Environ. Sci. Techno/. 1995, 29, 1215-1222
Photoreduction of Iron Oxyhydmxides and the Photooxidation of Halogenated Acetic Acids S I M O 0 . PEHKONEN,' RONALD L. SIEFERT, A N D M I C H A E L R. H O F F M A N N * Environmental Engineering Science, W. M.Keck Laboratories, California Institute of Technology, Pasadena, California 91 125
The photolytic reduction of ferrihydrite (am-Fe203*3H20), lepidocrocite (y-FeOOH), goethite (a-FeOOH), hematite (a-Fe203), maghemite (y-Fe203) and iron-containing aerosol particles (Feaerosol)in the presence of a series of halogenated acetic acids has been investigated. The fastest rates of photoreduction of Fe(lll) to Fe(ll) were achieved with ferrihydrite as an electron acceptor and fluoroacetic acid as an electron donor. The relative rates of photooxidation of the monohalogenated acetic acids with ferrihydrite in order of decreasing reactivity were as follows: FCH2C02H > CICH2C02H > BrCH2C02H > ICH2C02H; for multiple substituents the relative order of reactivity was as follows: FCH2C02H > F2CHC02H > F3CC02H. With respectto the iron oxide electron acceptors, the relative order of reactivity toward monohaloacetate oxidation was am-Fe203-3H20 > y-Fe203 > y-FeOOH 1 a-Fe,Os L Feaerosol > a-FeOOH. Strong kinetic isotope effects observed for the photooxidation of CICD2C02H suggestthat the oxidation of the mono- and disubstituted haloacetic acids proceeds via hydrogen-atom abstraction by surface-bound hydroxyl radicals to produce haloacetate radicals, which in turn yield the corresponding halide and glycolic acid. Fully halogenated haloacetic acids appear to be oxidized via a photo-Kolbe mechanism to yield the corresponding halo acids and C02.
Introduction Iron is one of the most abundant elements in the earth's crust, where it is present both as Fe(I1)andFe(II1)(I).Crustal iron is readily transferred into the atmosphere bywind and to a lesser extent by volcanic activity. When entrained in the atmosphere, iron {as Fe(I1) and Fe(II1)) plays an important role in the multiphasic atmospheric chemistry ofS(IV)(2-9)andavarietyoforganiccompounds (10-14). Iron(II1)-containingaerosols, which may serve as cloud and fog condensation nuclei, are expected to participate in electron-transfer reactions within atmospheric water droplets (15). TheFe(II)/Fe(III)redoxsystemalso hasachemical and a catalytic effect on other chemical species of interest in natural water systems. For example, iron is an important component in the redox chemistryof chromium in estuarine water (16) and in the redox chemistry of copper in atmospheric water (I 7). Halogenated acetic acids are predicted to be one of the major degradation products of hydrochlorofluorocarbons (HCFCs) and hydrofluorocarbons (HFCs) in the troposphere. The HCFC and HFC families of compounds are being considered as substitutes for chlorofluorocarbons (CFCs) (18). Degradation in the troposphere is initiated by hydroxyl radical attack and concluded with the formation of halogenated acetic acids and the corresponding hydrohalic acid. Many halogenated acetic acids are toxic to organisms and plants (19-21) and may present a threat to the biosphere if their concentrations were to increase from present levels to concentrations in the range of 5-50 nM. Schwarzenbach (22)has detected chlorinated acetic acids in rainwater in Switzerland. Since several chloroacetic acids are used in the pharmaceutical industry, they are detectable at trace levels in some surface waters (23). Chlorinated and brominated acetic acids are also present in many municipalwater supplies as byproducts of chlorination (2426). Laboratory research to date has indicated that the diand trihaloacetic acids in the aqueous phase are resistant to chemical degradation (27)and are expected to be stable in the troposphere because their UV absorbance falls off rapidly at it 250 nm, and therefore their direct photolytic decomposition is negligible. Hu et al. (28)have determined that gas-to-water exchange of halogenated acetic acids is limited only by gas-phasediffusionto water droplet surfaces. In this paper, we present results of experiments that probe the coupled processes of photoreduction of iron oxide polymorphs and photooxidation of halogenated acetic acids. Our principal experimental objectives were to determine the kinetics and mechanisms of photoreduction {Fe(III) Fe(I1)) in a series of iron oxide polymorphs as a function of the physicochemicalproperties of the electron donor and wavelength. Results of experiments with synthetic iron oxide polymorphs are compared to the photoreduction of ambient aerosol particles in the presence of halogenated acetic acids.
-
* To whom correspondence should be addressed: e-mail address:
[email protected]; FAX: 818-395-3170. t Present address: Department of Civil and Environmental Engineering, University of Cincinnati, Cincinnati, OH 45221-0071.
0013-936)(/95/0929-1215$09.00/0 0 1995 American Chemical Society
VOL. 29, NO. 5, 1995 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
1215
60
TABLE 1
,
I
Characteristics of Hdogei~atelAcetic Acids haloacetic acids fluoroacetic chloroacetic brornoacetic iodoacetic dichloroacetic trichloroacetic difl uoroacetic trifluoroacetic
p&,
a
2.59 2.86 2.9 3.18 1.26 0.64 1.22 0.23
db
log(BFalll)-HA)C
1.10 0.94 0.908 0.72 1.94 2.65 2.05 2.61
2.2 2.1 2.1 2.1 1.9 0.85d
a Refs 41 and 42. Refs 41 and 42; u* are values for the substituents attached to COOH. Values at 2 0 "C and at I=1.0 (33, 43). I = 3.0 M NaC104.
Experimental Section General experimental approaches, instrumentation, and procedures have been described previously by Pehkonen et al. (14) and Siefert et al. (13). The halogenated acetic acids used in this study are listed in Table 1; they have pK, values ranging from 0.2 to 3.0. The photoreduction of ferrihydrite in the presence of fluoroacetate was carried out in the presence of 3 atom % HPO. After 2 h of irradiation, an aliquot was withdrawn, filtered, dried, and redissolved in ethyl ether. After isolation in ethyl ether, the glycolic acid was derivatized by diazomethane to yield the methyl ester of 2-methoxyacetic acid. The 180-labeled methyl ester of 2-methoxyacetic acid was then quantified by GUMS (HP GC 5880 Series 11, MS Engine 5989A). The concentration of Br- was measured with an ionselective electrode (Orion Model 94-35 halide electrode). For anoxic control reactions, water in the irradiation vessel was purged with NZfor 60 min prior to the experiment, and the vessel was kept closed during the irradiation. Experiments with deuterated chloroacetic acid (Cambridge Isotope Laboratories) were performed similarly to the nonlabeled chloroacetic acid (i.e., 3 mM chloroacetic acid, ferrihydrite, lepidocrocite, maghemite, or hematite suspensions {resulting in 130 pM Fetot}at pH -4.25). The incident photon flux of the light source was measured using Aberchrome 540 { (E)-a-(2,5-dimethyl-3furylethylidene)(isopropylidene)succinic anhydride} (29) and by digital irradiance measurements. The light intensity of the 450-W Xenon light source is similar to the light intensity measured on a cloudy summer morning in Los Angeles between 8 AM and 11 AM from 380 to 750 nm (-15-40 x 1015 (quanta s-l) cm-2). In most cases, a minimum of three replicate experiments was performed. The Fe(I1) concentrations during the replicate experiments were averaged, and the relative standard deviation (100%xstandard deviatiodthe average) was used as the magnitude of the error bars in the graphs. Initial rates of photoreduction of Fe(II1) were determined by a linear least-squares fit when the data points formed a straight line from the beginning of the experiment to the end and by a linear least-squares fit of the first three data points in cases where linearity did not persist.
Results Relative Hfects of Electron Donor on Obeerved Photoreduction Rates. Figure l a shows the trends of photochemical reduction of ferrihydrite. Other iron oxyhydroxidesshow similar behavior and therefore are not plotted; the rates for arn-FezO3.3H20, y-FeOOH, a-FeOOH, a-Fez03,and y-Fe2O3 1216
ENViRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 5,1995
0
500
1000
1500
t (min)
FIGURE 1. Photoreductive dissolution of ferrihydrite with a variety of electron donors. [Fe(lll)lK,*.r = 13OpM, [electron donor]iai,l= 6.0 mM, Aimd from 320 to 800 nm, and pH 425. (a) Fluoroacetate, chloroacetate, bromoacetate, iodoacetate, acetate, and no organic electron donor. (b) Trichloroacetate, dichloroacetate, monochloroacetate, and acetate.
-
are summarized in Table 2. As can be seen in Figure la, there are several obvious trends in the photoreduction of iron oxides in the presence of the halogenated acetic acids. Fluoroacetate reduced Fe(II1) at the fastest rate followed by chloroacetate and bromoacetate. In addition, after a rapid initial rate of Fe(II1) photoreduction in the presence of fluoroacetate,the rate of Fe(I1D photoreduction decreased significantly. The rate of Fe(II1) reduction in the presence of trichloroacetate and dichloroacetate was substantially lower than in the presence of chloroacetate (Figure lb). Kinetic Isotope Effects. Insight into the mechanism of electron transfer was provided by photoreduction experiments with deuterated chloroacetic acid, CDZClCOOH, that were carried out with lepidocrocite, hematite, maghemite, and ferrihydrite. The deuterium kinetic isotope effect for the observed rate of Fe(II1) photoreduction was found to be 5.5 for ferrihydrite and maghemite and 2 for lepidocrocite and hematite. Experiments with 180-labeledwater were carried out to determine the pathway leading to one of the photooxidation products, glycolate. No 180-labeled glycolate was detected. Role of Oxygen. In order to determine the role of oxygen in the reaction mechanism, the determination of the rate of ferrihydrite photoreduction and the rate of Br- photoproduction were carried out both with deoxygenatedwater and with water in equilibrium with air. A significantly reduced rate of iron-induced Br- photoproduction was
TABLE 2
Initial Rates (nM m i d ) of Photoreduction of Iron Oxyhydroxides with Halogenated Acetic Acids goethite
hematite
0.2 0.13
fluoroacetate chloroacetate bromoacetate iodoacetate dichloroacetate trichloroacetate difluoroacetate trifluoroacetate no added donora
lepidocrocite
ferrihydrite
maghemite
FeasroIDl
18 1.3 0.8 0 0.7 0.5 0.7
70 35 27 4.2 4.4 3.1 5.5
39 2.2 1.0 0 1.2 0.55
12
0.2
2.2
14 2.3 1.3 0.8 0.6
4.8 0.02
0.01
10
1.3 0
Proton-assisted iron oxide dissolution. 101
.
I
.
0
0
200
400
800
800
lo00
340
320
t (mln)
FlGURE2. Photoproduction of&- withferrihydrite and bromoacetate as the electron donor in deoxygenated and air-equilibrated water and photoproduction of Br- without iron. [Fe(lll)]infiia~= 130 pM, [bromoacetateli.~,~= 6.0 mM, [bromidelinitiai= 0, kinadfrom 320 to 800 nm, and pH 4.25.
-
observed in a deoxygenated aqueous solution compared to a solution in equilibrium with air (Le.,decrease from 54 to 17 nM/min in the Br- production rate) (Figure 2). However, no further increase in the rate of photoreduction was observed as the oxygen concentration was increase from 0.25 to 1.2mM. Bromoacetate was shown to undergo photolysis in the absence of metal oxide polymorphs; however, the rates were found to be faster in the presence of the iron oxides as shown in Figure 2. Effect of the Wavelength. The wavelength dependency of the photoreduction of ferrihydrite with trifluoroacetate was studied from 325 to 400 nm with a bandwidth of 6 nm. The results are presented in Figure 3. The rate of Fe(II1) photoreduction decreases by a factor of 3 from 325 to 400 nm. The trend of decreasing rate of photoreduction from 325 to 400 nm is consistent with the onset of a bulk-phase 02-to Fe(II1) charge-transfer transition common to all of the iron oxide polymorphs (30, 31). Similar trends have been observed in previous studies of iron oxides (14, 32). The calculated quantum yield at 325 nm using the standard equation described elsewhere (30) is 4 x this value is apparent because the absorbance used in the calculation includes light scattering. Effect of the Iron Phase. The relative reactivities of the iron oxide polymorphs with respect to the haloacetic acids are summarized in Table 2. The initial photoreduction rates in the presence of fluoroacetate at pH = 4.25 for ferrihydrite, lepidocrocite, hematite, goethite, and maghemite were measured, and the following order of reactivity was found to be ferrihydrite > y-FezOs > y-FeOOH 2 a-FezOs >
360
380
4m
420
Lambda (nm)
FIGURE 3. Photoreductive dissolution of ferrihydrite with trifluoroacetate as a function of wavelength. [Fe(lll)]inki,l= 130pM.[electron donor]i,ir.l = 6.0 mM, bandwidth = 6.5 nm, and pH 4.25.
-
0
200
400
800
800
lo00
1200
1400
1800
t (mln)
FIGURE 4. Photoreductive dissolution of ambient iron-containing aerosols in the presence of fluoroacetate, chloroacetate, formate, and no donor. [Fe(lll)]infiiai 13pM, [electron donorlinitiai = 100pM, Aimd from 320 to 390 nm, and pH 4.25.
- -
a-FeOOH. The observed order of reactivity is similar to that observed for formate as an electron donor (14). Natural Iron-ContainingAerosols. Preliminary results indicate that the photochemical behavior of natural ironcontaining aerosol suspensions (Fe content -4%) in the presence of halogenated acetic acids leads to a significant rate and extent of Fe(II1) reduction and photooxidation of the haloacetic acids. This is illustrated by the experiments of ambient iron-containing aerosols with fluoroacetate and chloroacetate as shown in Figure 4. In these experiments, the initial rates of photoreduction were found to be similar to those for formate as an electron donor (14). Control VOL. 29. NO. 5. 1995 /ENVIRONMENTAL SCIENCE & TECHNOLOGY
1
1217
” ” ’
2.0
-1.5
i,
+L . p l d o C r 0 t l t .
1
L 0.7
0.8
0.a
1
1 .I
1.2
@
FIGURE 5. LFER relationship between Iog[kpholo,.dudon]and cr+ for the photoreduction of ferrihydrite, lepidocrocite, and maghemite with fluoro-, chloro-, bromo-, and iodoacetate.
experiments were carried out to test the possibility that organic compounds originating from the ambient aerosol were contributing to Fe(II1) reduction in addition to fluoroacetateand chloroacetate. These experiments showed that the initial rate of Fe(II1) photoreduction was smaller, when no additional electron donor was introduced to the reaction vessel (pH at 4.2 maintained by perchloric acid). Effect of Ambient Competing Ligands. The photoreduction rates of ferrihydrite (at 40 pM total Fe) were examined in the presence of 200 pM acetate [known to photoreduce iron oxides at a much lower rate (14)compared to chloroacetate] and 100 pM chloroacetate at pH 4.25. Acetate and ferrihydrite were allowed to equilibrate for 60 min in the stirred reaction vessel prior to the addition of chloroacetate and the start of the irradiation. No apparent change in the rate of ferrihydrite photoreduction was observed compared to the case with ferrihydrite and chloroacetate as the sole electron donor.
-
Discussion The difference in the rate of Fe(II1) photoreduction between chloroacetate and bromoacetate can be attributed to the greater electronegativityof chlorine than bromine. Due to its greater electronegativity, chlorine will withdraw more electron density away from the adjacent carbon atom than will bromine or iodine, thus stabilizing a radical formed from an H-atom abstraction by a suficial ‘OH radical. Stabilization of a free radical intermediate is suggested by the LFER of Figure 5 in which lOg(kphotoreduction)is plotted against the Taft parameter, 8,for ferrihydrite, maghemite, and lepidocrocite. The relative stability of the surface complexes of chloroacetate and bromoacetate with >FeOH surface sites cannot explain the observed differences in photoreduction rates {i.e., j31: 1 = 126 for Fe(02CCH2C1)2f and PI: 1 = 120 for Fe(02CCH2Br)2+(33)) In addition, the similarity of pK,’s of chloroacetic and bromoacetic acid (2.86 and 2.90) indicates that both are dissociated approximately to the same extent at a given pH. The continuous decrease in the rate of ferrihydrite photoreduction in the presence of fluoroacetate (Figure la) can be explained by the apparent buildup of F- in the reaction vessel (F- is an expected product based on the production of Br- in the case of bromoacetate and the concurrent production of glycolate as observed by GCI MS). Since F- forms strong complexes with Fe(II1) (33),it 1218 ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 5,1995
is likely to bind strongly at reactive surface sites and thus slow down the photoreduction of matrix-bound iron. This effect is not seen during the photooxidation of chloroacetate and bromoacetate with the corresponding production of C1- and Br-. Chloride and Br- are known to form very weak complexes with Fe(II1) (331, and thus the rate of Fe(II1) photoreduction remains relatively constant throughout an experiment. Similar trends were observed for the other iron oxyhydroxides. A value of kH‘lkD‘ = 5.5 for the kinetic isotope effect in the case of ferrihydrite and maghemite indicates that the transition state has a higher degree of bond breaking (Le., the hydrogen atom is bound about equally to two interacting species in the transition state (the surface OH radical and the carbon of the chloroacetic acid}),while the lower value of 2 indicates that the transition atate in the case of lepidocrocite and hematite does not have a high degree of bond breaking and is either product-like or reactant-like. These results are consistent with the fact that the hydrogenatom abstraction is the rate-determining step, and they explain the lower reactivity of dichloroacetate since in the case of dichloroacetate there is only one hydrogen atom available for abstraction. Since no 180-labeledglycolate was detected, the attack on the acetate radical by “vicinalwater” and not bulk H2I80 water is suggested. “Vicinal water” is the ordered water layer immediately adjacent to solid surfaces (34). Since ferrihydrite is a very porous iron oxide, the amount of the “vicinal water” i s relatively high compared to other iron oxides (34). We can conclude from our results that the rate of the attack of the acetate radical by a water molecule is faster than the rate of exchange ofwater molecules between vicinal and bulk water. The experiments with deoxygenated water indicate that dissolved 0 2 does not effect the rate of Fe(II1) photoreduction but plays a role in the photoproduction of Br(Figure 2). In addition, the rate of Fe(I1)production is very similar (within experimental error) to the net rate of production of Br-. This indicates that for every Fez+ produced there is one Br- produced; therefore, the product of bromoacetate photodegradation, glycolate, does not appear to participate in the photoreduction of Fe(II1).This is most likely due to the fact that bromoacetate is present at a much higher concentration than glycolate (6 mM vs 5-20 pM) and because glycolate and bromoacetate have similar constants for Fe(H20)s3+complexation {logpl. = 2.9 vs log PI: = 2.1 for bromoacetate} (33). On the basis of the experimental observations, we propose a general mechanism (Figure 6) for the photooxidation of surface-bound electron donors and the photoreduction of surficial Fe(II1). Based on this general reaction mechanism (35,361,the following stepwise mechanism can be written: >Fe(III)OH + BrCH,CO,-
2>Fe(III)OH
+
+ H’
-
+
>Fe(III)O,CCH,Br H2O (1)
+
>Fe(III)OH’+ >,Fe(II)OH-
(2)
+
>Fe(III)OH’+ >Fe(III)02CCH,Brrds >Fe(III)OH,’ > Fe(III)O,CHBr (3)
0 0
Fe(W Fe(lll)
Oxygen detachment of Fe(ll) from surface
*
+ e&’-
Fe(ll)aq
OH *+
H
I ,Br
H-C
-
i
H HO I H C ‘’
I
I
H I ,Br H
Ht
+ HO,’
FIGURES. Pmposed mechanism forthe photoreduction of imn oxidesandthe concomitantphotooxidation of haloacetic acid (Le.. bmrnoacetate as a specific example).
>Fe(III)O,CHBr
+ H,O - >Fe(III)OH,++ BrCH,CO;
d[Fe(II)I,, --y=
dt
k3Zabs@,[>Fe(III)O,CCH,Brl
k-, + k3[~Fe(III)0,CCH,Brl
(8)
On the basis of this rate expression, we can define the quantum yield forFe(II1)photoreduction as the ratio of the rate of Fe(l1) production to the rate of photon absorption as follows:
60 O+
,C-C-Br
0-
/ 0-O
\H
-laa
OH -
o+ /H ,C-C-OH 0-
+ Br- + HO;
(7)
(9)
Substitution of eq 8 into eq 9 yields
\H
Using a similar steady-state kinetic analysis as outlined by Hoffman et al. (35).we obtain the following rate equation for the rate of production of Fe(I1) in solution:
vllabs
&=
k,k_~’~,[>Fe(III)O,CCH,Brl
1 + ~k_,-’[>Fe(III)0,CCH2Br]
(10)
If we assume that 62 = 1 and K = k3/k+, then eq 10 can VOL. 29.
NO. 5.1995 I ENVIRONMENTAL SCIENCE &TECHNOLOGY m 1219
be simplified as follows:
''
=1
K [ >Fe(III)O,CCH,Brl
+ K[>Fe(III)O,CCH,Brl
0.1
(11)
For a given set of conditions, in which the total reactive surface area is constant and the bulk solution concentration ofXCH2C02- is also constant, such that K[>Fe(III)02CCH2Brl Fe(III)O2CCH2Br1labs. During the initial stages of photoreduction (i.e,, t < 500 min), the rate of production of Fez+is linear, and the slopes of the lines in Figure la,b yield values of K . When deuterated chloroacetic acid, C1CD2CO2H,is substituted as an electron donor, we see that kH'lkD' = 5.5 in the case of ferrihydrite and maghemite and 2 in the case of hematite and lepidocrocite. The general mechanism of eqs 2-7 should also be applicable to ferrihydrite, especiallyin light of recent reports of the semiconducting-like characteristics of ferrihydrite (bernalite) in the perovskite crystal structure (37, 38). The comparison of the rate expression obtained in this study to the expression obtained by Siffert and Sulzberger (32) illustrates the differences between two proposed mechanisms. Based on our study, the rate-determining step appears to be the abstraction of hydrogen atom from the surface-bound haloacetate in contrast to the detachment of surface Fe(I1) from the crystal lattice as proposed by Siffert and Sulzberger (32). Furthermore, valence-band hole transfer to a surface trap is a precursor step to the rate-determining step (rds) in our system. Finally, the chromophore in our system appears to be the bulk iron oxide polymorph and not the Fe(II1)-haloacetate surface complex as proposed by Sulzberger and co-workers (32) and Stumm and co-workers (39). These differencesindicate that the mechanism of iron oxide carboxylate photochemistry is dependent on the chemical nature of the electron donor and the iron oxide and that one mechanism cannot be extended readily to other systems. Experiments with varying concentrations of the iron oxide were carried out to determine whether the observed rate of Fe(II1) photoreduction is directly proportional to the concentration of the surficial {>Fe(III)OH}. The observed results with lepidocrocite, hematite, and ferrihydrite in the presence of chloroacetate as the electron donor indicate that the kinetics involved are first order with respect to the surficial [{Fe(III)OH}],since the rate of Fe(111) photoreduction doubled with the doubling of the concentration of total iron (up to -130 pM) for all three iron oxides (Figure 7a). Above a concentration of 130p M of total Fe, the rate leveled off (most likely due to increased light scattering of the iron oxide suspension). The rate of Fe(II1)photoreduction is plotted against the concentration of chloroacetate in Figure 7b. This dependence follows a Langmuir-Hinshelwood profile. Based on this Langmuirian behavior, the concentration of [>Fe(III)02CCH2Brlwas calculated to be 1.8 x loT5M for a solution concentration of 6 m M bromoacetate in the presence of ferrihydrite with a specific surface area of 250 m2/g. Similar calculations were carried out for the other haloacetates. The depen1220
ENVIRONMENTAL SCIENCE & TECHNOLOGY / VOL. 29, NO. 5, 1995
i
i
-
b
c 0 0
0.2
0.4
0.8
0.8
1
1.2
Rdatlw Ught lntor~alty
FIGURE 7. (a) Dependence of the rate of iron photoreductionon the concentration of iron oxide for hematite, lepidocrocite, and ferrihydrite in the presence of chloroacetate. The concentration of iront, was varied, the concentration of chloroacetate was 3.0 mM, and pH 4.25. (b) Dependence of the rate of femhydrite photoreductionon the concentration of chloroacetate. The concentration of iront, was 130 pM, and pH 4.25. (c) Dependence of the rate of iron photoreduction on the absorbed light intensity in the case of lepidocrocite and chloroacetate as the electron donor. The concentration of iron, was 810 pM, the concentration of chloroacetate was 3.0 mM, and pH 4.25.
-
-
-
dence of the rate of lepidocrocite photoreduction with chloroacetate on the absorbed light intensity is plotted in Figure 7c. This observation is consistent with the linear dependence on absorbed light intensity predicted by eq 8. The trihaloacetates, unlike the mono- and dihaloacetates, appear to be photooxidized by a photo-Kolbe mechanism to yield CF3*radicals, which then react with the O2 to produce CF302.radicals as follows:
TABLE 3
-
Surface Ionization Constants: PK, Is, Specific Surface Area$ (A), Average Particle SizesC (D ), Nominal Band-Gap Energies (E,”), and the Enset of 02Fell’ Charge Transitionsd for the Iron Oxide bolynorphs
a
iron oxide
PL1
PK ~ z
PKW
A (m2/a)
Dp(nm)
hematite goethite lepidocrocite ferrihydrite maghemite
6.7 6.2 6.0 7.3
10.3 9.0 8.0 8.9
8.3 7.5 7.1
30 40 80 250 130
60 1500 1300 370 350
8.0
(eV)
-
L(02- Fe3+) (nm)
2.02 2.10 2.06
380 -340 -340 -285
Refs 44 and 45. Refs 14, 40, and 44. Based on photon correlation spectroscopy measurements ( 7 4 ) . Refs 14 and 31.
>Fe(III)OH
+ CH,CO,- + H+
-
>Fe(III)O,CCF, >Fe(III)OH
+ H,O
(13)
+ >Fe(III)O2CCF3+ H,O ki
>Fe(II)OH>Fe(II)OH-
+ >Fe(III)OH,+ + CO, + CF3 (14)
+ 6H,O - > + Fe(H,O);+ + OHCF3 + 0, - CF30,’
2CF30,’
-
-
CF,OOCF,
+ 0,
+ FCFO + 0, CF,O + H,O - 2HF + CO,
2CF,O,’
CF,OF
(15) (16)
(17) (18) (19)
Evidence for this proposed mechanism includes the observed decrease of the rate of Fe(II1)photoreduction due to the buildup of F-. The apparent order of reactivity of the iron oxides is an indication of the importance in the strength of Fe-0 bond in the different iron phases; there appears to be a weak correlation between photochemical reactivity and the calculated { >FeOH) concentrations (14). Photochemical reactivity (Table2) correlates relatively well with the surface area of each iron oxide (Table 3) if goethite is not included. The reason for the nonlinearitywith goethite is not known. Ferrihydrite has a weakly defined structure (37,38,40),and thus bonding between iron and oxygen in ferrihydrite is weaker than in the other iron oxyhydroxide polymorphs. Our results are consistent with the results of Leland and Bard (31) for the photooxidation of sulfite and oxalate in the presence of a variety of iron oxide polymorphs. Furthermore, the difference in the p G l values (Table 3) (and therefore the ratio of >FeOH and >FeOH2+surface groups) does not seem to be an important factor in the photoreactivity, since goethite has a much lower reactivity compared to lepidocrocite (Table 21, even though their surface pKal values are very similar. When both acetate (log ,f?1:1 = 3.4) and chloroacetate (log,f?1,1=2.1) were present as competitive electron donors, the rate of photoreduction of ferrihydrite by chloroacetate was unaffected. However, ligands with a much larger stability constants for complex formation with Fe3+,such as C2042- (log = 7.6), may affect the rate of Fe(II1) photoreduction and the rate of concomitant photooxidation of haloacetates.
Conclusions Halogenated acetic acids are oxidized photochemically to the corresponding halide, glycolic acid, and C02 in the presence of avariety of iron oxide polymorphs and ambient
iron-containing aerosol particles. Fluoroacetic acid reduces solid-phase Fe(II1) to aqueous-phase Fe(I1) at the fastest rate; trihaloacetic acids seem to be less photoreactive than dihalo- and monohaloacetic acids. Strong kinetic isotope effects observed for the photooxidation of ClCDZC02H suggest that the oxidation of the mono- and disubstituted haloacetic acids proceeds via hydrogen-atom abstraction by surface-bound hydroxyl radicals to produce haloacetate radicals, which in turn yield the corresponding halide and glycolic acid. Fully halogenated haloacetic acids appear to be oxidized via a photo-Kolbe mechanism to yield the corresponding halo acids and COZ. These mechanisms appear to be consistent with the observed trends of reactivity with respect to the photochemical reduction of solid-phase Fe(II1) (Le., fluoroacetic acid > chloroacetic acid > bromoacetic acid =- iodoacetic acid and monochloroacetic acid dichloroacetic acid). The observed kinetic isotope effects indicate that the principal chromophore is the bulk iron oxide phase.
Acknowledgments The authors wish to thank Prof. J. Morgan and Mr. Scot T. Martin for helpful discussions and Mr. Sam Webb for his help in the photoreduction experiments. Support for this research has been provided by a grant from the National Science Foundation, Division of Atmospheric Sciences, Atmospheric Chemistry Section (ATM 9015775).
Literature Cited (1) Taylor, S. R.; McLennan, S. M. The Continental Crust: Its
Composition and Evolution, 2nd ed.; Blackwell Scientific Publications: London, 1985; pp 9-52. (2) Jacob, D. J.; Hoffmann, M. R. J. Geophys. Res. 1983, 88, 6611. (3) Martin, L. R.; Hill, H. W. Atmos. Environ. 1987, 21, 1487. (4) Hoffmann, M. R.; Jacob, D. J. In Kinetics and Mechanism of the Catalytic Oxidation of Dissolved SOz in Atmospheric Droplets: FreeRadical, Polar and PhotoassistedPathways, 1st ed.; Calvert, J. G., Ed.; Butterworth Publishers: Boston-London, 1984; pp 101 - 172. (5) Weschler, C. J.; Mandich, M. L.; Graedel, T. E. J. Geophys. Res. 1986, 91, 5189. (6) Graedel, T. E.; Mandich, M. L.; Weschler, C. J. J. Geophys. Res. 1986, 91, 5205. (7) Conklin, M. H.; Hoffmann, M. R. Environ. Sci. Technol. 1988,22, 899. (8) Jacob, D. J.; Gottlieb, E. W.; Prather, M. J. J. Geophys. Res. 1989, 94, 12975. (9) Martin, L. R.; Hill, M. W.; Tai, A. F.; Good, T. W.J. Geophys. Res. 1991, 96, 3085. (10) Zuo, Y.; HoignB, J. Environ. Sci. Technol. 1992, 26, 1014. (1 1) Pehkonen, S. 0.;Erel, Y.; Hoffmann, M. R. Environ. Sci. Technol. 1992, 26, 1731. (12) Erel, Y.; Pehkonen, S. 0.;Hoffmann, M. R. J. Geophys. Res. 1993, 98, 18423. (13) Siefert, R. L.; Pehkonen, S. 0.;Ere1,Y.;Hoffmann, M. R. Geochim. Cosmochim. Acta 1994, 58, 3271. (14) Pehkonen, S. 0.; Siefert, R. L.; Erel, Y.; Webb, S.; Hoffmann, M. R. Environ. Sci. Technol. 1993, 27, 2056. (15) Faust, B. C. Environ. Sci. Technol. 1994, 28, 216A.
VOL. 29, NO. 5,1995 I ENVIRONMENTAL SCIENCE &TECHNOLOGY
1221
(16) Kieber, R. J.; Helz, G. R. Enuiron. Sci. Technol. 1992, 26, 307. (17) Sedlak, D. L.; Hoigne, J. Atmos. Enuiron. PartA 1993, 27, 2173. (18) Hampson, R. F.; Kurylo, M. J.; Sander, S. P. Evaluated Rate Constants for Selected HCFCs and HFCs with OH and O(lD): World Meteorological Organization Global Research and Monitoring Project, Report No. 20; Scientific Assessment of Stratospheric Ozone: 1989;Vol. 11,Appendix: AlternativeFluorocarbon Environmental Acceptability Study Report. (19) Mead, R. J.; Moulden, D. L.; Twigg, L. E. Aust. 1.Biol. Sci. 1985, 38, 139. (20) Wu, T. J.; Mcarthur, N. H.; Harms, P. G. Ann. N. Y. Acad. Sci. 1991, 633, 626. (21) Clarke, D. D. Neurochem. Res. 1991, 16, 1055. (22) Schwarzenbach, R. P. Personal communication, 1994. (23) Helgeson, H. C. Appl. Geochem. 1992, 7, 291. (24) Rook, J. J. Environ. Sci. Technol. 1977, 11, 478. (25) Miller, J. W.; Uden, P. C. Enuiron. Sci. Technol. 1983, 17, 150. (26) Christman, R. F.; Horwood, D. L.; Millington, D. S.; Johnson, J. D.; Stevens, A. A. Enuiron. Sci. Technol. 1983, 17, 625. (27) Wine, P. H.; Chameides, W. L. In World Meteorological Organization Global Research and Monitoring Project, Report No. 20; Scientific Assessment of Stratospheric Ozone: 1989;Vol. 11, Appendix: Alternative Fluorocarbon Environmental Acceptability Study Report. (28) Hu, J. H.; Shorter, J. A.; Davidovits, P.; Worsnop, D. R.; Zahniser, M. S.; Kolb, C. E. 1.Phys. Chem. 1993, 97, 11037. (29) Heller, H. G.; Langan, I. R. 1.Chem. SOC.Perkin Trans. 1 1981, 341. (30) Faust, B. C.; Hoffmann, M. R. Environ. Sci. Technol. 1986, 20, 943. (31) Leland, J. K.; Bard, A. J. 1.Phys. Chem. 1987, 91, 5076. (32) Siffert, C.; Sulzberger, B. Langmuir 1991, 7, 1627. (33) Smith, R. M.; Martell, A. E. Critical Stability Constants, 2nd ed.; Plenum: New York, 1976; Vol. I. (34) Schwarzenbach, R. P.; Gschwend, P. M.; Imboden, D. M. Environmental Organic Chemistry, 1st ed.; Wiley: New York, 1992; p 289.
1222 m ENVIRONMENTAL SCIENCE &TECHNOLOGY / VOL. 29, NO. 5 , 1 9 9 5
(35) Hoffman,A. J.; Carraway, E. R.; Hoffmann, M. R. Enuiron. Sci. Technol. 1994, 28, 776. (36) Carraway, E. R.; Hoffman, A. J.; Hoffmann, M. R. Enuiron. Sci. Technol. 1994, 28, 786. (37) Birch, W. D.; Pring, A.; Reller, A.; Schmalle, H. W. Naturwissenschaften 1992, 79, 509. (38) Birch, W. D.; Pring, A.; Reller, A.; Schmalle, H. W. Am. Miner. 1993, 78, 827. (39) Stumm, W.; Sulzberger, B. Geochim. Cosmochim.Acta 1992,56, 3233. (40) Schwertmann, U.; Cornell, R. M. Iron Oxides in the Laboratory, 1st ed.; VCH: Weinheim, Germany, 1991; pp 6-12. (41) Schwarzenbach, R. P.; Gschwend, P. M.; Imboden, D. M. Environmental Organic Chemistry, 1st ed.; Wiley: New York, 1992; p 356. (42) Perrin, D. D.; Dempsey, B.; Serjeant, E. P. pKa Prediction for Organic Acids and Bases, 1st ed.; Chapman and Hall: London, 1981. (43) Perrin, D. D. Stabiliy Constants of Metal-Ion Complexe'i-Part B Organic Ligands, 1st ed.; Pergamon Press: New York, 1979. (44) Dzombak, D. A.; Morel, F. M. M. Surface ComplexationModelingHydrous Ferric Oxide, 1st ed.; John Wiley & Sons: New York, 1990. (45) James, R. 0.;Parks, G.A. In Characterization ofAqueous Colloids by Their Electrical Double-Layer and Intrinsic Surface Chemical Properties in Surface and Colloid Science Vol. 12; Matijevic, E., Ed.; Plenum: New York, 1982; p 185.
Received for review July 18, 1994. Revised manuscript received January 9, 1995. Accepted January 17, 1995.@
ES9404400 @Abstractpublished in Advance ACS Abstracts, March 1, 1995.
