case, the hydrogen production was found to depend on the pH of the solution2’ or on the actual concentration of unprotonated sulfide i ~ n s . ’ ~ * ~ l 5. Temperature Dependence of Hydrogen Generation. The temperature dependence of the rate of hydrogen formation is shown in Figure 12 for unplatinized zinc sulfide (sample I) suspended in 0.35 M Na2S03 and in 0.24M Na2S, respectively. In the range studied, the efficiency of the hydrogen production increases linearly with the temperature and the temperature coefficient is almost identical for the two reactions (0.71 mL of H2/(h deg) in 0.35 M Na2S03and 0.79mL of H2/(h deg) in 0.24 M Na2S). The dispersion of the results obtained with a sample of platinized ZnS (method A, sample I, 1.7% Pt) is about 5-10 times smaller. Moreover, the temperature coefficient of such a sample suspended in 0.35 M Na2S03is also identical (0.76mL of H2/(h deg)). This fact suggests that the efficiency of the process is controlled by the anodic reaction. Changing the temperature leads to shifts in the energy levels of both the semiconductor and electrolyte and causes changes in the electrochemical reaction rates. It is known59that only small shifts of the flat-band and redox potentials of the electrolyte species are observed with increasing temperature. Therefore, the enhancement of photoactivity must be attributed to the increase in the exchange rate of the electrolyte species at the ZnS surface. As the studied temperature range and the temperature coefficient are both small, the Arrhenius plot of the described measurements yields also straight lines which have identical correlation coefficients.60 The calculated activation enthalpies for hydrogen formation are equal to 1.23 and 4.55 kcal/mol in the presence of sulfide and sulfite ions, respectively. These values are consistent with an adsorption-desorption-determining process. 6. Influence of the Physicochemical Properties of ZnS Powders on Photoactiuity. The influence of the physicochemical properties of various ZnS powders on their photocatalytic activity has been determined in three different electrolytic media.

(57) “Handbook of Chemistry and Physics”, 58th ed.; CRC Press: Cleveland, OH, 1977; p D-151. ( 5 8 ) Glasstone, S.; Hickling, A. J . Chem. SOC.1933, 36, 829-836.

(59) Butler, M. A,; Ginley, D. S. Nature (London) 1978, 273, 524-525. (60) The accuracy of the measurements does not exclude an eelT dependence.

E 100-

+ W L

PH

-

Figure 11. Rate of hydrogen formation (-) and concentrations of hydrosulfide (-- -) and unprotonated sulfide (---) ions vs. pH of the solution. Reaction temperature: 60 O C . Sample I (400 mg of unplatinized

ZnS) is dispersed in 90 mL of a 0.24 M solution of Na2S.

determined in the case of SO$-and S2-oxidation. In Figure 10 are represented the rate of hydrogen formation of unplatinized zinc sulfide (sample I), dispersed in solutions containing sulfite ions, and the concentration of dissociated SO:- ions as a function of the pH of the solution. In such a medium, the pH dependence of the hydrogen formation is similar to the curve obtained with platinized cadmium ~ulfide.’~A maximal photoactivity is observed between pH 9 and 10. The decrease of the activity observed above pH 10 is attributed to the weaker concentration of protons. Below pH 7.5,the hydrogen production decreases sharply and reaches a zero value at pH 4.5. Since the second pK of the sulfurous acid corresponding to the equilibrium

HS03- ~i H+ + SO3”

5910

Reber and Meier

The Journal of Physical Chemistry, Vol. 88, No. 24, 1984

A

TABLE 111: Influence of the Platinization Method on the Rate of Hydrogen Formation in Various Electrolyte Solutions (Sample I)

Sample V (11 m2/g ZnSl

H, formation rate (mL/h) in 0.24 M

type of photocatalyst A

B C

Na2S and 0.35 M

treatment untreated ZnS ZnS/ 1.7% Pt photochemically reduced ZnS/1.6% Pt thermally reduced ZnS/1.6 Pt thermally reduced and subsequently irradiated in acetic acid

'Reaction temperature:

Irradiation t i m e [minl

-

Figure 13. Hydrogen production (mL) vs. illumination time (min) for 400 mg of unplatinized ZnS dispersed in 90 mL of a 0.24 M solution of Na2S (reaction temperature: 60 "C): triangles, sample V (Ssp= 1.1

m2/g ZnS); squares, sample VI (Ssp= 2.3 m2/g ZnS); circles, sample 111 (Ssp= 119 m2/g ZnS). TABLE II: Influence of the Physicochemical Properties of Various ZnS Samples on the Rate of Hydrogen Formation H2 formation rate (mL/h) in specific 0.35 M 0.24 M Na2S and 0.24 M surface Na2P sample area, m2/g Na2S03 0.35 M Na2S03' V VI I I1 I11 IV

1.1 2.3 6.1 30 119 162

88 37 44 41 134 305

158 128 185 161 169 247

146 146

146

Reference 48. It is shown in Figure 13 representing the hydrogen evolution obtained in the presence of sulfide ions with three different ZnS powders that the hydrogen production is exactly the same for the three samples, independent of the size, the structure, and the impurity content of the particles. In solutions containing S2- ions, the quantum yield of hydrogen formation seems to be independent of the quality of the ZnS sample and the ratio between hydrogen production and disulfide reduction apparently depends only on the respective concentrations of sulfide and disulfide ions. The photocatalytic activity of various ZnS samples with very different physicochemical properties suspended in solutions containing sulfite ions or a mixture of S2-and SO,2- ions is given in Table 11. Results obtained in the S2-/SO?- mixture reveal that, upon addition of S032-ions to the sulfide solution, the rate of hydrogen formation is no longer independent of the quality of the ZnS sample. However, there is no clear correlation between physicochemical properties of the powders and their photoactivity. The weaker rate of hydrogen generation obtained with sample VI is perhaps due to the fact that this powder contains microcrystals of two different crystallographic structures. On the other hand, the higher photoactivity observed with sample IV results from a direct oxidation of sulfite ions to sulfate and dithionate, a process which occurs with a very high rate in the case of this ZnS sample. A quite different behavior is observed in solutions containing only sulfite anions. In such a medium, samples I11 and IV yield significantly higher rates of hydrogen formation. The quantum yield of hydrogen formation is 0.90 for sample IV irradiated with monochromatic light (313 nm) in a 0.35 M Na2S03solution. This remarkable effect can be attributed to the large specific surface

60

0.35 M

NazS03 Na2S03b

0.24 M

Na2Sb

42 44

185 214

166 179

38

124

133

52

165

131

"C. Reference 48.

area of these two Although the degree of photolytical decomposition of the powders described here seems to be slightly dependent on their specific surface area (see section 8), we do not attribute the strong differences of activity observed to this factor but to the size of the ZnS particles and to the area of the semiconductor-electrolyte interface. These results lead to the conclusion that ZnS and CdS powders show quite a different photocatalytic behavior. Photochemical production of hydrogen with suspensions of platinized cadmium sulfide strongly depends on the specific surface area of such powders. Powders of cadmium sulfide with a specific surface area larger than 6.7 m2/g have been found to be very little or not active in all the solutions described in this work.62 However, a similar behavior-increasing activity with decreasing specific surface area-has been observed with ZnS powders of very low specific surface area ((2 m2/g; see sample V and ref 63). 7. Influence of Platinum Deposition. Platinum was deposited on the surface of the ZnS microcrystals in essentially two different ways, i.e. photochemical and thermal reduction of Pt4+ions. The irradiation of ZnS suspensions containing hexachloroplatinic acid changes the color of the powder from white to more or less gray. The analysis of the solutions has shown the disappearance of the ions. The platinum ions and the presence of Zn2+ and Sod2influence of the deposition of platinum particles at the surface of ZnS microcrystals on the rate of hydrogen generation is shown in Table I11 for three different electrolytic media. As expected from the previous section, the photoactivity of unplatinized zinc sulfide is quite high in all three solutions. The additional deposition of platinum by the photochemical way does not significantly improve the charge transfer reaction. On the contrary, the thermal reduction of platinum even decreases the rate of hydrogen formation, except for the oxidation reaction of sulfite ions to sulfate with photocatalyst of type C. This interesting behavior of zinc sulfide is very different from the one observed with cadmium sulfide. In the absence of any charge transfer catalyst (Pt, Rh, Ni, Ru02, etc.) deposited on its surface, CdS does not present a significant activity. Moreover, its photoactivity is strongly dependent on the platinum concentration, the highest activities being observed for concentrations higher than 2% Pt.36 As shown in Figure 14 for two different electrolytic media, the concentration of photochemically reduced platinum has only a small influence on the rate of hydrogen formation. The small Pt concentration dependence suggests that platinum does not really (61) The particularly high quantum yield observed for sample IV may also result from the fact that this sample has been precipitated in an excess of ZnZ+ ions. This procedure could lead to a larger concentration of mobile positively charged point defects, interstitial ZnZt ions, or S2- vacancies. (62) Reber, J.-F., unpublished result. The specific surface area of the investigated CdS samples was comprised between 0.2 and 101 m2/g. (63) The rate of hydrogen formation in 0.35 M Na,S03 has been measured for ZnS samples of very low specific surface area. It is equal to 166, 186, and 130 mL of H2/h for ZnS powders having specific surface areas of 0.60,0.76, and 0.84 m2/g, respectively.


TABLE IV: Molar Fraction of ZnS Photolytically Reduced as Metallic Zinc After 60-min Irradiation in a 0.35 M Solution of Na2S03

ZnS

,501


specific surface

1o4(molar

area, m2/g

fraction of ZnO)

sample V

6.1

IV

5-

2

1

2.8 3.1

* 1st illuminotion 2nd

.

1

t

0.6

30 119 162

I11

6.

0.5 0.4 2.8

t.1 2.3

VI I I1

3rd

I 05

10

15

20

-

Pt concentration [%I Figure 14. Influence of the Pt concentration (wt %) on the rate of hydrogen formation (mL/h) obtained with sample I (400 mg of ZnS, method B) dispersed in 90 mL of a solution containing 0.35 M Na2S.03 (curve 1) and 0.24 M Na2S and 0.35 M Na2S03(curve 2). Reaction temperature: 60 'C. 4

40'~ e

.3 z

v05

OO

10

15

20

25

30

Irradiation time Ihoursl Figure 15. Molar fraction of ZnS (sample I) photolytically reduced as metallic zinc in a 0.35 M solution of Na2S03as a function of the irradiation time (h). Reaction temperature: 55 "C.

act as an eelectron transfer catalyst for the hydrogen formation. Irradiation of zinc sulfide suspended in the presence of a scavenger for positive holes like SO?- ions leads to the formation of reduced species, probably metallic at the surface of ZnS particles. It has been recently shown that this well-known photolytical reduction of ZnS plays a determining role in the hydrogen generation process with colloidal zinc sulfide.65 8. Photolytical Reduction of Zinc Sulfide. It has been recently shown65that, upon irradiation in the presence of a scavenger for positive holes, colloidal zinc sulfide undergoes a photolytical reduction according to ZnS

+ H+ + Ah

hu

Zno

+ HS- + Ah2+

where Ah denotes a scavenger for holes. The aim of this section is to demonstrate the existence of a photolytical reduction process in ZnS powders irradiated in the presence of sulfite ions. The concentration of the reduced species formed has been determined by reacting the irradiated ZnS powder with methylviologen according to the method described in the ~

~~

.6

Irradiation time [hours]

-

10

12

Figure 16. Hydrogen production (L) at 60 "C as a function of the irradiation time (h). Sample IV (1.0 g of unplatinized ZnS) is dispersed in 400 mL of a 1.0 M solution of Na2S03: circles, first run; squares, second run; triangles, third run.

Experimental Section. Assuming that the reduced species are small clusters of metallic we estimated the amount of zinc photolytically reduced assuming that two radical cations of methylviologen are produced per zinc atom. Figure 15 shows that the irradiation of a ZnS powder (sample I) in the presence of sulfite ions really leads to the formation of reducedlreducing species by surface trapping of conduction-band electrons, probably as metallic zinc clusters. A maximal value mol of ZnO per mole of ZnS is reached after equal to 3.5 x 30 min of irradiation. Table IV shows that all the samples used in this work undergo more or less a photolytical reduction to a degree which seems to depend slightly on the specific surface area of the ZnS powders. In hydrogen production experiments with a standard photoreactor, a short induction period of about 2 min, which may correspond to the phase of photolytical formation of Zno nuclei, is observed prior to hydrogen ev~lution.~' 9. Long-Term Stability. A good photocatalyst must show simultaneously a high activity and a sufficient long-term stability. The aim of this section is to examine the efficiency of zinc sulfide for the production of larger amounts of hydrogen over a longer period of illumination. Zinc sulfide of sample IV has been chosen for this purpose because it shows the highest photoactivity for sulfite oxidation. A 1.O-g sample of unplatinized zinc sulfide, dispersed in 400 mL of a 1.0 M Na2S03solution, was illuminated until several liters of hydrogen were produced. As shown in Figure 16 (circles), the rate of hydrogen production slightly decreased from 500 to 475 mL/h after 9.5-h irradiation. At this point, 4.45 L of hydrogen (0.18 mol) had been produced, and the final solution contained about 0.3 ppm Zn, corresponding to a total amount of 0.12 mg of Zn. As the production of 1 mol of hydrogen consumes about 1.3 mol of S032-(see section l), more than half of the initial amount of Na2S03had been oxidized. As mentioned before (see section 3), the decrease of hydrogen generation could result from the decreasing concentration of SO?- ions. However, the deac-

~~~

(64) Sviszt, P. Phys. Status Solidi A 1971, 4, K113-Kl18. (65) Henglein, A. "Book of Abstracts", Jaenicke, W., Ed.; 34th Meeting of the International Society of Electrochemistry, Erlangen, Germany, Sept 1983; p 11.4.

(66) Gergely, G.; Hangos, I. In "Festkorperphysik"; Akademie-Verlag: West Berlin, 1961; p 327. (67) The lamp used in this section was about 10 times less intense than the one in the standard photoreactor.


5912

tivation of the photocatalyst cannot be a priori excluded. Therefore, the irradiated ZnS suspension was centrifuged, redispersed in a fresh solution of 1.O M Na2S03,and subsequently' illuminated under identical conditions. After the production of 5.80 L of hydrogen (0.24 mol) in 12.5 h, the isolation and the redispersion of the photocatalyst were repeated and the suspension of zinc sulfide was irradiated a third time. No deactivation could be detected within the limits of experimental error (see Figure 16, squares and triangles). No trace of zinc68was found in the final solutions of the second and third experiment. The production of 15.8 L of hydrogen in 34 h implies a turnover number of 63 molecules of H2 per ZnS molecule. Assuming that the photolytically reduced zinc acts as electron transfer catalyst and that its concentration is about 3 X mol % (see section 8), this value corresponds to a turnover number of 2 X los molecules of H2 per Zno atom.

Discussion Zinc sulfide is a n-type semiconductor characterized by a wide bandgap of 3.66 eV. This last property makes it very unattractive for solar energy conversion because only light of wavelength shorter than 340 nm is absorbed and, consequently, can induce photochemical reactions. Moreover, it is well-known that ZnS is not stable in aqueous solutions in which it undergoes anodic photocorrosion leading to the formation of sulfur and/or sulfate ions.29,30Our results show that sulfide and sulfite ions are actual hole scavengers which stabilize the ZnS surface very efficiently against anodic photocorrosion. In water at pH 7, the energy level of the valence band of ZnS is located at +1.8 V ("E) and thus lies quite below the standard redox potentials of the anodic reactions used in this work.

-

+ 2h+ SZ2- Eo = -0.52 V ("E, pH 14) S2-+ S032-+ 2h+ Sz032- Eo = -0.12 V S03z-+ 2 0 H - + 2h+ S042- + H 2 0 Eo = -0.92 V 2S032- + 2h+ S2OS2- Eo = -0.25 V 2s'-

-

-

Reber and Meier conduction bands resulting from its wide bandgap confers a poor selectivity to zinc sulfide. Therefore, many sometimes undesirable reactions occur with this semiconductor. When a reaction product formed at one of the poles of the system is not removed from the solution, it may react at the opposite side, so that a cyclic reaction tends to establish itself. Disulfide and thiosulfate ions react according to such a process. The increasing concentration of these ions, which compete which the proton reduction, tends to suppress the hydrogen generation process. Deposition of a metal like platinum at the surface of the microcrystals does no't significantly improve the charge transfer reaction. The quite acceptable charge transfer assured by zinc sulfide is ascribed to the presence of reduced species, probably metallic zinc nuclei, resulting from a cathodic photodecomposition process and acting as electron transfer catalyst. One could argue that the redox potential of the conduction-band electrons of ZnS is so negative that this semiconductor does not need any electron transfer catalyst. However, the fact that zinc oxide behaves identically although the conduction-band edge of this semiconductor is much lower (-0.23 V ("E) at pH 7)72suggests that the photolytical reduction of ZnS plays an important role in the electron transfer process. As shown in Figure 10, protonated sulfite ions are not efficiently oxidized and therefore the reaction has to be carried out in a medium more alkaline than pH 7, where the concentration of dissociated SO:- ions is high. Since hydrogen is generated with a constant rate for hours, the oxidation products of S032-ions, sulfate, and dithionate ions are not competing with protons in the reduction step.73*74 Although the formation of SZOs2-is therions, modynamically less favorable than the production of Sod2half of the SO:- ions yield dithionate ions. According to Henglein and c o - ~ o r k e r swho ~ ~ -have ~ ~ studied the oxidation of SO:- ions, the scavenger action of SO?- ions may essentially consist in the one-hole oxidation

S 0 3 2 - + h+

+ 2h+

-

Zn2+ + S EdWomp(wurzite) = +0.26 V ("E, pH 7); Ed,,,(spha1erite)

The values of the redox potentials and the fact that these reducing agents interact with the ZnS surface explain why they are efficient hole scavengers, preventing the anodic photocorrosion of ZnS. The energy level of the conduction band of ZnS is located a t -1.85 V ("E, pH 7). The corresponding flat-band potential, equal to -1.65 V, is in any case located much above the reduction potential of protons, even in a strong alkaline medium ( vH/H+/i/lHl = -0.83 V ("E) at pH 14). Moreover, recent measurements with a single crystal of ZnS70 have shown that in aqueous solutions its flat-band potential undergoes a cathodic shift linked to the pH according to Vfi (V VS. N H E ) = -1.26 - 0.055pH Although no measurements have been reported heretofore, a cathodic shift of the flat-band potential of ZnS, depending on the concentration of S2-or SO3%ions, as observed for CdS," can be expected, because of the strong interaction of these ions with the crystal surface. The reduction potential of the conduction-band electrons of ZnS should allow an efficient hydrogen generation. However, the very large energy difference between its valence and (68) Zn detection limit: 0.1 ppm. (69) Kroger, F. A. J . Electrochem. SOC.1978,125, 2028-2034. (70) Fan, F.-R. F.; Leempoel, P.; Bard, A. J. J. Electrochem. SOC.1983, -130. - - ,1x66-1875. - - - - - - -.

(71) Inoue, T.; Watanabe T.; Fujishima, A.; Honda, K. Bull. Chem. Soc.

Jpn. 1979, 52, 1243-1250.

SO32S03-

= +0.19 V

SO,-

yielding SO3- radicals which desorb and subsequently react with each other by both combination and disproportionation

The redox potentials of these reactions are also significantly higher than the anodic decomposition potential of Z I I S . ~ ~ ZnS

-

and SO3

-

-

+ H20

s~o~~-

+ so3

~ 0 ~ 2 -

+

S042- 2H+

The yield of hydrogen formation concomitant to the S032-oxidation depends on the quality of the ZnS powders, the specific surface area of the samples being certainly a determining factor. In fact, high yields of hydrogen are obtained with ZnS powders of very large or very low specific surface areas. Regarding their activity, the ZnS powders can be classified as follows: (a) In the range of large specific surface areas, the photoactivity increases with specific surface area. (b) In the range of low specific surface areas, the photoactivity increases with decreasing specific surface area. (c) In the range comprised between these two domains, the yield of hydrogen formation is low and approximately constant. Photochemical reactions at the semiconductor-electrolyte interface can occur if charge carriers generated by absorbed light can reach the interface during their lifetime and find suitable reaction partners, like protons for electrons and donor molecules for holes. Moreover, a low surface recombination rate is an essential condition for the carriers reaching the interface to react with suitable adsorbed species. ( 7 2 ) Gerischer, H. In "Topics in Applied Physics"; Seraphin, B. O., Ed.; Springer-Verlag: West Berlin, 1979; Vol. 31, p 121. (73) Von der Linde, R. Thesis, Marburg, 1902. (74) Coursier, J. Anal. Chim. Acta 1952, 7, 77-94. (75) Lilie, J.; Hanraman, R. J.; Henglein, A. Radiat. Phys. Chem. 1978, 11, 225-227. (76) Henglein, A. Radiat. Phys. Chem. 1980, 15, 151-158.



As the recombination process in the bulk is governed by the diffusion length of the minority carriers, it is expected that the probability for charge carriers to reach the interface increases as the particle size decreases. This behavior is observed in the range of large specific surface areas where the increase of photoactivity with the specific surface area suggests that the efficiency of ZnS is governed by recombination processes in the bulk. In the range of low specific surface areas the ratio bulk to interface is significantly larger, and consequently, the efficiency of the charge separation process depends on the presence of a depletion layer. In the range where the ZnS microcrystals are too thin to build up a complete depletion layer, it is expected that the efficiency of the charge separation increases with increasing size of the semiconductor particles until they reach the thickness of the space charge layer. Such a behavior is observed with our ZnS powders of low specific surface area. Hydrogen formation in solutions containing sulfide anions occurs with an efficient rate a t the beginning of the reaction, independent of the quality of the ZnS sample and of the solution pH. However, disulfide ions resulting from the oxidation of S2ions compete with protons in the reduction process. Therefore, hydrogen is generated with a decreasing rate which depends on the respective concentrations of sulfide and sulfide ions according to

thiosulfate ions.77 However, the quantum yield of hydrogen formation is low in this process due to the cathodic reduction of thiosulfate ions. The addition of H2P02-reducing ions, in which the standard redox potential is -1.65 V (NHE), also prevents the formation of disulfide ions. In contrast to SO?-,sulfide ions, which mediate the oxidation of hypophosphite ions to phosphite, are not consumed as long as the solution contains hypophosphite ions. Since phosphite ions resulting from the anodic reaction are less easily reduced than protons, hydrogen is produced very efficiently and almost linearly with irradiation time. As the reactivity of zinc sulfide powders depends strongly on their specific surface area, the choice of an appropriate photocatalyst (sample IV for instance) is of primary importance for an acceptable efficiency.

r(H2) (mL/h) = K(0.17

+ log ([S2-]/[So] - 1))

Moreover, as sulfide ions absorb light in a large part of the absorption range of ZnS, the flux of photons reaching the ZnS surface is smaller than in the presence of S032-ions. Thus, the production of hydrogen with zinc sulfide from solutions containing sulfide ions is not very attractive. The formation of disulfide ions can be prevented by adding sulfite ions to the Na2S solution. Irradiation of zinc sulfide suspended in solutions containing a mixture of S2-and SO:- ions leads to the formation of thiosulfate ions. In this case, both S2and S032-ions are consumed during the course of the reaction. The rate of hydrogen formation, which is quite high at the beginning of the reaction, tends again to decrease due to the competitive reduction of thiosulfate ions. Thiosulfate ions can also be further oxidized to sulfite and subsequently to sulfate ions. The overall process

s2-+ so3’-+ 5 H z 0

5H2

+ 2So4’-

which is achievable with zinc sulfide can produce 5 times more hydrogen than cadmium sulfide which is not able to further oxidize

Conclusion The irradiation of ZnS powders of appropriate quality produces hydrogen with a very high quantum yield (0.9). The high photoactivity obtained is ascribed to reduced species, probably metallic zinc nuclei, which act as an electron transfer catalyst. The resulting cheap and very active photocatalysts do not need the additional deposition of an expensive charge transfer catalyst like Pt or Ru02. In contrast to other photocatalysts activated by deposition of a noble metal, no visible deactivation occurs when the semiconductor surface is protected against anodic photocorrosion. The wide bandgap of ZnS allows the generation of hydrogen in less reducing media (e.g. thiosulfate). On the other hand, its wide bandgap is a serious handicap in the conversion of solar energy. For this purpose, an efficient spectral sensitization of zinc sulfide has to be achieved. Investigations in this direction are in progress in our laboratories. Acknowledgment. We express our gratitude to Dr. R. Kubler and Mr. J. Pave1 for the numerous analyses. We also thank Dr. A. Burkhard for the X-ray diffraction measurements and Mr. R. Muller for BET measurements on ZnS powders. The careful experimental work of Ms. C. Schobi is greatly appreciated. Registry No. MV+, 25239-55-8; ZnS,1314-98-3; Na2S03,7757-83-7; S2-, 18496-25-8; HZPOT, 15460-68-1; S2032-, 14383-50-7; H2, 1333-74-0;

S;-, 16734-12-6; HPO3-, 38201-36-4; P043-, 14265-44-2; SO>-, 14808-79-8; S202-, 14781-81-8; H20, 7732-18-5: Pt. 7440-06-4: Zn. 7440-66-6; methylviologen, 1910-42-5. (77) Borgarello, E.; Desilvestro, J.; GrBtzel, M. Helv. Chim. Acta 1983, 66, 1827-1834.

Photochemical production of hydrogen with zinc sulfide suspensions

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