The removal of substituted phenols by a photocatalytic oxidation process with cadmium sulfide

Wat. Res. Vol. 24, No. 5, pp. 543-550, 1990 Printed in Great Britain.All rights reserved

0043-1354/90 $3.00 + 0.00 Copyright© 1990Pergamon Press plc

THE REMOVAL OF SUBSTITUTED PHENOLS BY A PHOTOCATALYTIC OXIDATION PROCESS WITH CADMIUM SULFIDE* ALLEN P. DAVISand C. P. HUANGt @ Department of Civil Engineering, University of Delaware, Newark, DE 19716, U.S.A. (First received June 1988; accepted in revised form October 1989)

Abstract--The visible light photo-irradiation of aqueous cadmium sulfide produces conduction band holes that oxidize organic compounds. The oxidation rates of some substituted phenols depend upon the type of substituent as well as the degree of substitution. These rates do not correspond with those expected for electrophilic oxidation. Evidence for the rate dependence on the degree of phenol adsorption is presented which includes a correlation with the octanol-water partition coefficient as well as pH-surface charge dependencies. Key words--photocatalytic oxidation, cadmium sulfide, phenols, chlorophenols, oxidation, surface

reactions

INTRODUCTION

Phenols and phenolic compounds are common pollutants of the aquatic system. All are soluble to a certain degree; especially in the deprotonated form. The chlorination of natural waters for disinfection produces chlorinated phenols due to the partial oxidation of the aromatic macromolecules that constitute natural organic matter (Larson and Rockwell, 1979). Additionally, chiorophenols occur as products in the partial degradation of pesticides (Crosby and Wong, 1973). Pentachlorophenol was commonly utilized as a wood preservative. Anthropogenic sources, such as heavy industry and mining, also contribute chloro- and nitro-phenols to the aqueous system (Wegman and Van Den Broek, 1983; Milner and Gouldner, 1986). Activated carbon adsorption is the generally accepted method for phenolic pollutant removal from the aqueous phase. However, this process produces spent carbon as a waste by-product. Therefore, it is most desirable to examine mineralization technologies for detoxification that chemically treat the pollutant, i.e. oxidation. Recently, a considerable amount of research has been undertaken examining the ability of colloidal semiconductors, under appropriate illumination, to oxidize toxic organic compounds in water and hence detoxify the solution (Barbeni et al., 1987; Matthews, 1987). Virtually all of this work has been completed using titanium dioxide as the photocatalyst due to the fact that under the experimental conditions used, this *Paper presented at the International Conference on Physicochernical and Biological Detoxification of Hazardous Wastes, Atlantic City, N.J., 3-5 May 1988.

tAuthor to whom all correspondence should be addressed.

material is the most active. However, the greatest drawback arising from the utilization of TiO2(s) is that with a bandgap of 3.0 eV, it is activated by light only in the ultraviolet range. Therefore, any practical application must utilize artificial u.v. light, or rely on the small percentage of u.v. light that is available in solar irradiation. It is for this reason that we are examining cadmium sulfide as an alternative photocatalyst. CdS(s) contains a bandgap corresponding to excitation by visible light (2.4eV) and has been shown to photocatalytically oxidize phenol (Davis and Huang, 1989). Cadmium sulfide is thermodynamically stable with a solubility product of 10 -28.3 (Garrels and Christ, 1978). Semiconductor materials absorb light energy which results in the excitations of electrons from the valance band of the crystal to a higher energy state in the conduction band (Gratzel, 1983). To preserve electroneutrality, an electron vacancy or hole remains in the valence band. Under the electric field created by the charge transfer resulting from immersing an n-type semiconductor in aqueous solution, the holes are energetically attracted to the particle surface where they act as strong oxidizing agents. The photocatalytic oxidation of phenol and a few chlorinated phenols in TiO2(s) dispersions have been examined independently by other researchers. A thorough study of phenol oxidation has been presented by Okamoto et al. (1985a, b). Barbeni et aL (1984, 1986a, b, 1987) and Matthews (1986, 1987) examined the photocatalytic oxidation of some chlorinated phenols, but no mechanistic interpretation of the variation in rate with substitution was presented. Additionally, variable parameters such as light spectrum and intensity as well as reactor geometry make comparisons between researchers impossible and

543

544

ALLEN P. DAVIS and C. P. HUANG Table I. Properties of substituted phenols Compound Phenol 2-Chloro 3*Chloro 4-Chloro 2,4-Dichloro 2,6-Dichloro 3,4-Dichloro 2,4,6-Trichloro Pentachloro 3-Bromo 3-Nitro 3-Methyl

UV-VIS 2 (nm) 271 274 274 280 284 276 283 288 291) 275 310 272

Oxidation rate log Ko,t @ pH 7 ( × 10 5)++ 1.46 1.88 2. t 7 2.18 2.49 2. I0 2.42 1.56 2.75 (3.17§) 2.40 2.77 2.25 3.38 238 3.38 (3.66i) 2.54 5.01 4.3$; 2.63 t 34 2.00 0.40 1.99 0.34

pK* 9.9 8.3 8.8 9.1 7.8 6.8 8.6 6.1" 4.7* 8.8 8. I 10.1

*Martel and Smith (1974). tLeo et al. (1971). ~M/h. §Schellenberg et al. (1984). IbXie and Dyrssen (1984). i n c o n s i s t e n t initial c o n c e n t r a t i o n s i m p e d e c o m p a r i sons o f w o r k by the s a m e a u t h o r s . A n i n v e s t i g a t i o n into the m o d i f i c a t i o n in r e a c t i o n kinetics resulting from various p h e n o l s s h o u l d p r o v i d e valuable insights into the m e c h a n i s m o f such a p h o t o c a t a l y t i c r e a c t i o n o n CdS(s). It is necessary to d e t e r m i n e if this m e c h a n i s m is identical to that a c c e p t e d for TiO2(s ) as well as to d o c u m e n t the variances d e m o n s t r a t e d by the different p r o p e r t i e s o f the o r g a n i c pollutant. Such a p r o c e s s m a y have i m p l i c a t i o n s in w a t e r a n d w a s t e w a t e r t r e a t m e n t systems as well as in the d e c o n t a m i n a t i o n o f h a z a r d o u s waste sites. MATERIALS AND METHODS The phenols were obtained from the Aldrich Chemical Co. with 98% purity or greater. The cadmium sulfide is powdered, electronic grade, 99.999% purity (Aldrich). The BET surface area is found to be 1.5 m2/g. A 10 3M phenol solution was prepared with an ionic strength of 5 × 10 -2 M Na:SO 4. 2,4,6-Trichlorophenol and pentachlorophenol were initially prepared in basic solution to enhance solubility. A catalyst concentration of 5 g/l was used. The solution pH was adjusted with NaOH or H2SO 4. The reaction vessel was identical to that described previously (Davis and Huang, 1989). The illumination was provided by 300W ELH projection lamps (General Electric). ELH lamps provide a spectrum near that of natural sunlight (Matson et al., 1984). The vessel was double-jacketed with temperature controlled at 25°C. A Pyrex glass cover was used to allow maximum visible light passage. A light intensity of 700 W/m s was employed as measured by a YSI radiometer. Oxygen was constantly purged through the system. 02 is necessary for efficient phenol oxidation by scavenging the photo-produced electrons (Davis and Huang, 1989). The organic concentration is measured on a Hitachi Perkin-Elmer UV-VIS spectrophotometer at a wavelength of maximum absorption (Table 1). All compounds except pentachlorophenol (PCP) were acidified to the protonated species before quantification. PCP determinations were completed on the phenolate ion at pH approx. 7 due to solubility difficulties. Periodic samples were analyzed on a Hewlett-Packard Model 5890/5970 GC/MS to detect any intermediate products and as a check on the u.v. determinations using direct aqueous injection of the acidified sample. The following are the GC/MS analytic parameters: the column is a HP-1 capillary column (cross-linked methyl silicone gum) with a 12 m length, i.d. 0.2 mm and a film

thickness of 0.33 gm. The temperature range was from 100 to 300°C at a ramp rate of 25°C/min over a run time of 8.5 min. The injection and detector temperatures were 200 and 280°C, respectively. Helium served as the carrier gas at a flow rate of 21.6 cm3/min with the splitless on. Initial oxidation rates were estimated from the slope of the removal curve. Carbon dioxide is measured by the precipitation of BaCO3 from gas purged through the vessels and into a series of 3 sealed flasks containing Ba(OH)2/NaOH. Chloride was detected by an argentometric titration method (APHA, 1980). Adsorption studies were completed in Pyrex glass test tubes also at a concentration of 5g/l CdS(s), 10-3M organic. The tubes were completely wrapped in aluminum foil to avert any light impingement and prevent any oxidation. They were allowed to equilibrate by shaking overnight. After pH measurements, the samples were filtered. The supernatants were acidified and the remaining concentrations of phenols measured. In each case the concentration is referred to a blank submitted to identical conditions. RESULTS F i g u r e 1 d e m o n s t r a t e s t h a t the r e m o v a l is a p h o t o sensitive process. A lighted p h e n o l i c s o l u t i o n s h o w s small o r g a n i c toss due to p h o t o c h e m i c a l d e c o m p o sition a n d / o r volatilization. A d d i t i o n o f CdS(s) to the

3 - CnLorophenot O O-,,

1.0 ju~ u ~ . j u _

o

0.Q 8 0,6

o

0.4

;: 2L,,

Q2

tx CotoLyst w / L i g h t

NO

'ght

/ 0

i

0

[ 8

~_L_

Z

16

J-

24

l 32

I

I 40

l

[ 48

Time (h)

Fig. 1. Photocatalytic oxidation of 3-chlorophenol using CdS(s). Experimental conditions: CdS(s) 5g/l; phenol 10-3M; light intensity 700W/m2; pH 7; ionic strength 0.05 M Na2SO4; temperature 25°C.

Photocatalytic oxidation of substituted phenols

100

200

545

300

m/z Oh

'2

'3

14

15

'3

'4

'5

'4

'5

2h

'2 8h

'2

'3 Time [min)

Fig. 2. Mass spectrum and chromatogram for photocatalytic oxidation of 2,4,6-trichlorophenol under high light intensity. Experimental conditions: CdS(s) 5 g/l; phenol 10-3 M; light intensity 2000 W/m2; pH 7; ionic strength 0.05 M Na2SO4; temperature 25°C. dark phenolic solution demonstrates an initial removal due to adsorption, but no continuous removal as compared with lighted, catalytic sample. The adsorption is rapid enough that equilibrium occurs by the initial data point, i.e. within the first half hour of the reaction time. Therefore, any deviation pH 7 1.0 08

o Phenot D 2- ChLorophenot ~ 2.4-Dichtoroph.not v 2,4,6 - Trichtorophenot

~ ~'x~Ct-' ~ ] ~ - = , ~ - ~ ~

o 0.6 o

• pe°too.toro...oL

~

~

u O.4 0.2

0

i

I 8

t

i 16

i

I h 24 Time (h)

I 32

l

I 40

t 48

Fig. 3. Photocatalytic oxidation of substituted phenols using CdS(s). Experimental conditions: CdS(s) 5g/l; phenol 10-3M; light intensity 700W/m2; pH 7; ionic strength 0.05 M Na2SO4; temperature 25°C.

from the first reading is attributed to oxidation. Over the course of the photocatalytic oxidation process, large amounts of bases are needed to maintain the pH at its original value. Results from GC/MS examination revealed only the original phenol compound in detectable quantity. Reaction intermediates were not isolated either due to the specific analytical methodology used, because they are present in undetectable concentrations, or because they are in some degree adsorbed to the catalyst. A mass spectrum and chromatogram for 2,4,6-trichlorophenol removal under high light intensity is shown in Fig. 2. Results after 8h of reaction show 80-93% of the removed phenol was obtained as CO2 and 97-98% recovery of the stoichiometric chloride from 2,4,6-trichlorophenol. Figure 3 shows the relative removal of chlorinated phenols at pH 7. The concentration differences at short time durations, e.g. half hour, are influenced to a certain degree by adsorption reaction that takes place rapidly onto the CdS(s) surface. In general, the greater the chlorine substitution, the higher the extent of oxidation; pentachlorophenol being oxidized most efficiently, while the unsubstituted phenol is the most

546

ALLEN P. DAVISand C. P. HUANG Meto

2 - ChLorophenoL

substituted

phenols

1.0

1.0

0.8 0.6

o

(-~ 0 . 4

o6°'i °

o 0.4

OpH7

0.2

v 3- NitrophenoL 0 / 0

I

I

I 8

I 16

J

I 24

I

I 32

I

~ 40

0.0

I 48

q 0

I 8

J -

i 16

f i

I 24

Time

T i m e (h)

I

I 32

I

l 40

~

J 48

(h)

Fig. 4. The effect of pH on the photocatalytic oxidation of 2-chlorophenol. Experimental conditions: CdS(s) 5 g/l; phenol 10 3 M; light intensity 700W/m2; ionic strength 0.05 M Na2SO4; temperature 25°C.

Fig. 6. Photocatalytic oxidation of recta-substituted phenols. Experimental conditions: CdS(s) 5g/l; phenol 10-3M; light intensity 700W/m2; pH 7; ionic strength 0.05 M Na2SO4; temperature 25°C.

resistant to degradation. Therefore the increase of the electron withdrawing groups produces an easier substance to oxidize. Similar results are found at other solution pH values. The effect of solution pH for 2-chlorophenol photocatalytic oxidation is presented in Fig. 4. Comparable results are found for other compounds such as 2,4-dichlorophenol, 3,4-dichlorophenol and 2,4,6trichlorophenol. For the most part, solutions with a lower pH provide a more efficient oxidative removal.

It is clear that at very high pH, i.e. 11, the oxidation process is consistently severely inhibited. A comparison of the removal of some isomers of mono- and dichlorophenols is depicted in Fig. 5. Oxidation kinetics of various meta-substituted phenols are presented in Fig. 6. The rates of oxidation for the halogen substitutions are greater than those for both the nitro and the methyl replacements.

The photocatalytic oxidation of organic compounds is mechanistically described as follows: the excitation of the semiconductor by bandgap illumination creates conduction band electrons and valence band holes (Gratzel, 1983):

M o n o - chLorophenols 1.0

o

(o)

0.8

0.6

o

DISCUSSION

CdS(s) + hv --, CdS(s) + e + h ~.

( 1)

The electronic gradient produced by solid-solution interfacial equilibration may attract the holes to the surface; otherwise they may recombine with the electron:

OA

0.2 GO

I 0

I 8

I

I 16

I 24

I

Time

k

I 32

I

l 40

I 48

(h)

1.0

0.8 e

(b)

o 2,6

Dichlorophenol

2h+rf + CdS(s) ~ Cd 2+ + S.

-. 0.6 0.4 Q2 0.0 / O

~

I 8

I

I 16

I

I 24

~

(2b)

With cadmium sulfide the possibility of photocorrosion also exists (Henglein, 1982):

o 2 , 4 Dichlorophenol

~

(2a)

h ÷ + e ~ recombination.

pH 7

~.

h ÷ --, h~urt

I 32

I

i 40

I 48

Time (h)

Fig. 5. Photocatalytic oxidation of isomers of mono- and dichlorophenols. Experimental conditions: CdS(s) 5g/l; phenol 10-3M; light intensity 700W/m2; pH 7; ionic strength 0.05 M Na2SO4; temperature 25°C.

(3)

Work on the photo-oxidative dissolution of CdS(s) has been conducted by Hsieh (1988). The electron is forced into the bulk of the solid, and to maintain electroneutrality, eventually reacts at a surface defect in a dark section of the semiconductor (Kobayashi et al., 1983). Oxygen will scavenge this electron, becoming reduced (Harbor and Hair, 1978) and thus averting recombination: 02 + e- ~ Of O2+e

~O~-.

(4a) (4b)

Photoeatalytic oxidation of substituted phenols One possible fate for the surficiai hole is to react directly with an adsorbed organic compound as proposed for benzene oxidation (Hashimoto et al., 1984): OH

OH

h~urf + ~

_._

Cl

~

(5)

CI

Subsequent reaction with water produces a dihydroxyphenolic radical (Hashimoto et al., 1984): OH

OH

+

H20

H+

CI'

C

(6)

H

Repetition of hole oxidation and water reaction is similar to hydroxyl radical oxidation and complete mineralization should follow a pathway similar to the hydroxyl mechanism. With TiO2 (s) it is assumed that the hole will react with adsorbed water to form hydroxyl radicals (Okamoto et al., 1985a):

hs+rf Jr" H20 --, OH" + H +

(7)

The hydroxyl radical, a strong oxidizing agent, can react with the phenol, i.e. OH

OH

1 OH* + ~ CI

(8)

t, C1

OH

The eventual products of reactions 6 and 8 are identical. The production of acid over the course of the reaction is explained by reactions 6 and 7. The superoxide from the oxygen reduction (reaction 4) forms hydrogen peroxide and these substances act as homogeneous oxidizing agents for the toxic compounds (Okamoto et al., 1985a; Bieski et al., 1985): 02- + H + --, HO~

(9a)

0 2 + HO~ --+O: + H O f

(9b)

HO2 + H ÷ --+H202.

(9c)

The rates of homogeneous oxidation of aromatic organic compounds are dependent upon the electron density of the aromatic ring which is altered by ring substitution. Halogen substitutions as well as nitro functional groups are electron withdrawing that extract electron density from the ring. The methyl group is electron donating which increases the aromatic electron density and renders the compound more susceptible towards the electrophilic characteristics of oxidizing agents, resulting in a larger oxidation rate. Conversely, slower rates are noted with

547

electron withdrawing substituents. These electronic characteristics are quantified by the Hammett relationship which relates aromatic reaction rates to thermodynamic substituent effects and assigns quantitative values to the ring substituent groups. The reactivity of aromatic compounds towards oxidation by ozone (Hoign6 and Bader, 1983a), hydroxyl radicals (Anbar et al., 1966) and even direct anodic oxidation (Suatoni et al., 1961) have each been shown to follow the Hammett relationship. The Hammett substituent values are additive for multiple replacements, which explain why compounds with multiple electron withdrawing group substituents, such as pentachlorophenol, are very resistant towards oxidation. However, the present photocatalytic oxidation data is inconsistent with that of direct oxidation theories. Pentachlorophenol and trichlorophenol show the best oxidative removal, while methyl-phenol is one of the most persistent compounds. It should be noted that the oxidation rate by ozonolysis increasing in the order: trichlorophenol > dichlorophenol > chlorophenol has been presented (Gilbert, 1976). However, this order was obtained by considering only one solution pH in which the phenol/phenolate ratio for each compound was different. The higher chlorinated species were removed faster due to the higher concentration of the more reactive phenolate. Nucleophilic substitution would demonstrate a difference in rates between ortho and para in comparison to that of meta substitution due to the conjugation available in the former isomers. Similarly, a large difference would be expected between rates of nitro and methyl substituents, both of which are not found. The adsorption of substituted phenols onto CdS(s) has been determined (Huang et al., 1988) and adsorption density is plotted against initial oxidation rates of the substituted phenols (Fig. 7). Some relationship between the two seems apparent. The highly substituted phenols which are more strongly adsorbed have

/ /

%3

/

x

n

~, / t a /

21

yQ

o

/

~o

=1

0

f/ 0

/~

/ t

mpH5 "" pH 7

/

•

epH9 opH

A I 10

J

I 20

i

I 30

t

i 40

t

11 I 50

I 60

Adsorption density [ / ~ M / m 2)

Fig. 7. Photocatalytic oxidation rate of various phenols as a function of their dark adsorption density.

548

ALLENP. DAVISand C. P. HUANG

the largest photocatalytic oxidation rate. In most cases, deviations from the linear relationship are towards more efficient oxidation which may result from a small amount of OH" production (reaction 7) by the semiconductor particles. It has been noted that the adsorption of noncomplexing organic compounds onto solid surfaces is associated with the hydrophilicity of the compound (Westall, 1987). The adsorption is therefore a function of both the octanol-water partition coefficient, Kow, of the organic chemical as well as the organic carbon content of the adsorbent, foc: R ---~Rads;

Kp

(10)

log Ko = a log Kow + log foc + b.

OH

CI

CI

The rate dependence on adsorption as shown in Figs 7 and 8 demonstrate that this equilibrium is established. This will result in a Langmuir type rate dependence that has been found for phenol oxidation on both CdS(s) (Davis and Huang, 1989) and TiO2(s ) (Okamoto et al., 1985a): rate = k [phenolads]

k [phenol] 1 + k [phenol]"

i I

0 0,0

I 0.:5

~

I 1.0

4

i 1.5

,~ 2.0

Log

t

I 2.5

t

t 3.0

J

t 3.5

t 4.0

Kow

Fig. 8. Photocatalytic oxidation rate as a function of octanol-water partition coefficient.

(11)

The adsorption of unsubstituted phenol onto hydrous surfaces such as TiOz(s) (Okamoto et al., 1985a), FeOOH (Yost and Anderson, 1984), MnO2 (Stone, 1987) and CdS(s) (Park, 1987) has been found to be quite small. However, an increase in the degree of chlorination results in a larger extent of adsorption of molecular phenols onto sediments (Schellenberg et al.. 1984). This adsorption is explained using equations (10) and (11) and expounding that the octanol-water partition coefficient increases with the increase in chlorine substitution. A plot of initial rates against the octanol-water coefficient (Table 1) is given in Fig. 8. The rate for both pH 5 and 7 for most compounds is plotted where the phenols are in the molecular form. Although a linear relationship may not be justified, a qualitative trend of oxidation rate corresponding to Ko~ is noted. It therefore must be concluded that the charge transfer between the phenolic compounds and the semiconducting surface is the rate limiting step in the oxidation of phenols on illuminated CdS(s) (reaction 5). The rate is limited by the surface concentration of organic molecules and is dependent upon an adsorption equilibrium on the surface before this step. OH

/

(13)

The rate constant, k, is a function of oxygen concentration, light intensity and catalyst concentration (Davis and Huang, 1989). This kinetic development

for the charge transfer is similar to that demonstrated by Stone and Morgan (1987) for the electron transfer between the solid and adsorbed organic in the reductive dissolution of manganese oxides. Weak, nonspecific interactions are responsible for the phenol adsorption of CdS(s) (Park, 1987; Huang et al., 1988) and this will inhibit the charge transfer. The differences in rates between the various substituents is relatively small, less than one order of magnitude. In contrast, the substituent effect of the reductive dissolution of manganese oxides, which was found to be dependent upon the electronic characteristics of the aromatic ring, results in a rate variation of three orders of magnitude (Stone, 1987). A trend of increasing CO2 production rate with increasing chloride substitution was noted by Matthews from the decomposition of many organic compounds under illuminated TiO2(s) (Matthews, 1986); results similar to those were observed in this study. Other workers, however, have presented TiO2(s) results that follow the tendencies expected for electrophilic attack. Under comparable conditions, Barbeni et al. (1986b) noted a longer half-life for pentachlorophenol as compared to 4-chlorophenol. Similarly, less time was needed to half the concentration of 3,4-dichlorophenol vs that for the more highly substituted 2,4,5-trichlorophenol, all at pH 3. Similarly, in a recent work by Matthews (1987), apparent first order rate constants for 2 monochlorophenols were less than that obtained for the unhalogenated compound. The adsorption dependent hypothesis is also supported by the effect of pH. The mono-substituted phenols show little change in degree of adsorption except at high pH where phenol deprotonation occurs. The oxidation of trichlorophenol exhibits the most pronounced pH dependence and also demonstrates the greatest degree of variation in adsorption over the pH range at which oxidation was examined. An examination of the pH dependence of the oxidation rate is depicted in the log (rate) vs pH plot (Fig. 9). The order with respect to proton concentration (slope) is low; 0.1 or less for neutral pH values

Photocatalytic oxidation of substituted phenols -4.4

549

becomes negatively charged, and the phenolate becomes the dominate organic form.

-4.6

A

-4.8

CONCLUSIONS n 2- chLoro

- 5.0

3 , 4 - dichLoro

X~

• 2,4,6 - t r i c h L o r o

-5.2

"

X

~

- 5.4 -5.6

4

I 5

i

I 6

i

I 7

I

I 8

I

I 9

J

I

10

I

11

pH

Fig. 9. Relationship between pH and the photocatalytic oxidation rate of substituted phenols.

(5 ~
(14)

where Ep, 0 is the reduction potential in a solution at pH equal to zero. Similarly, the photo-corrosion of CdS(s) via holes is increased by a decrease in pH. The energetics of the semiconductor bands denoted by equation (14) result in a very small pH dependence and may be the cause of the slight increase of phenol oxidation at low pH. Pentachlorophenol as well as 4-chlorophenol photocatalytic oxidation using TiO2 (s) both exhibited faster rates in alkaline solution, where each was in the anionic form, in comparison to the rates of the undissociated molecular compound at lower pH (Barbeni et aL, 1984, 1986a). These results are consistent with hydroxyl radical formation on the TiO2(s) catalyst and are in direct contrast to the data of the present work. The oxidation rates do not demonstrate abrupt changes as the suspension pH varies on either side of the organic pKa value which would occur if homogeneous oxidation is the rate determining step. Ozonation rates of the conjugate base of the phenols are orders of magnitude faster than that of the protonated species (Hoign6 and Bader, 1983b). The reaction on CdS(s) at high pH is slower, consistent with the decrease in adsorption as the CdS(s) surface

It is generally accepted that the photocatalytic oxidation of organic compounds in the presence of titanium dioxide proceeds via the formation of OH" radicals from adsorbed water reacting with the surficial holes. In contrast, as is demonstrated here for photocatalytic oxidation using cadmium sulfide, the organic compound must be adsorbed onto the semiconductor for the reaction to occur, whereby it reacts directly with the holes. It is possible that cadmium sulfide is not hydroxylated to the degree of oxides and not able to produce sufficient amounts of OH" for this mechanism to dominate. Another rationalization may be that CdS(s), being a visible light semiconductor, has a valence band potential of only 1.9 V (NHE) (Maruska and Ghosh, 1978). This value corresponds to the potential of the surficial holes. TiO2(s) has a more anodic valance band energy of 3.05 V (NHE) (Maruska and Ghosh, 1978). The formation of OH" species requires potentials (Hickling and Hill, 1950) which may not be thermodynamically possible with cadmium sulfide: h + + H20--* OH" + H+ E0 = 2.8 V (NHE)

(15)

h + + H O - ~ OH'E0 = 2.0 V (NHE).

(16)

Under low pH conditions, the oxidation of the cadmium sulfide becomes significant and aqueous cadmium is present. Methods to inhibit this corrosion are currently under investigation. Other compounds which are known visible light semiconductors have been investigated, but extremely small or no photocatalytic oxidation has been measured. Acknowledgements--The research on which this paper is

based was supported in part by the United States Department of the Interior as authorized by the Water Research and Development Act of 1978 (P.L. 95.467). Contents of this publication do not necessarily reflect the views and policies of the United States Department of the Interior, nor does mention of trade names and commercial products constitute their endorsement by the U.S. Government. We thank Jaime Tseng for help with the GC/MS. The award of a University Competitive Fellowship to A.P.D. by the graduate office of the University of Delaware is greatly appreciated. REFERENCES

American Public Health Association (1980) Standard Methods for the Examination of Water and Wastewater.

American Public Health Association, Washinton, D.C. Anbar M., Meyerstein D. and Neta P. (1966) The reactivity of aromatic compounds towards hydroxyl radicals. J. phys. Chem. 70, 2660-2662. Barbeni M., Parmauro E., Pelezzeti E., Borgarello E., Gr~itzel M. and Serpone N. 0984) Photodegradation of 4-chloropbenol catalyzed by titanium dioxide particles. Nouv. J. Chim. 8, 547-550.

550

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