Environ. Sci. Technol. 1989, 23, 209-213

Oxidation of H,S with H202 in Natural Waters

Frank J. Millero, ' Arthur LeFerriere,+ Marina Fernander, Scott Hubinger, and J. Peter Hershey

University of Miami, Rosenstiei School of Marine and Atmospheric Science, 4600 Rickenbacker Causeway, Miami, Florida 33149

H The oxidation of H2S with H202 has been studied as a function of pH (2-13), temperature (5-45 "C), and ionic strength (0-6 m) in seawater and NaCl solutions. The rate constant, k (min-' M-I), for the oxidation d[H2SIT/dt = -k[H2SIT[H2O2] a t pH = 8 was found to be given by log k = 8.60 - 2052/T + 0.084Z1/2 in seawater and NaCl so- lutions ( a = 0.07). The energy of activation (AE' = 39 f 2 kJ mol-') for the reaction was independent of ionic strength. The rate increased from pH = 2 to 8 due to the increased rate of oxidation of HS- compared to H2S. The rate constant for the oxidation of H$ was found to be zero at every temperature. The rate constant for the oxidation of HS- was found to be 12.0 f 0.5, 36.2 f 0.4, and 211 f 5 min-' M-', respectively, at 5, 25, and 45 "C. The energy of activation (al' = 51 f 3 kJ) for HS- was higher than the AE' for the overall reaction due to the effect of tem- perature on K1, the dissociation constant of H2S. Above pH = 8.0, the rate slowly decreased in a near linear man- ner. The formation of H02-, which does not react easily with HS-, could be part of the cause of this decrease.

Introduction The oxidation of hydrogen sulfide, H2S, with hydrogen

peroxide, H202, has been studied by a number of workers (1-5). The most extensive measurements in aqueous so- lutions were made by Hoffmann (5). His interest in the oxidation of H2S by H202 was related to the use of H202 to eliminate odor in sewage waters and other anaerobic environmental systems. The elimination of H2S can also alleviate corrosion in sewer lines. These earlier studies (5) were confined to low ionic strength solutions at 25 "C over a pH range of 2-8.

We became interested in the oxidation of H2S by H202 due to the formation of H202 (6-9) in various natural waters that may contain H2S. Although the oxidation of H$ by H202 in surface waters may not be competitive with O2 oxidation, it may be important in rainwater and marine aerosols.

In the present paper, we present measurements of the rates of oxidation of H2S by H202 as a function of pH (2-13), temperature (5-45 "C), and ionic strength (0-6 m) in seawater and NaCl solutions. These results should provide reliable kinetic data valid for a wide range of natural conditions.

Experimental Section Reaction Vessel. The reactions at pH I 8 were carried

out in a 500-cm3 jacketed glass beaker under a nitrogen blanket. The temperature was controlled to k0.05 "C with a Forma bath. Solution samples were taken from the closed vessel with a Rainin automatic pipet. For the measurements a t pH <8 a closed 3.5-L plexiglass vessel with no head space was used (IO), preventing the loss of H2S. The reaction vessel was also controlled with a Forma constant-temperature bath to h0.05 "C. Samples were

'Present address: Chemistry Department, R h d e Island College, Providence, RI 02908.

taken from the vessel by using a PVC piston to displace aliquots into a 10-cm3 pipet. The solutions were stirred with a magnetic stirrer. Both vessels were cleaned with 2 M HNOB between runs to minimize the catalytic effect of trace metals.

Chemicals. All chemicals were reagent grade. The NaHS stock solutions (0.1 M) were prepared by dissolving rinsed Na2S.9H20 crystals in degassed water. The sulfide concentration was determined by potentiometric titration with Pb(N03)?, following the emf with a sulfide ion se- lective electrode and a double-junction reference electrode. The solutions were stored under nitrogen before use. The water used in all the experiments was ion-exchanged Millipore Super Q (18 MQ). The NaCl solutions were prepared by weight. The seawater was collected in the Gulf Stream 10 miles off the coast of Miami and filtered through a 0.45-pm filter. The salinity of the seawater was determined with a Guildline Autosal conductance bridge and the Practical Salinity scale. The hydrogen peroxide stock solutions (0.1 M) were prepared from 30% (-9.8 m) Baker solutions of H202. The stock solutions were standardized by iodometric titrations. For the measure- ments from pH = 6.5 to 13.0, a 0.01 M borax (Na2B4O7- 10H20) buffer was used; from pH = 4 to 6,O.Ol M acetate buffers were made with reagent-grade acetic acid, HC1, and NaOH. Solutions at very low and high values of pH were made with concentrated HC1 and NaOH. The pH was experimentally determined with glass and calomel elec- trodes and a Metrohm pH meter. The electrode system was calibrated with National Bureau of Standards (NBS) buffers for water and tris(hydroxymethy1)aminomethane (Tris) buffers for seawater. The pH values for the NBS standards at a given temperature were taken from Cov- ington et al. (11). For seawater the values of pH (on the free proton scale) for Tris in seawater and NaCl solutions were taken from Millero et al. (12, 13). The addition of as much as 0.05 M borax was found to produce the same results as with the 0.01 M buffer within experimental error.

Sulfide Analysis. The rate of reaction was followed by analyzing the total sulfide concentration, [H2SIT, with a modified methylene blue technique (IO, 14). The ab- sorbance of the methylene blue was determined at 670 nm with a Cary 2200 spectrophotometer. The color reagent was prepared by dissolving 1.79 g of N,N-dimethyl-p- phenylenediamine dihydrogen chloride and 3 g of iron(II1) chloride hexahydrate in 6 M HC1 to a final volume of 500 cm3 (14). Typically, a 5.0-cm3 aliquot was withdrawn from the reaction vessel and added to 15 cm3 of water to quench the reaction, followed immediately by the addition of 0.8 cm3 of color reagent and dilution to 25 cm3. This particular sequence was found to yield the most reproducible results compared to other sequences, e.g., adding the aliquot di- rectly to the color reagent. In experiments at pH = 2 or 3, the peroxide concentration was increased to speed up the rate of the otherwise very slow reaction. The color faded on standing, so the methylene blue was spectro- photometrically analyzed within 5-10 min of color de- velopment. This fading was caused by the high peroxide concentration. The methylene blue color developed from a solution of pH = 2 and low peroxide concentration was

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5.0 1.

0.0

-0.5 -x

9 -1.0 m

25pM 50pM

T lWpM A 2WpM

-

-

-

0.00 0.05 0.10 0.15 -I 0.20

Time/hr Flgure 1. Values of In [H2SIT versus time for the oxkiation of H2S with H,O, for different values of [H,S]; (pH = 8.0, in 0.01 M borax at 25 "C).

0.5 I I

-2.0 -1.5 -4.0 !_p" -3.5 -3.0 -2.5 -2.0 -1.5 -1.0

log [H202I0

Flgure 2. Values of log k,' versus log [H,O2]' for the oxidation of H,S with H202 (k,' is the pseudo-firstorder rate constant) (pH = 8, in 0.01 M borax at 25 "C).

stable overnight. In experiments at higher ionic strengths (4-6 m), the methylene blue color was less intense for a given sulfide concentration than for experiments at lower ionic strengths. This appeared to be caused by the com- plexing of Fe(II1) with the large excess of chloride. Beer's law, however, was found to be valid over the entire ionic strength range.

Results and Discussion

pseudo-first-order rate constant (kl') in In our first series of measurements we determined the

d[H,S]~/dt = -k,'[HzSl~ (1)

for the oxidation of H2S in water. The measurements were made at a pH = 8.0 (0.01 M borax) and 25 "C as a function of [H2SITo = 25-200 pM at [H2O2l0 = 5000 pM and as a function of [H2O2l0 = 500 pM-60 mM at [H~SITO = 25 pM where the superscript (0) indicates initial concentration. The pseudo-first-order behavior as a function of [H2S]0 is shown in Figure 1. The value of k,' was found to be 0.13 f 0.02 min-'. This first-order behavior relative to [H2SITo agrees with the earlier findings of Hoffmann (5). The order with respect to H20z was examined by plotting log kl' versus the log [H2O2l0 as shown in Figure 2. The slope is 0.94 f 0.04, which is essentially first order. If the re- action is assumed to be first order, the value of k = kl'/ [H2O2l0 = 30 f 5 min-I M-l is found for the 27 measure- ments made at pH = 8 and 25 "C. These findings are also in reasonable agreement with the results of Hoffmann (5) who found k = 10-87 min-' M-l between pH = 6.8 and 8.1

210 Environ. Sci. Technol., Voi. 23, No. 2, 1989

Table I. Effect of Ionic Strength on the Rate of Oxidation of HzS with HzOzn

5 "C 25 "C 45 "C I log k

2.28 0 1.24 0 1.66 0 0.02 1.32 0.06 1.69' 0.02 2.25 0.07 1.29 0.07 1.73 0.07 2.30 0.36 1.32 0.36 1.70 0.36 2.28

1.32 0.41 1.71 0.41 2.20 0.72b 1.22 0.62 1.75 0.50 2.18 0.73 1.38 0.72' 1.75 0.62 2.19 1.00 1.18 1.00 1.73 0.72' 2.17 1.25 1.32 2.00 1.76 0.73 2.21 2.00 1.28 3.00 1.79 1.00 2.16 3.00 1.38 4.00 1.80 2.00 2.20 4.00 1.38 4.50 1.84 3.00 2.27 5.00 1.46 5.00 1.89 4.00 2.33 6.00 1.47 1.83 5.00 2.37

5.50 1.90 6.00 2.36 6.00 2.04 2.36

I log k I log k

2.04 2.00

[H2S]~O = 25 pM, [ H Z O ~ ] ~ = 500 pM, and pH = 8.0. ' Seawater.

2.04 E

l . o \ . . . . I . . " I . . ' . , I . . . , . . . .

0.0 0.5 1 .o 1.5 2.0 2

JI 5

Figure 3. Values of log k versus I", for the oxidation of H,S with H,02 in seawater and NaCl at 5 (0), 25 (O), and 45 "C (B) (pH = 8.0, in 0.01 M borax) and seawater (+).

and at 25 "C. In all our subsequent discussions we will examine the overall rate constant k for

d[HZS]T/dt = -k[H2SIT[H202] (2)

In our second series of measurements, we examined the effect of ionic strength on the rate of oxidation at a pH = 8 and at 5, 25, and 45 "C. These results are given in Table I and shown plotted versus the square root of ionic strength in Figure 3. Over most of the ionic strength range the values of log k were found to be linear functions of Ill2 independent of whether the measurements were made in seawater or NaC1. The slopes were almost independent of temperature and ranged between 0.04 and 0.12 (average of 0.08 f 0.04 from 5 to 45 "C). The slopes were smaller than the values of 0.44 f 0.06 found in our earlier mea- surements for the oxidation of H2S by O2 (10). Extensive measurements (15) made in seawater, using the emf technique (5) as a function of salinity, give results that agree with this study and also demonstrate that the rate constants are nearly independent of ionic strength.

We also measured k as a function of ionic strength at pH = 3 and 13. The results at pH = 13 gave log k = 1.33 f 0.01 mi& M-l for four measurements between I = 0 and 3 m. At a pH = 3 in dilute solutions below 0.04 M, no ionic

Table 11. Effect of Temperature on the Rate of Oxidation of Has with HpOan

water seawater (0.7 M) NaCl (6 M) t 'C l o g k t 'C log k t ' C log k 5.0 1.24 5.0 1.22 5.0 1.47 6.0 1.15 15.0 1.59 16.0 1.68

17.2 1.52 25.0 1.75 25.0 2.02 25.0 1.66 35.0 2.00 35.0 2.16 31.2 1.92 45.0 2.17 45.0 2.36 37.1 2.16 45.0 2.28

IH,SlTo = 25 uM, IH90910 = 500 uM, and DH = 8.0.

2.5

2.0

1 .5

1 . 0 ' . ' . ' . ' . ' . ' . ' 3.0 3.1 3.2 3.3 3.4 3.5 3.6

1 ooo/T Flgure 4. Values of log k versus l I T ( K ) for the Oxidation of H2S with H202 for seawater (A) and NaCl (0) SOIUtlOn.

strength dependence was found; however, at I = 3.0 m, the rate was 10 times faster than at I = 0. We attribute this increase in rate to the presence of trace metals. All of our runs at pH = 8-13 were made with enough borax to com- plex these trace metals and suppress the catalytic effect. An experiment a t pH = 11 without borax was completed within 5 min compared to 1.5 h with 0.01 M borax. These results support our contention that the effect of ionic strength on the rates of oxidation are independent of pH if the catalytic effects of trace impurities are avoided.

In our next series of measurements we examined the effect of temperature on the rate of oxidation of H2S by H202. These results are given in Table I1 and shown plotted versus the reciprocal of the absolute temperature (1/T) in Figure 4. The energies of activation for seawater and NaCl were hE* = 39 f 2 kJ mol-l and were found to be independent of ionic strength. The energy of activation in HzO2 is lower than the value (Ah!' = 57 f 4 kJ mol-l) found in our earlier work (10) for the oxidation of H2S with 0 2 .

All of our measurements a t pH = 8.0 in seawater and NaCl have been fitted to the equation

(3) log k = 8.60 - 2052/T + 0.084 Ill2 with a u = 0.07 in log k. This equation should be valid for most natural waters from 0 to 50 "C and to I = 6.0 near a pH of 8.0.

The effect of pH on the rate of oxidation of H2S with H202 was determined from pH = 2 to 13 at 5 , 25, and 45 O C . These results are given in Table I11 and shown in Figure 5. Our results a t 25 "C from pH = 5 to 8 are in good agreement with the results of Hoffmann (5) (see Figure 6). At lower values of pH, his results are faster than ours. This may be due to problems with the emf technique he used. For the slower reactions of H2S with 0 2 or H202, the emf technique may yield unreliable results due to problems with the electrode response ( I O ) .

Table 111. Effect of pH on the Rate of Oxidation of H2S with HzOz

5 oc 25 O C 45 o c pH log k pH logk pH log k 4.0 -2.398 3.0 -2.301 3.0 -1.409 4.7 -1.595 4.0 -1.623 4.0 -0.878 6.0 -0.466 5.4 -0.538 4.7 -0.014 7.0 0.624 6.0 0.135 5.9 1.999

0.541 6.5 0.685 6.78 2.023 7.8 1.313 6.84 1.196 6.8 1.928 8.5 1.160 6.9 1.232 7.41 2.247 9.4 1.000 7.37 1.42 8.0 2.324

0.93 7.5 1.442 2.264 10.0 0.706 1.576 9.03 2.059 10.4 0.368 8.0 1.761 10.0 2.012 10.5 0.490 1.576 11.0 1.905 11.0 0.380 1.66 12.0 1.421 11.2 0.317 8.5 1.561 1.493 11.5 0.460 1.556 13.0 1.273 12.1 0.560 9.0 1.492 1.337 12.3 0.464 9.5 1.306 12.5 0.270 1.239

10.0 1.194 10.5 1.176 11.0 1.185 11.5 1.133 12.0 1.088 13.0 0.636

3.0

Y 2*oi 1 .o

0 - -1.0

O . 4

- 5 . 0 1 . I . I I I I I 8

8 10 12 14 0 2 4 6

PH Figure 5. Effect of pH on the log k for the oxidation of H2S with H202 at 5 (0), 25 (0), and 45 O C (A). Analytical expressions for llnear flts are taken from eq 4 and 5.

3.01- I 2.0 1

::::I1 , \ , , , I , I , I , I -3.0

8 10 12 14 0 2 4 6

PH Figure 6. Values for log k versus pH for the oxidation of H2S with H,O$ (0) present work; (0) Hoffmann; (-) fitted curve accounting for a,-; (--) fitted curve accounting for aHB and aHzoZ.

The effect of pH on the oxidation of H2S at various temperatures can be divided into two linear portions: from

Environ. Sci. Technol., Vol. 23, No. 2, 1989 211

120-

100-

N 80- n w I

8 60-

U

'=. 40 - 20 -

6.0

5.0-

c

4.0- C -

3.0-

2.0

K l / [ H + I Flgure 7. Values of k / a H e s versus K,/ [H+] for the oxidation of H,S with H,02 at 25 "C.

pH = 2 to 7.5 and from pH = 7.5 to 13. The increase between 2 and 7.5 has been fitted to (a = 0.18)

(4)

and the decrease between 7.5 and 13 has been fitted to (a = 0.13)

(5) These linear fits of log k as a function of pH are given in Figure 5. The effect of temperature on the results above pH = 7.5 is the same as found for the results in NaCl at pH = 8 (eq 3). The effect of temperature on the results below pH = 8 is different because of the effect of tem- perature on the ionization of H2S that occurs in this pH range.

From a pH = 2 to 8, the rate increases in a near-linear manner with increasing pH. This is related to the ioni- zation of H2S

H2S - H+ + HS- (6)

and indicates that the HS- species is more reactive than H2S (5, IO). It is noteworthy that the log k determined in this study does not appear to level off at low pH values as was found by Hoffmann (5) for the oxidation of HzS by H202. The leveling off a t low pH was also observed for the oxidation of H2S by oxygen ( I O ) . The leveling off at low values of pH can be related to the difference in the rates of oxidation of H2S and HS- according to the equa- tion (5, 10)

log k = 6.38 - 3420/T + 0.902pH

log k = 12.04 - 2641/T - 0.186pH

= kOaHzS + h a H S (7)

where ai represents the fraction of species i and ko and kl are the rates of oxidation of H2S and HS-, respectively. Combining eq 7 with the dissociation constant K1 gives

h/aHzS = k0 + klK1/[Htl (8)

If eq 8 is valid, a plot of k /aHzS vs Kl/[H+] should yield a straight line with intercept equal to ko and a slope equal to k,. A plot of k/aHzs versus Kl/[H+] at 25 "C is shown in Figure 7. The intercept ko = 0 a t all temperatures within the experimental error. The slopes give kl = 12.0 f 0.5, 36.2 f 0.4, and 211 f 5 min-l M-', respectively, at 5, 25, and 45 "C. The values of In kl a t 25 "C have been plotted versus 1/T in Figure 8 and fitted to

(9)

which gives an energy of activation of AEl* = 51 f 3 kJ mol-' for the oxidation of HS-.

In kl = 25.0 - 6306/T

* h

I ' I ' I '

3.1 312 313 3.4 3.5 3.6 7

This is larger than the value given earlier a t a pH = 8 (AE* = 39 f 2 kJ) because it does not contain terms due to the effect of temperature on the dissociation of H2S. The overall rate constant, k, is related to kl by k = aHskl. Thus, the differential of k with respect to T contains terms due to the effect of T on a H S and k,. The effect of tem- perature on aHS is related to the AH" for the ionization

The simplest mechanism suggested by the pH depen- dence between 2 and 8, therefore, would involve HS- and Hz02 in the rate-determining step with the HS- concen- tration dependent on the pH according to the following equations:

HS- + HzOz - products (10)

rate = -k1[HS-][H202] (11)

of HZS.

By substituting the expression for [HS-] the following rate equation can be derived:

rate = -klKl[H,SlT[H2O2I/(Kl + [H+I) (12) where the second-order rate constant should be equal to klKl/(Kl + [H+]). Using a value for kl = 36 min M-l and pKl = 6.98 at 25 OC (6) , Figure 6 illustrates the pH de- pendence based on this simple mechanism. This mecha- nism predicts that when [H+] >> K,, that is, a t the lower values of pH, the slope of log K vs pH should be a straight line with a slope of 1. The experimental values of log k are within experimental error of this line at all three tem- peratures.

The experimental points are lower than the calculated curve at pH values above 8.0 and fall off in a near-linear manner. A t pH I 11 the [H2Oz] is reduced due to its ionization:

H2Oz = H+ + HOz- (13)

where pKHaO = 11.6 (16). The formation of HOz- causes the rates to decrease above a pH - 8.0 apparently due to the slow reaction of HS- with HOz-.

If we assume that the reaction between HS- and HOz- is small, the decrease in k can be attributed to

k = hi ~ H S ~ H ~ O ~ (14)

aHzOz = [H+1 /([H+1 + KH202) (15)

The 25 "C results (dotted line in Figure 6) show that the addition of the correction for Hz02 ionization does improve the fit above pH = 10, but does not explain the nearly

where

212 Environ. Sci. Technol., Vol. 23, No. 2, 1989

Environ. Sci. Technol. 1909, 23, 213-218

2.0-

1.0-

Y

0 IT 0.0- -

-1.0-

-2.0-

0

1 - 6 h lb 12 14

-3.0 1 . I 3

0 2 4

PH Flgure 9. Values for log k versus pH for the oxidation of H2S with H20,: (A) 45 OC; (0) 5 OC; (-) fitted curve accounting for aHS- and

linear offset above pH = 8. The 5 and 45 "C results shown in Figure 9 look slightly better. Obviously, this simple explanation does not completely explain the near-linear dependence above pH = 8. Other factors such as the formation of polysulfide ions (HS-,) and S2- may be im- portant.

Over the entire range of pH and temperature studied there was little or no dependence of the rate on ionic strength, that is, no salt effect. This result is consistent with the simple mechanism in that the slow step does not involve two ions. With a t least one neutral molecule in the slow step, transition-state theory predicts that the rate of the reactions will be independent of ionic strength.

To propose more detailed steps beyond those in the simple mechanism described above (eq 14) would be speculative. It is likely that the reaction involves several elementary steps involving one-electron transfers (5), but it is not known at this time what the detailed mechanism is. Studies need to be performed to identify possible in- termediates, e.g., elemental sulfur, S2032- or other poly- thionates, Sot- and S042-. Free radicals are always likely intermediates in redox reactions and may be identified

(0s 15).

with appropriate scavengers or by using ESR (electron spin resonance).

Millero et al. (10) found that the half-life for the oxi- dation of sulfide in seawater with oxygen was 30 h at 25 "C. If one uses the [H202] = 1.0 X lo-' M found (7-9) for surface seawater, then the half-life for the oxidation of sulfide by peroxide in seawater would be 2800 h.

M, the oxidation of H202 with H2S becomes competitive with 02. Such con- centrations of H202 are found in rainwaters (17); thus, peroxide oxidation of H2S may be more important than oxygen oxidation in aerosols or rainwaters.

Registry No. HzOz, 7722-84-1; HzS, 7783-06-4; SH, 15035-72-0.

At concentrations of H202 >

Literature Cited Classen, A,; Bauer, 0. Ber. Dtsch. Chem. Ges. 1883,16,1061. Wasserman, A. Justus Liebigs Ann. Chem. 1933,503,249. Feher, F.; Heuer, E. Angew. Chem. A 1949,59, 237. Satterfield, C. N.; Reid, R. C.; Briggs, D. R. J . Am. Chem. Soc. 1954, 76, 3922. Hoffmann, M. R. Enuiron. Sei. Technol. 1977, 11, 61. Millero, F. J. Mar. Chem. 1986, 18, 121. Moffett, J. M.; Zika, R. G. Mar. Chem. 1983,13, 239. Zika, R. G.; Saltzman, E.; Cooper, W. J. Mar. Chem. 1985, 17, 265. Zika, R. G.; Moffett, J. M.; Copper, W. J.; Petasne, R. G.; Saltzman, E. Geochim. Cosmochim. Acta 1985,49, 1173. Millero, F. J.; Hubinger, S.; Fernandez, M.; Garnett, S. Environ. Sei. Technol. 1987, 21, 439. Covington, A. K.; Bates, R. G.; Durst, R. A. Pure Appl. Chem. 1985,57,531. Millero, F. J. Limnol. Oceanogr. 1986, 31, 839. Millero, F. J.; Hershey, J. P.; Fernandez, M. Geochim. Cosmochim. Acta 1987,51, 707. Cline, J. D. Lirnnol. Oceanogr. 1969, 14,454. Millero, F. J.; Garnett, S., unpublished results, 1984. Gray, R. D. J . Am. Chem. Soc. 1969,91,56. Kok, G. L. Atmos. Environ. 1980,14, 653.

Received for review March 29, 1988. Accepted September 15, 1988. We wish to acknowledge the support of the Office of Naval Research (Grant N00014-87-G-0116) and the Oceanographic Section (Grant OCE-8600284) of the National Science Foun- dation for this study.

An Assessment of the Importance of Direct Solar Degradation of Some Simple Chlorinated Benzenes and Biphenyls in the Vapor Phase

Nigel J. Bunce,* James P. Landers, Jo-Anne Langshaw, and Jamie S. Nakal

Department of Chemistry and Biochemistry, University of Guelph, Gueiph, Ontario, Canada N1G 2W1

rn Quantum yields of decomposition of some representative chlorobenzenes and chlorobiphenyls in the vapor phase are of comparable magnitude with those previously observed in hydrocarbon solvents. Combination of quantum yield data with tabulations of solar intensities indicates that PCBs should degrade in sunlight with a half-life of several days, but that chlorinated benzenes will be much less photolabile. This is a consequence of poor spectral overlap between the solar spectrum and the absorption spectra of these substances, rather than of intrinsically low quantum yields for decomposition.

1 . Introduct ion This study has been carried out in an attempt to esti-

mate the rates of direct photodecomposition of vaporized chlorinated benzenes and chlorinated biphenyls (PCBs) in the atmosphere upon absorption of sunlight. Most previous studies of the photochemistry of these chlorinated pollutants have been done in solution, usually in organic solvents (1).

The literature on the vapor-phase photochemistry of these compounds is very sparse. Ichimura and Mori (2) photolyzed chlorobenzene vapor in the presence of ethane; they obtained benzene, hydrogen chloride, and butane as the products. These were rationalized as forming by the sequence of eq 1-4, analogous to the photochemistry ob- served in alkane solutions.

P h C l L Ph' + C1' (1)

0013-936X/89/0923-0213$01.50/0 0 1989 American Chemical Society Environ. Sci. Technol., Vol. 23, No. 2, 1989 213