J . Phys. Chem. 1987, 91, 3328-3333
3328
Photooxidation of Organic Impurities in Water Using Thin Films of Titanium Dioxide Ralph W. Matthews Division of Energy Chemistry, Commonwealth Scientific and Industrial Research Organization, Sutherland, N S W 2232, Australia (Received: October 20, 1986)
Results of the destruction of organic solutes in a simple, thin film TiO, reactor are described. The reactor was illuminated with a 20-W blacklight UV fluorescent tube and the aqueous stream containing the organic solute flowed past the stationary photocatalyst. In the continuous recirculation mode, the destruction rate of each solute obeyed approximately first-order kinetics. The reaction rate constant decreased with increasing solute concentration. The times for 50% destruction of 500 cm3 of 10 pM solutions of each of the solutes salicylic acid, phenol, 2-chlorophenol, Cchlorophenol, benzoic acid, 2-naphtho1, naphthalene, and fluorescein were 7.1, 7.2, 8.2, 8.7, 6.9, 8.5, 4.3, and 6.4 min, respectively. It was found that the observed apparent first-order dependence and the change in rate constant with concentration could be explained in terms of the integrated form of the Langmuir adsorption isotherm. A marked dependence of the destruction rate on flow rate was observed and an expression developed which allows the calculation of the destruction curve with good precision at any solute concentration and flow rate. A corresponding curve was observed for the formation of carbon dioxide from salicylic acid solution. It was shown that hydroxylation of the aromatic ring to give salicylic acid is a minor reaction path in the destruction of benzoic acid. The maximum quantum yield for the destruction of salicylic acid at 25 OC was found to be 0.022. The activation energy for the photooxidation of salicylic acid was determined to be 11.0 h 0.8 kJ mol-'.
Introduction
TABLE I: Analytical Wavelengths for Solutes Studied
Many organic compounds are decomposed in aqueous solution in the presence of titanium dioxide powder illuminated with near-UV lightl4 and there is interest in this process as a means of purifying Although phenols are formed as intermediate compounds in the p r o c e s ~ , ' ~these - ~ ~ and other organic compounds including chlorinated organic compounds have been shown to be completely mineralized to C 0 2 and Furthermore, solar illumination may also be used to drive the reaction. The Langmuir adsorption isotherm has been invoked to explain many of the r e s ~ l t s . ~ - ' ~ It * ~has ' + also ~ ~ been reported recently that the destructions of phenolz4 and salicylic acidz5 follow
H20.9J09143'69z0,2'
compd salicyclic acid phenol 2-chlorophenol 4-chlorophenol benzoic acid 2-naphthol naphthalene fluorescein
L, 294
267 264
278 235 272 27 5 486
A,"' 412 300 298 313 315 356 335 570
149zz
(1) Carey, J. H.; Lawrence, .I.; Tosine, H. M. Bull. Enuiron. Contam. Toxicol. 1976, 16, 697. ( 2 ) Kraeutler, B.; Bard, A. J. J. Am. Chem. SOC.1978, 100, 5985. (3) Izumi, I.; D u m , W. W.; Wilbourn, K. 0.; Fan, F.-R. F.; Bard, A. J. J. Phys. Chem. 1980, 84, 3207. (4) Kawaguchi, H.Enuiron. Technol. Left. 1984, 5 , 471. (5) Oosawa, Y.J. Phys. Chem. 1984,88, 3069. (6) Matthews, R. W. Water Res. 1986, 20, 569. (7) Hustert, K.; Kotzias, D.; Korte, F. Chemosphere, 1983, 12, 55. (8) Herrman, J.-M.; Mozzanega, M.-N.; Pichat, P. J. Phorochem. 1983, 22, 333. (9) Hsiao, C . - Y . ;Lee, C.-L.; Ollis, D. F. J. Catal. 1983, 82, 418. (10) Pruden, A. L.; Ollis, D. F. (a) Enuiron. Sci. Technol. 1983, 17,628; (b) J. Catal. 1983, 82, 404. (11) Ollis, D.F.; Hsiao, C . - Y . ;Budiman, L.; Lee, C.-L. J. Cufal. 1984, 88, 89. (12) Nguyen, T.; Ollis, D. F. J. Phys. Chem. 1984, 88, 3386. (13) Ollis, D. F.Enuiron. Sci. Technol. 1985, 19, 480. (14) Barbeni, M.; Pramauro, E.; Pelizzetti, E.: Borgarello, E.; Gratzel, M.; Seroone. N. Nouu. J. Chim. 1984. 8. 547. (15) Hidaka, H.; Kubota, H.; Gratzel, M.; Serpone, N.; Pelizzetti, E. N o w . J. Chim. 1985, 9, 67. (16) Barbeni, M.; Pramauro, E.; Pelizzetti, E.; Borgarello, E.; Serpone, N. Chemosphere 1985, 14, 195. (17) Matthews, R. W. J. Chem. SOC.,Faraday Trans. 1 1984, 80, 457. (18) Hashimoto, K.; Kawai, T.; Sakata, T. J. Phys. Chem. 1984,88,4083. (19) Takagi, K.; Fujioka, F.; Sawaki, Y . ;Iwamura, H. Chem. L t f . (Chem. SOC.Jpn) 1985, 913. (20) Okamoto, K.-I.: Yamamoto, Y . ;Tanaka, H.; Tanaka, M.: Itaya, A. Bull. Chem. SOC.Jpn. 1985, 58, 2015. (21) Matthews, R. W. J. Catal. 1986, 97, 565. (22) Ahmed, S.;Ollis, D. F. Solar Energy 1984, 32, 597. (23) Matthews, R. W. Ausr. J . Chem., in press. (24) Okamoto, K.-I.; Yamamoto, Y . ;Tanaka, H.; Itaya, A . Bull. Chem. SOC.Jpn. 1985, 58, 2023
0022-3654/87/2091-3328$01.50/0
first-order kinetics and that the apparent rate constant depends on the initial concentration of solute. Since the effect of solute concentration is of obvious importance in any process for the removal of solutes from water, it is of interest to determine the basis for the observed first-order dependence and to determine to what extent other solutes might behave similarly. It is also evident that in any water purification process, filtration and resuspension should be avoided if possible and, in the case of titanium dioxide, this can be done by fixing the photocatalyst to a stationary support. It was therefore an object of the present investigation to maintain the photocatalyst as a stationary phase, use water as the mobile phase, and determine the rates of destruction of various solutes at different concentrations. Experimental Section
Chemicals used were at least laboratory reagent grade and were obtained from reputable suppliers. Degussa P25 titanium dioxide, primarily anatase, BET surface area 50 m2 g-], 30 nm mean particle size was used as the photocatalyst, and 400 mg was suspended in 80 cm3 of water by sonication and sucked by vacuum into a 7-m-long borosilicate glass tube, 4 mm i.d., 80 cm3volume, wound in a 65-turn spiral. Approximately 5 cm3 of the suspension remained in the spiral on draining. The end of the spiral was closed and vacuum applied while warm air from a hair dryer was blown around the spiral. The procedure was repeated to build up successive layers of TiOz on the inside surface of the tube. A 20-W NEC blacklight fluorescent lamp (TIO), 32.5 mm diameter, 588.5 mm long, fitted snugly into the spiral and the whole unit was mounted in a standard domestic 20 W fluorescent lamp holder. Solutions (500 cm3) were circulated through the spiral via a reservoir and peristaltic pump in a loop. In the energy of activation experiments, the reservoir was surrounded by a (25) Matthews, R. W. Solar Energy, in press.
0 1987 American Chemical Society
Photooxidation of Organic Impurities in Water
The Journal of Physical Chemistry, Vol. 91, No. 12, 1987 3329
L5
LO
-
co
35
0
-s X
J
30
0
U
2
25
*
-1
u
a
23
20 15
0.1
10
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I
,
I
I
0
1
8
12
16
20
ILLUMINATION TIME ( m i d
0
L
8
12
16
ILLUMINATION TIME (min 1
Figure 1. Effect of Ti02 loading on salicylic acid destruction; 500 cm3 of 50 pM salicylic acid, flow rate 120 cm3min-': 0, no TiOz;0 , 2 5 mg TiO,; 0 , 50 mg Ti02,0.1 pM subtracted from points to facilitate plotting; 0 , 75 mg Ti02, 0.2 pM subtracted from points to facilitate plotting; C), 100 mg Ti02. 0.3 pM subtracted from points to facilitate plotting; 0 , 125 mg Ti02, 0.5 pM subtracted from points to facilitate plotting.
thermostatically controlled water bath. Analysis. Samples of 3 cm3 were withdrawn for analysis from the reservoir a t regular intervals during each run and returned to the reservoir after measurement. An exception was made in the case of benzoic acid which required an additional reagent before measurement. In this case, corrections were made for the diminishing volume in the reservoir. All analyses were done with a Perkin-Elmer LS-5 luminescence spectrometer using calibration curves and the excitation and emission wavelengths given in Table I. The method of Armstrong et a1.z6was used for the benzoic acid analysis. Carbon dioxide was determined gas chromatographically on a Hewlett-Packard Model 5750 instrument fitted with Porapak Q columns and thermal conductivity detectors. Helium was the carrier gas. A gas loop was connected in parallel with the liquid loop so that the gas bubbles were drawn through the spiral. The gas was sampled at regular intervals via an automatic sampling valve controlled by a Hewlett-Packard Model 3390 sampling integrator. Calibration was done with measured volumes of carbon dioxide added to the gas loop. Actinometry. Potassium ferrioxalate solution and the analytical method of Hatchard and Parkerz7were used to measure the photon flux. All solutions were prepared by using water from a Millipore Waters-Milli Q water purification system and contained equilibrium concentrations of oxygen.
Results TiOz Layer. The effect of the amount of TiOz on the surface of the glass tube is shown in Figure 1 . A salicylic acid solution (500 cm3), of initially 50 pM, concentration was pumped through the reactor at 120 cm3 m i d . In the absence of Ti02, the decrease in salicylic acid concentration upon illumination was trivial. When a layer of 25 mg of T i 0 2 was present, a marked decrease in concentration occurred with illumination time. Further increase (26) Armstrong, W. A,; Black, B. A,; Grant, D. W. J. Phys. Chem. 1960, 64, 1415. (27) Hatchard, C. G.; Parker, C. A. Proc. R. SOC.A 1956, 235, 518.
Figure 2. Effect of initial salicylic acid concentration on the rate of destruction; 500 cm3 solution, 75 mg TiO,, flow rate 120 cm3 m i d .
in the thickness of the TiO, layer resulted in an increase in the destruction rate but at a diminishing rate of increase. A layer of 75 mg of Ti02 was used in all subsequent experiments. The surface area of the inside of the spiral was 877 cm2, so the layer of TiOl on the surface was approximately 85 pg cm-2. Solute Concentration. The effect of initial solute concentration on the rate of destruction of salicylic acid is shown in Figure 2. The lines drawn through the data points were calculated from the integrated Langmuir expression as explained later in the Discussion section. It can be seen that there is an approximately linear decrease in the logarithm of the salicylic acid concentration with illumination time. The expression
is a good approximation to the data over a wide range of concentrations. Similar approximately linear plots were obtained for the solutes listed in Table I. The slopes of these plots, obtained by the method of least squares, are given in Table 11. Flow Rate. Increasing the flow rate through the reactor caused an increase in the rate of solute disappearance. Results for salicylic acid of initially 50 pM concentration are shown in Figure 3. At the higher flow rates the downward curvature of the lines drawn through the points is more noticeable. Effect of Temperature. The destruction rate of 500 cm3 of initially 50 p M salicylic acid circulaing through the reactor at 120 cm3 min-I was determined at different temperatures. The slopes of the In [salicylic acid] values vs. illumination time plots were determined by the method of least squares and the In k' values vs. the reciprocal of the absolute temperatures are plotted in Figure 4. Carbon Dioxide. The yield of carbon dioxide from 200 cm3 of salicylic acid solution at 25 and 50 pM concentrations is shown in Figure 5. The dashed lines show the yield of C 0 2corresponding to total oxidation of the salicylic acid. Actinometry. The measured rate of ferrous ion formation in 500 cm3 of 0.006 M potassium ferrioxalate circulated through the reactor, without Ti02, was 833 pM min-'. Taking $(Fe2+) to be 1.28,28a concentration change of 65 1 pM m i d is calculated for 500 cm3 of solution in which the photon-induced process has a quantum yield of 1.0. (28) Demas, J. N.; Bowman, W. D.; Zalewski, E. F.; Velapoldi, R. A. J. Phys. Chem. 1981, 85, 2766.
Matthews
3330 The Journal of Physical Chemistry, Vol. 91, No. 12, 1987 TABLE 11: k’ Values from Expression 1 for Different Solutes, Flow Rate 120 cm3 m i d . Temwrature 25 O C ~
solute salicyclic acid
concn, WM 1 2 4
IO 20 30 40 50 100 1 2
phenol
IO 20
2-chlorophenol 4-chlorophenol benzoic acid
50 100 10
IO IO 20 50 1 2 5 10 20 50 100 1 2 5
2-naphthol
naphthalene
IO
20
min
0.1017 f 0.0007 0.1024 f 0.0008 0.0989 f 0.0007 0.0975 f 0.0020 0.0857 f 0.0018 0.0857 f 0.0014 0.0811 f 0.0013 0.0695 f 0.0008 0.0519 f 0.0013 0.116 f 0.001 0.102 f 0.002 0.0966 f 0.0019 0.0742 f 0.0017 0.0562 f 0.0002 0.0408 f 0.0005 0.0843 f 0.0013 0.0793 f 0.0017 0.1002 f 0.0016 0.0934 f 0.0007 0.0704 f 0.0016 0.0930 f 0.0007 0.0848 f 0.0010 0.0889 f 0.0008 0.08 12 f 0.0005 0.0671 f 0.0010 0.0617 f 0.0015 0.040 f 0.0009 0.174 f 0.0015 0.176 f 0.0007 0.165 f 0.0008 0.160 f 0.0005 0.154 f 0.0007
6.81 6.77 7.01 7.1 1 8.08 8 09 8.54 9.97 13.35 5.97 6.79 7.17 9.34 12.3 17.0 8.22 8.73 6.92 7.42 9.84 7.45 8.17 7.80 8.53 10.3 11.2 17.1 3.98 3.94 4.20 4.33 4.50 5.50 5.29 5.54 5.82 6.41 7.61
0.126 f 0.0012
50 1 2 5 10
fluorescein
to 5,
k‘, m i d
0.131 f 0.0013 0.125 i 0.0009 0.119 f 0.0015 0.108 f 0.0018 0.091 f 0.0020
20
I
\o
\ 3.1
I
I
3.2
3.3
I
lOOOlT Figure 4. Arrhenius plot for salicylic acid, 500 cm350 pM,flow rate 120 cm3min-l, 75 mg Ti0,. In k’values were determined by the method of least squares.
1.8
I
_
1.61 1.11
-
-
-
... 4 ~
1 4
0.01 0
I
I
1
I
I
I
1
8
12
16
20
21
ILLUMINATION TIME (min)
t 0
i
c
8
12
16
ILLUMINATION TIME(min)
Figure 3. Effect of flow rate on the rate of destruction of 500 cm3of 50 salicylic acid, 75 mg Ti02: 0, 30 cm3 min-I; 0 , 60 cm3 min-I; 0, 120 cm3 min-’; 0 , 180 cm3 min-I; 0 , 300 cm3 min-I; 0, 420 cm3 min-I.
pM
Discussion
Solute Concentration. All solutes demonstrate a linear relationship between In [solute] and illumination time to a good approximation up to 50% decomposition. The k’ values from expression 1 also show a decrease with increasing solute concentration (Table 11), and at high flow rates the plots of In
Figure 5. Carbon dioxide yield from salicylic acid solution (200 mL) passed through spiral reactor coated with 75 mg of TiO, at 120 mL min-l; 20-W BLB lamp. 0, 25 p M ; 0,50 p M ; dashed lines show yield
corresponding to 100% oxidation to C02. [salicylic acid] vs. illumination time show a noticeable curvature downward. From studies on the effect of solute concentration on the rate of carbon dioxide formation from various solutes in UV-illuminated aqueous suspensions of -Ti02 powder,23it was observed that the rate of C 0 2 formation, R ( C 0 2 ) was given by a form of the Langmuir adsorption isotherm
The Journal of Physical Chemistry, Vol. 91, No. 12, 1987 3331
Photooxidation of Organic Impurities in Water
r-
TABLE IV: k 2 Values for 500 cm3 of 50 pM Salicyclic Acid Recirculated at Different Flow Rates through the Reactor
10
20
30
LO
4.55 & 0.03 6.09 f 0.05 7.38 f 0.1 1 9.85 f 0.05 11.7 f 0.10 11.6 f 0.10 13.0 f 0.10 13.6 i 0.10
c
2
L
0
k,, W M m i d
30 60 120 180 240 300 360 420
t
0
5
flow rate, cm3 m i d
50
Figure 6. Plot of values versus O.5[SIo values for salicylic acid at different initial concentrations. Data from Table 11. TABLE III: k, and k 2 Values for Solutes, Flow Rate 120 cm3 m i d , Temperature 25 OC solute k , , pM-' k2,pM m i d klk2,mi& salicyclic acid 0.014 f 0.001 7.63 f 0.41 0.106 f 0.014 4.61 f 0.28 0.107 f 0.016 phenol 0.023 f 0.002 ~
benzoic acid 2-naphthol naphthalene fluorescein
0.017 f 0.002 0.017 i 0.002 0.0107 f 0.0008 0.0315 f 0.0008
6.71 f 0.56 5.46 f 0.43 16.3 f 01.0 4.20 f 0.12
0.114 f 0.093 f 0.174 f 0.132 f
0.023 0.018 0.023 0.007
where [SI is the solute concentration and k, and k2 are constants related to the adsorption and reaction properties of the solute. The Langmuir expression has been useful in explaining other results from photocatalytic studies in UV-illuminated TiOz suspensions.8-1321-23 The integrated form of expression 2 is t=
[SI0 1 1 -In - + -([SIo klk2
[SI
k2
- [SI)
(3)
where t is the time in minutes for the initial concentration of solute, [SIo,to decrease to [SI. It is seen that, at low initial concentrations of S, the second term in the expression becomes small compared with the first term and under these conditions In
[SI0 [SI
- k,k2t = k't
(4)
0
100
200
300
COO
500
FLOW RATE ( m l min-1) Figure 7. Relationship between k2 and flow rate; 500 cm3 of 50 pM salicylic acid; data from Table IV.
expression 2, the constant kl reflects the proportion of solute molecules which adhere to the surface; k2 reflects the limiting rate of reaction at maximum coverage for the experimental conditions. It is seen from the k, values in Table I11 that maximum adsorption occurs with fluorescein and the least with naphthalene, but the limiting rates, shown by the k, values, occur in the reverse order. Thus, although fluorescein is more strongly adsorbed, the destruction rate is not as great as for naphthalene. It is also seen that the product, klk2, agrees well with the k'values obtained for the various solutes a t low concentrations (Table 11). Effect of Flow Rate. Turning now to the flow rate data, it is assumed that the proportion of solute adsorbed on the photocatalyst to that remaining in the aqueous stream is independent of the flow rate; that is, k, is constant although the flow rate may change. Therefore, the value of 0.0139, previously found, is taken as k l for the data in Figure 3 and the expression describing the data becomes
It is also seen from expression 3 that at to,5when [S]/[SIo = 0.5,
(5) Therefore, if expression 3 is a reasonable approximation to the data, a plot of the to.5 values against the O.5[SIovalues for the various initial concentrations should yield a straight line whose slope is I l k 2 and whose slope/intercept is k,/0.693. This plot is shown in Figure 6 for the salicylic acid data in Table 11. The slope and intercept of the plot yielded values of 0.01 39 f 0.001 1 pM-I for kl and 7.63 f 0.41 pM m i d for k2. The lines drawn through the points in Figure 2 were calculated from expression 3 upon substitution of the above numerical values of k, and k,. Similar treatments of the data from the other solutes yielded the k l and k2 values listed in Table 111. In the derivation of
The method of least squares was used to find k2 and the results are given in Table IV. The lines passing through the data points in Figure 3 were calculated from expression 6 with the values of k2 from Table IV. The values of k2 plotted against the flow rate in Figure 7 show a regime at low flow rates in which the k2 value is increasing approximately linearly with the flow rate and a high flow rate regime in which the k2 value is approaching a limiting value, independent of flow rate. This type of dependence is described by the expression k*PFR k2=-- 1 + P F R . - A
(7)
where k* is the limiting value of kZ, is a proportionality constant,
3332 The Journal of Physical Chemistry, Vol. 91, No. 12, I987
Matthews
and FR is the flow rate in cm3 min-I. Application of the method of least squares to the data in Figure 7 assuming expression 7 gave values of 0.0136 f 0.0027 for p and 14.7 f 2.2 for k*. The line drawn through the points is that calculated by substitution of these values in expression 7. It is noted that 0 = k, within experimental error. This may at first sight seem a surprising result, but may be rationalized in terms of expression 7 being of the same form as the Langmuir expression 2 and the amount of solute transported to the stationary reaction surface of a flow through reactor being proportional to the flow rate. Since k2 is proportional to the flow rate, expression 3 may be rewritten to include the flow rate so that -
I
P
This may be rearranged to
Expression 9 permits the calculation of the destruction curves at any initial solute concentration and flow rate within the experimental conditions, with reasonable precision, in terms of only two constants, k , and k*. Energy of Activation. The energy of activation obtained from the Arrhenius plot in Figure 4 was 11.O f 0.8 kJ mol-'. This value agrees well with the 10 kJ mol-' found by Okamoto et aLzo for the destruction of phenol in aqueous solution containing UV-illuminated Ti02 powder. It also agrees well with that found for OH radical reactions.29 On the basis of an energy of activation of 11.0 kJ mol-', a temperature increase of 54 "C would be required to double the rate constant k'. Quantum Yield. From expression 2 it can be seen that at high solute concentrations, since k,[S]>> 1, the rate of solute destruction, -d[S]/dt, is equal to k2. From expression 7 the limiting value of k2 was found to be 14.7 pM min-I. Therefore, if it is assumed that all the radiation emitted by the NEC blacklight tube is absorbed by the Ti02, the maximum quantum yield for the destruction of salicylic acid a t 25 "C is 0.022 (14.7/651). The low quantum yield reflects electron/hole pair recombination as the dominant energy wasting reaction. It appears that there is scope for considerable improvement here using doped Ti02 or noble metal coatings to facilitate electron/hole separation. Reaction Mechanism. The details of the photooxidation reactions occurring with aromatic molecules adsorbed on the surface of bandgap illuminated TiOz in the presence of aerated water are yet to be elucidated. Hashimoto et a1.I8 have proposed reaction schemes involving both positive holes and hydroxyl radicals reacting directly with the aromatic molecule. Carbon dioxide is a major end product in the Ti02photocatalyzed oxidation of each of the solutes in Table I.6 At least part of the oxidation proceeds via hydroxylation of the aromatic ring'7s20and it has been shown by Takagi et aI.I9 that the hydroxylation mechanism is pH dependent with the oxygen atom in the phenol arising from solvent water at lower pHs. The results in Figure 8 show the formation of salicylic acid in 500 mL of benzoic acid circulated through the reactor at 120 cm3 min-I. I t is expected that t h e m-a n d p-hydroxy isomers would show similar formation curves.17 The maximum salicylic acid concentration formed in 50 p M benzoic acid is 0.28 pM after 12 min illumination. The maximum concentration was not achieved in the 100 pM benzoic acid after 24 min illumination, but the curve appears to be heading for a plateau of 1.1 pM salicylic acid. Since after 12 min illumination the benzoic acid concentration decreased from 50 to 18 pM and less than 1% has been converted to salicylic acid, it is evident that the reactions giving salicylic acid ( a n d the other hydroxy acids) represent a minor reaction pathway. It is also evident from the shape of the curves in Figure 8 that salicylic acid is rapidly destroyed by the (29) Elliott, A. J.; Simsons, A. S . Radiat. Phys. Chem. 1984, 24, 229.
"'r
0.0 I I 0 L 8
I
I
I
I
12
16
20
24
ILLUMINATION TIME (min 1
Figure 8. Salicylic acid formation vs. illumination time; 500 cm3 of benzoic acid, flow rate 120 cm3 min-I, 75 mg TiOz: 0, 100 pM benzoic acid; 0 , 50 pM benzoic acid.
0
8
16
24
ILLUMINATION TIME ( m i n l
Figure 9. Destruction of 500 mL of 50 pM salicylic acid and conversion to carbon dioxide in spiral reactor coated with 75 mg of Ti02. Circulation rate 120 mL min-', 20-W BLB lamp. 0,salicylic analysis from fluorimetry, 0, calculated from data in Figure 5 assuming the time to oxidize 500 mL of salicylic acid is 5/2 times the time to oxidize 200 mL.
indiscriminate photooxidation reactions occurring at the surface. The higher concentration formed in the 100 MM benzoic acid apparently results from the increased protection of salicylic acid in the presence of the higher concentration of benzoic acid.30 From the data for the 50 WMsalicylic acid in Figure 5 , the time for the conversion of the same fraction to C 0 2 in 500 cm3 of (30) Matthews, R. W . J. Chem. SOC.,Chem. Commun. 1983, 177.
J. Phys. Chem. 1987, 91, 3333-3336 solution was calculated and compared with the disappearance of salicylic acid measured fluorometrically. The results plotted in Figure 9 show that at 50% destruction of salicylic acid approximately 41% has been converted to CO,; that is, most of the salicylic acid is cleanly converted to CO,; the intermediate products are also converted to C 0 2 eventually.
3333
stants determined for a number of the solutes. A marked dependence of the destruction rate on the flow rate was observed and an expression derived which enables the production of the destruction curve at any solute concentration and flow rate. Ferrioxalate actinometry indicated a maximum quantum yield of salicylic acid oxidation of 0.022 at 25 OC. The activation energy for the photooxidation of salicylic acid was determined to be 1 1.O f 0.8 kJ mol-I.
ConcIusions The rates of photocatalytic oxidation of a number of aromatic compounds in aqueous solutions over UV-illuminated thin films of Ti02in a flow-through reactor have been determined. The rates in each case obeyed approximately first-order kinetics with the rate constants decreasing with increasing solute concentration. The data were more rigorously interpreted in terms of the integrated form of the Langmuir expression and the k , and k2 con-
Acknowledgment. I thank Dr. Ian Hoare for assistance with the computer calculations. Registry No. C02, 124-38-9; Ti02, 13463-67-7;salicylic acid, 69-72-7; phenol, 108-95-2; 2-chlorophenol, 95-57-8; 4-chlorophenol, 106-48-9; benzoic acid, 65-85-0; 2-naphthol, 135-19-3; naphthalene, 91-20-3; fluorescein, 151-67-7; water, 7732-18-5.
Effects of Pressure on Retention in Supercritical Fluid Chromatography Clement R. Yonker,* Robert W. Gale, and Richard D. Smith Chemical Methods and Separations Group, Chemical Sciences Department, Pacific Northwest Laboratory,t Richland. Washington 99352 (Received: July 2, 1986; In Final Form: February 9, 1987)
A simple thermodynamic relationship is described relating solute retention, pressure, and density for supercritical fluid chromatography (SFC). Solute retention in SFC has been predicted based upon the partial molar volume of the solute in the mobile phase and the polymeric stationary phase at infinite dilution. Calculated chromatographic retention for naphthalene and biphenyl has been compared to experimental data for capillary SFC. Predictions of the thermodynamic model accurately fit the retention data, leading to further insight into the retention mechanism in SFC.
Theory Solution retention in SFC as a function of pressure' is given by
Introduction Retention in supercritical fluid chromatography (SFC) is a complex function of temperature, pressure, density, and solute concentration. Van Wasen and Schneider have described the effects of pressure on solute retention, (a In k ' / a P ) , for SFC,'.* while Yonker et al.) have derived a simple thermodynamic relationship between temperature and solute retention for the temperature and pressure regime appropriate to SFC. These results have shown solute retention to be primarily dependent on the partial molar volumes of the solute in the mobile and stationary phases at infinite dilution, the isothermal compressibility, volume expansivity, and heat capacity of the fluid. Yonker et aL3 have applied their approach to calculate solute retention as a function of temperature and shown reasonable agreement with experimental data. The prediction of SFC retention requires knowledge of the above parameters having an impact on solute retention. At constant temperature the isothermal compressibility of a fluid and the partial molar volume of the solute in the mobile phase can be estimated from a two-parameter, cubic equation of state, such as the Peng-Robinson equation of state (EOS).4 This EOS gives more accurate results near the critical point$.5 which is the basis for choosing it in this work. In this article we explore the contributions of the solute partial molar volume in the mobile and stationary phases at infinite dilution to retention as a function of pressure. The Peng-Robinson EOS was used to calculate the partial molar volume of the solute in the mobile phase at infinite dilution and the isothermal compressibility of the supercritical fluid. Estimation of the partial molar volume of the solute in the stationary phase at infinite dilution permits the calculation of solute retention for comparison to experimental data. Through this approach, insight into the dependence of solute retention upon pressure or density in SFC can be obtained.
(a In k ' / d P ) T =
I
(1)
where k'is the dimensionless retention factor of the solute, R is m the solute the gas constant, Tis temperature, r l m p ' " and ~ l s P ~are partial molar volumes in the mobile and stationary phases at infinte dilution, respectively, and K is the isothermal compressibility of the fluid solution. The solute partial molar volume in the fluid phase can be described by
where Vis the volume of solution and n, is the number of moles of solute. For the solutes of interest (naphthalene and biphenyl) the mole fraction in the fluid phase is very low ( x l m P
