Environ. Sci. Technol. 1994, 28, 2161-2169
Redox Interactions of Cr(V1) and Substituted Phenols: Kinetic Investigation Michael S. Eiovitzt and William Fish'
Department of Environmental Science and Engineering, Oregon Graduate Institute of Science & Technology, 20000 N.W. Walker Road, P.O. Box 91000, Portland, Oregon 97291-1000 The kinetics of the reduction of Cr(V1) to Cr(II1) by substituted phenols in aqueous solution were studied over environmentally relevant ranges of reactant concentrations, pH, temperature, and ionic strength. At a fixed pH, the reaction was first order with respect to both the concentration of phenol reductant and the total concentration of monomeric Cr(V1). Reaction rates increased as much as 4 orders of magnitude from pH 5 to pH 1. The apparent reaction order with respect to [H+l varied between 0.2 and 2 over this same pH range, but also depended on the structure of the phenol. Comparison of the reactivities of 14 substituted phenols revealed a strong substituent effect. A t pH 2, rates for the 14 phenols spanned more than 5 orders of magnitude with reactivity generally increasing with the electron-donating character of the substituent (methoxy > methyl > chloro, aldehyde > nitro). Substituent effects are described by a linear correlation of the second-order rate constants ( ~ A ~ o Hwith ) the phenol half-wave potential (Elp). Representative reaction half-lives fall within the time scales of major environmental transport processes and suggest that the kinetics of these interactions are of significant concern in assessing sites contaminated with chromium and phenolic compounds.
Introduction Few hazardous waste sites are adequately characterized as single-component systems. Complex mixtures of organic and inorganic materials are susceptible to chemical reactions that alter the toxicity and mobility of the parent compounds. Transport models of these systems rarely include abiotic oxidation-reduction (redox) transformations because little information is available about the kinetics and mechanisms of this important class of reactions. Widespread industrial use of redox-active compounds such as hexavalent chromium [Cr(VI)I and phenols has resulted in numerous releases of these compounds into the environment. Previous research on the redox reactions of these compounds has focused on synthetic organic or industrial processes, so few data are available for the relatively dilute conditions typical in the environment. More information is needed about the rates and mechanisms of reactions under environmentally relevant conditions. Cr(V1) is a common industrial oxidizing agent and a significant environmental contaminant (1-3). The two environmentally stable oxidation states, Cr(V1) and Cr(1111, exhibit very different toxicities and mobilities. Cr(111)is relatively insoluble in aqueous systems and exhibits little or no toxicity ( 4 ) . In contrast, Cr(V1) usually occurs as highly soluble and highly toxic chromate anions (HCr04or Cr02-) (5)and is a suspected carcinogen and mutagen (6, 7). * Corresponding author; e-mail address:
[email protected]. Present address: U.S. Environmental Protection Agency, Environmental Research Laboratory, College Station Road, Athens, GA 30613; e-mail address:
[email protected]. 0013-936X/94/0928-2181$04.50/0
0 1994 American Chemical Society
Phenols are common in the environment, stemming from both natural and anthropogenic sources. Many phenols exhibit toxic effects at low concentrations (81, and 11 appear on the EPA list of priority pollutants (9). Phenolic compounds are relatively reactive and are commonly used in industrial processes. Furthermore, phenolic moieties are a major constituent of natural organic and humic matter. Mixtures of Cr(V1) and phenolic derivatives are found in the waste streams of several industries (10, 11). The wood products industry utilizes large quantities of creosotes, cresols, chlorophenols, and Cr-Cu-As (chromated copper arsenate) for wood treatment (12). Leather tanning facilities use many of the same compounds [phenol, chlorophenols, and Cr(V1) salts] as tanning agents (13). Even when phenolic co-contaminants are not mixed with Cr(VI), Cr(V1) may react with the phenolic moieties in natural organic matter (NOM). NOMmay be the primary electron source for Cr(V1) reduction in many soils and natural waters (14, 15). This paper reports on kinetic experiments that quantify the effects of reactant concentration, pH, temperature, p h F 1 structure, and ionic strength on reaction rates. The results obtained allow field-site investigations and modeling efforts to quantitatively incorporate redox reactions into transport models when assessing the fate of Cr(V1) and phenols in the environment.
Experimental Section Materials. Aqueous solutions of Cr(V1) and organic reactants were freshly prepared for each experiment from pure reagents or concentrated stock solutions. Pure phenols (Aldrich) were analyzed for significant impurities by gas chromatography with flame ionization detection and were purified by recrystallization if necessary. KzCrz07 (Aldrich, ACS primary standard) stock solution was the source of Cr(V1)for all experiments. Aqueous buffers from pH 1.0to pH 6.0 were prepared from phosphate buffer or phosphate and acetate buffer solutions adjusted with HCIO4 to within 0.01 pH unit of the desired value. Ionic strength was maintained with addition of KC104. All solutions were prepared using ultrapurified 18 MO water (Nanopure, Barnstead). Kinetic Experiments. Reactions were carried out over a wide range of Cr (VI),phenol, and proton concentrations. For experiments with excess phenol, reactions were initiated by mixing a small volume of concentrated Cr(V1) stock solution into aqueous phenol solutions. Reaction advancement was followed by analyzing small aliquots for Cr(V1) by a diphenylcarbazide (DPC) colorimetric procedure (16). All Cr(V1) measurements were compared to standards prepared in buffers without phenol. Reactions with an excess of Cr(V1)were initiated by adding a small volume of concentrated phenol stock solution to aqueous solutions of Cr(V1). Time-course concentrations of phenol were determined by HPLC using a reverse-phase column (Supelcosil25-cm LC-18; Supelco). UV detection Environ. Scl. Technol., Vol.
28, No. 12, 1994 2181
-0.5
-1.5
f
-2 E
Y
-2.5
6
I
-c
-3.5
2.0 I
0.02
-4.5 0
5
10
15
I
I
1
I
I
0.03
0.04
0.05
0.06
0.07
[4-methylphenoI] (M)
time (h) Figure 1. Pseudo-first-order rate plots of in [Cr(VI)] versus time for Cr(V1) reduction by excess 4MP for 15,25, and 35 "C. Slopes of the (defined lines are equal to the pseudo-first-order rate constants kexp in text). Values of ?from linear regressionswereall >0.999. Reactions performed at pH 2.0, I = 0.1 M, [4MP] = 0.05 M, and [Cr(VI)lo = 0.5 M.
was performed at either 280 nm or, when increased sensitivity was necessary, at the specific,,A of the compound. For experiments in which the molar ratio of Cr(V1) and phenol did not satisfy an excess condition (stoichiometric experiments), substrate concentrations were followed by HPLC and Cr(V1) concentrations were followed by both DPC and HPLC with UV detection of Cr(V1) at 280 or 350 nm. Reactions were carried out in glass vials with screw caps containing Teflon liners. Glassware was washed in 3 N HC1and rinsed thoroughly with purified water. All kinetic experiments were maintained at a constant temperature [15, 25, or 35 "C (k0.5 "C)I in a circulating water bath. Appropriate reaction blanks consisted of either Cr(V1)or the selected phenol without the addition of the other reactant in pH buffer solutions in both clear and amber vials. The blanks showed no depletion of Cr(V1) or degradation of phenol over the course of the sampling period. Light-control reactions in amber vials showed no difference in reaction rates compared to clear vials, indicating the absence of light-induced side reactions. All reactions and corresponding blanks and controls were performed in duplicate or triplicate. Solution pH was measured with a microcombination electrode (Corning) at the beginning and end of each reaction. The pH change was less than 0.1 unit in all cases, indicating a sufficient buffer capacity. Kinetic Analysis. Reaction kinetics were simplified in many cases by using one reactant concentration in large excess (20-500-fold) and by maintaining constant pH with unreactive buffers. Under these conditions, the reaction kinetics were fixed with respect to the concentration of H+and the excess reactant (pseudo-zero-orderconditions), thus isolating the rate dependence of the limiting reactant. Pseudo-first-order rate constants for the variable reactant, .Itexp, were determined from least-squares linear regression of first-order plots of In(C/Co) as a function of time. Activity corrections for secondary salt (ionic strength) 2162
Envlron. Scl. Technol., Vol. 28, No. 12, 1994
0 8
20 Flgure 2. Linear plot of pseudo-first-order rate constants, kSw,as a function of [4MP] for Cr(V1) reduction at pH 2.0 and 25 OC. [Cr(VI)lo = 0.5; [4MP] = 0.025, 0.0375, 0.05, 0.085, and 0.075 M.
effects on Cr(V1) speciation were calculated from the Davies equation (17). Results and Discussion
Oxidation-Reduction Kinetics. Cr(V1)reacts rapidly with phenolic compounds at moderately acidic pH values, causing reduction of Cr(VI), presumably to Cr(III), and conversion of phenols to oxidized products. Reactions of Cr(V1)and excess phenol are characterized by a nonlinear reduction of Cr(V1) with time. The rate of Cr(V1) reduction, indicated by the slope of the reaction curve, decreases with time due to the depletion of Cr(V1). The linearity of the semi-logarithmic first-order plots and statistical goodness-of-fit verified the suitability of the pseudo-first-order assumption (Figure 1). The redox interactions of Cr(V1) and 4-methylphenol (4MP) comprised the primary focus of the kinetic investigations. Cr(V1) reduction rates, as quantified by kexp, varied linearly with the concentration of excess 4MP, indicating that the reaction was first-order with respect to the reductant concentration (Figure 2). Thus, at a fixed pH, the overall reaction could be characterized by a secondorder rate expression d[Cr(VI)l/dt = -kAroHICr(VI)l[4MPl
(1)
where the second-order rate constant, kAroH, is defined as
Studies of Cr(V1) oxidation of primary and secondary alcohols (18-29) as well as tertiary alcohols (30-32), aldehydes (33-361, thiols (37),and organic acids (38)have demonstrated a first-order dependence with respect to both Cr(V1) and the organic reductant, as was observed here for Cr(V1) and phenols. Oxidations of phenolic compounds by non-chromyl oxidants have likewise displayed a first-order dependence on the phenol concentration (39-41 1. First-order dependencies of the oxidant and reductant are consistent among studies despite significant differences in the reactants, and the observations of overall
not linearly proportional to [Cr(VI)lht, and reaction rates are therefore expected to show a nonlinear variation with A 0.04 M Cr(VI) [Cr(VI)ltot. Experiments with excess Cr allowed speciation effects 0 0.06 M Cr(VI) to be examined because [Cr(VI)ltotand, therefore, Cr(V1) 0 0.08M Cr(VI) speciation were essentially constant over the course of the reaction. Reaction rates thus can be compared with the fixed concentrations of each Cr species. Values of kexp for 4MP were linear or nearly linear with respect to concentrations of the monomeric species Cr042-,“ 3 0 4 - , and HzCr04 (Figure 4). Slopes of the plots in Figure 4 normalize the pseudo-first-order rate constants to the change in a particular Cr(V1)species; we term these slopes the species-normalized second-order rate constants. It is not possible to conclusively identify the primary reactive Cr(V1) species simply from the goodness of linear fit of such a plot. However, if we compare the species-normalized second-order rate constants for 4MP oxidation to the 30 60 , apparent second-order rate constant for Cr(V1)reduction time (min) ( k h o ~calculated , by eq 2), only the HCr04- normalized (kHCrOd-) falls within the same order of rate constant Figure 3. Pseudo-first-order linearization of 4MP oxidation by various magnitude as ~ A ~ O Hand , the two values are virtually concentrations of excess Cr(V1) at pH 1.0, 25 ‘C. Reaction half-life identical for pH 1and 2 (Table 1). This finding strongly f , / 2 (Calculated as in 2/k,,,)and P were calculated by linear regression. Cr(V1) = 0.02 M: fIl2 = 21.7 min, P = 0.998: Cr(V1) = 0.04 M: f112 suggests that HCr04- is the active Cr(V1) species. = 13.6 min, P = 0.996; Cr(V1) = 0.06 M: fIl2 = 11.0 min, P = 0.996; In most experiments, the ionic strength was maintained Cr(V1) = 0.08 M: f I l 2 = 9.1 min, P = 0.997. at 0.1 M. However, in certain excess Cr(V1) experiments the high concentrations of Cr(V1)raised the ionic strength second-order kinetics suggest a bimolecular process in the as high as 0.64 M. It therefore was necessary to investigate rate-limiting step. the influence of ionic strength on Cr(VI)/4MP oxidation The first-order dependence on [4MPl was verified in a rates. Additions of 0.1,0.5,1.0, and 2.0 N NaC104or KC104 second series of experiments with eicess Cr(V1). For these to reactions had no significant effect on the reaction rate; pseudo-zero-order conditions with respect to Cr(VI), we conclude there is no primary salt effect. Additional reaction progress was monitored by the loss of 4MP for experiments with additions of CrC13 showed that Cr(II1) several concentrations of excess Cr(V1). Plots of ln[4MPl did not inhibit the rate of Cr(V1) reduction. vs time were linear over at least three reaction half-lives (Figure 3). However, the observed rates in these experiPhenol Speciation Effects. The acid-base chemistry ments were not linearly proportional to .total Cr(V1) of the phenols could contribute to a pH dependence of the concentration. The reaction rates instead increased less reaction rates. The deprotonated anion of a hydroxylated than the proportional increase in the initial Cr(V1) organic compound frequently has been observed to be a concentration (Figure 4). As discussed below, this phemore reactive reductant than the undissociated form. nomenon implies that 4MP reacts at different rates with Tratnyek and Hoign6 investigated a wide variety of different Cr(V1) species and that the nonlinear trend in phenolic compounds in both states and showed a reactivity kexpis a reflection of changing Cr(V1) speciation. Proper difference of up to 2 orders of magnitude for oxidation by interpretation of the kinetic data therefore depends on singlet oxygen (40) and up to 6 orders of magnitude for careful consideration of the active reactant species. oxidation with chlorine dioxide (44). For hexachloroiriCr(V1) Speciation Effects. Aqueous Cr(V1) exists as date(1V)oxidation of phenol and 2,&dimethylphenol,Cecil five species in this system: HzCr04, HCr04-, C T O ~ ~ - , and Littler (45)showed that the phenolate species was 6 Cr2072-, and HCrz07-. Previous reports have suggested orders of magnitude more reactive than the phenol. The that bichromate, HCr04-, is the primary reactive species increased reactivity of the anion presumably reflects a (18,29,30,33,34), although this has not been conclusively higher electron-donating potential due to its higher established for all systems (29,36,42,43).Our experiments electron density. However, phenolate anions are unlikely with excess Cr(V1)indicate that a monomeric species (Hzto be important in the present study. The pKa values of CrO4, HCrOd-, or ( 3 0 4 % ) is the primary reactant with the phenols used here are 29.0, with the exception of phenol. To achieve the necessary excess of [Cr(VI)I over 4-nitrophenol (7.16) and vanillin (7.4) (Table 2). Experi[4MPl within the constraint of the analytical detection mental solutions were poised at pH values 5-9 orders of limit for 4MP, relatively high concentrations of Cr(V1) magnitude below the PKa of the reactant phenol. Hence, (0.01-0.5 M) were required. In this concentration range, concentrations of phenolate were minute (0.98. The pH dependence of the empirical reaction order suggests the presence of a rapid acid-base equilibrium involving the reactants or the existence of parallel reaction pathways involving different stoichiometries of hydrogen ions. Acid-base equilibria are important only in a pH region close to the pK, of the reactant in question, usually (pK, - 2) < pH < (pK, + 2). For the pH region studied, the HCr04-/Cr042-acid-base equilibrium, HCrO4- = H+ + Cr042-;pK, = 6.52, meets this criterion, but only for pH >4. Hence, an antecedent protonation of Cr042- and subsequent reaction of HCr04- is unlikely to cause the variable proton dependence observed for pH 1-4. At low pH, the reaction order with respect to [H+l is most likely due to parallel reaction pathways involving different numbers of protons. In addition, if HCr04- is the sole reactive Cr(V1) species, and the only involvement of the proton is in the acid-base equilibria of HCr04- (e.g., in a preoxidation protonation of CrO12-to HCrOd-), a uniform pH dependence among the phenols would be expected. The different empirical order curves we observed for different phenols imply that some or all of the involvement of the proton is with the phenolic reactant or a reaction intermediate. Because of the strong pH dependence of the Cr(V1)phenol reactions, we explored the possibility of acid catalysis. Reactions were conducted with equimolar
concentrations of Cr(V1) and 4MP in pH 1.0 solutions of H3P04,H2S04,HC1, and HC104 without additional buffer. The solutions containing the strong, completelydissociated acids HC1, H2S04,and HC104 all reacted at the same rate. Reactions with H3P04 proceeded more rapidly. H3P04 is not completely dissociated at pH 1 and so has a higher concentration of total acid (approximately 1.4 M). Because the proton activity is the same for both the strong acid and weaker acid systems, faster rates with the H3P04 system must be due to the general acid-catalyzing effects of the undissociated form of the Brransted acid (H3P04). At lower concentrations of phosphate buffer (0.05 M), no general catalysis was observed: reactions at pH 1and pH 2 in HC104showed no difference in rates with the addition of 0.05 M phosphate. Thus, at the concentrations of phosphate buffer used in most experiments, general acid catalysis was minimal and did not affect reaction rates. Temperature Effects. The temperature dependence of Cr(VI)-4MP reactions was studied in the range 288308 K using an excess of reductant and a fixed pH. Pseudofirst-order rate constants, kexp, were determined for pH 1.0,2.0, and 3.0 and for T = 288,298,and 308K. Activation enthalpy and entropy (AH*and AS*)were calculated from (4) where h is the Planck constant, k is the Boltzmann constant, and R is the universal gas constant, by plotting ln(kexp/T)versus 1 / T (46) (Table 4). Our values (Table 5) are consistent with those reported for other systems employing similar reaction conditions and kinetic analysis, which implies the operation of similar reaction pathways prior to and including the rate-limiting step. Because the [H+l term or its function is not explicitly included in the kexp values, the activation parameters AH* and AS* must be taken as composite values for all reactions prior to and including the rate-limiting step (47,48)and may not describe the activation energy of an elementary reaction step. However, these values do allow us to examine the overall effect of pH on the thermodynamics of the reaction. All AS* values are similar and are independent of the initial proton concentration, whereas AH*values increase with initial pH. Specificacid catalysis may explain this result. By definition, a catalyst lowers the activation energy of the reaction (AH*)without affecting the overall thermodynamics (Le., AGO,,,) (49). An alternative explanation of the apparent shift in AH* is the hypothesis of parallel reaction pathways. Elevated activation energy associated with increasing pH can be caused by a shift from low energy to high energy pathway(SI.
Stoichiometry. The overall stoichiometry of the Cr(V1)-phenol reaction is complicated by a possible multistep mechanism and a potential for side reactions involving primary reaction products. To minimize side reaction effects, calculations were limited to data from the first 50% of the reaction. For a 1mM:5 mM initial ratio of Cr(V1) to 4-methylphenol, the reaction stoichiometry was calculated from the slope of [Cr(VI)I versus 14MPI. The value of A[Cr(VI)l/A[4MPl was 0.642, very close to 213. On average, each phenol molecule appears to be oxidized by two electron equivalents for each Cr(V1) reduced to Cr(II1). The net mean oxidation of the phenol could correspond to a single two-electron transfer or to two successive one-electron transfers. When equimolar Environ. Sci. Technol., Vol. 28, No. 12, 1994 2165
Table 3. Second-Order Rate Constants for Cr(V1) Reduction by 4-Methyl-, 4-Methoxy-, and 2,6-Dimethoxyphenol at Various Hydrogen Ion Concentrations and 25 "C khoH (M-'
[H+l (M)
PH
1.33 X 1.39 X 1.33 X 4.21 x 1.38 x 1.79 X 1.81 X 1.24 X
1.00
1.98 2.00 2.50 2.98 3.87 4.85 6.01
4-methyl
10-l
(5.15 f 0.06) X
10"
(1.10f 0.02) x 10-3
10-3 10-3 10-4
(1.21 f 0.01) x 10'' (4.57 f 0.04) X lo4
10-5 10-8
0
in
-3
0
22
-4
\
I
I
I
I
1
2
3
4
5
Table 4. Activation Data for Cr(V1) Reduction by 4-Methylphenol at pH Values of 1,2, and 38
(kJ mol-l) f SD AS: (J K-l mol-') 35.0 42.3 54.9
1.6 1.1
1.7
-183.4 -185.1 -181.6
f SD
1.0
0.8 0.8
[Cr(VI)], = 0.5 mM, [4MP] = 0.05 M, t = 15-35 "C. ~
initial concentrations (5 mM) of Cr(V1) and 4MP were used, a consistent 1:l stoichiometry was observed within the initial 50+% of the reaction. A 1:l stoichiometric ratio implies that each organic molecule is oxidized by an average of three electron equivalents, indicating products more oxidized than in the reaction for which the stoichiometry is 2:3. Apparently the 5-folddecrease in the relative concentration of the organic leads to its more complete oxidation. Because the reaction is first-order in [Cr(VI)I and [phenol] for both excess oxidant and excess reductant, the net stoichiometry leading up to the activated complex associatedwith the rate-limiting step is probably a constant 1:l ratio. Deviation from 1:l in the overall stoichiometry therefore must be due to intermediate reaction steps that occur after the rate-limiting step. 2168
10"
(8.58 f 0.12) X 10-* (5.40 f 0.04) X (1.45 f 0.01) X (7.28 f 0.08) X 10-3 (4.80 f 0.06) X
Environ. Scl. Technol., Vol. 28, No. 12, 1994
(kJ mol-') 35-54 49.0 53.2 26.6 49.4 24.7 29
-AS$(J K-l mol-')
ref
181-185 120.2 107.7 180.8 105.0 155.2 167
this work 29 29 29 43 43 49
6
+
1.288 X 10-l 1.273 X lo-* 1.273 X
10-2 10" lo4
I
I
-
PH Flgure 5. Plots of log kAroH versus pH for the reduction of Cr(V1) by excess 4-methyL 4-methoxy, and 2,6dlmethoxyphenol at t = 25 OC; I = 0.1 M. Slopes of the plots indicate the empirical reaction order = a[H+I2 with respect to [H+]. Solid lines are empirical fits to kArOH b[H+] 4- c, (4-methylphenol: a = 2.6, b = 5.8 X lo-*, c = 3.7 X 4-methoxyphenoi: a = 3.9 X lo2,b = 9.1, c = 1.8 X 2,6dimethoxyphenol: a = 1.0 X lo3, b = 20.8, c = 5.5 X
[H+] (mol L-9
*
(1.59 0.01) X (1.61 f 0.05) X (5.82 f 0.05) X (1.91 f 0.03) X
Empirical Rate Expression. The experimental results indicate a first-order dependence on both Cr(V1) and phenol concentrations; the order with respect to [H+l varied between 0.2 and 2. These findings are in accord with a second-order rate expression
4-methyl
-5
-6
(4.35 f 0.1) x 10-1
4-methylphenol ethanol isopropyl alcohol benzyl alcohol triphenylcarbinol benzhydrol glutathione
r
rn
(2.16 f 0.03) X 10-l
reductant
-2
-0
2,6-dimethoxy
Table 5. Comparison of Activation Data for Cr(V1) Reduction Reactions by Compounds with Hydroxyl or Sulfhydryl Functional Groups
-1
2I
8-l)
4-methoxy
where [HCr04-I and [ArOHl denote the concentrations of bichromate and total phenol, and the empirical secondorder rate constant, ~ A ~ O His, a function of pH. At a constant pH, eq 5 and a suitable value of khoH will accurately predict the rate of chromate reduction. An empirical rate expression of this simple form is practical from a computational perspective because it can be easily incorporated into transport models. However, the representation of ~ A ~ OasH a function of [H+l requires a quadratic formula, which somewhat negates the simplicity of the second-order rate expression. However, in a wellbuffered system in which pH varies little, K A ~ O Hcan be treatedas a constant. In addition, if either Cr(V1)or ArOH is present in sufficient excess, it may be reasonable to reduce eq 5 to an even simpler first-order expression. Substituent Effects/Reactivity Correlation Analysis. The position and composition of ring substituents greatly affects the reactivity of aromatic compounds. We investigated the comparative reactivities of 14 substituted phenols. Experiments were carried out at pH 2 under pseudo-first-order conditions ([ArOHl in excess of [Cr(VI)]). Values of kexpwere calculated from first-order kinetic plots, and second-order rate constants (kh01.1)were calculated using eq 2. Second-order rate constants for the 14 phenols range over 5 orders of magnitude (Table 2). The general reactivity trend is methoxy-substituted > methyl-substituted > unsubstituted > chlorophenol > nitrophenol. Reaction rates have been correlated to half-wave for several phenol oxidations including potentials (E112) oxidations by singlet oxygen (40)and Mn(II1,IV)(52).Halfwave potentials are an experimentally derived redox
1
0
-1
-2
-3
-4
€,,*
(V vs. SCE)
Flgure 6. Correlation of log kAm to ,Eil2(Vvs SCE) for the reduction of Cr(V1) by various substltuted phenols at pH 2.0 and 25 OC; identifier numbers correspond to the compounds and data in Table 2. Llnear regression for all compounds,excluding the outliers 2,ddimethylphenol (5), 2,4,d-trlmethylphenoI (6), 2,6dimethoxyphenol(E), and 4-nltrophenol (not shown) yielded P = 0.948. ,Ei,* values are taken from ref 54.
parameter and are believed to represent the first oneelectron oxidation potential of the phenol. Half-wave potentials for 10of our phenols were available from Suatoni et al. (53), who reported Ell2 values for 40 substituted phenols determined by anodic voltometry in 50 % aqueous 2-propanol buffered at pH 5.6 in 0.5 M acetate buffer. For the Cr(V1) system, Ell2 was a good predictor of log kArOH (Figure 6, Table 2): log kArOH = (-16.87 f 1.4)E112 + (6.25 f 0.73)
(6)
(r2= 0.948; s = 0.44; n = 10). With the exception of three outliers, a very good linear fit is achieved with rates ranging 5 orders of magnitude. The three points not included in the regression line in Figure 6 are all di-ortho-substituted phenols that have reactivities significantly less than those predicted by their reported Elp. If the hydroxyl group is the reaction center, di-ortho substitution may sterically hinder oxidation by Cr(V1) in a way not reflected in the polarographic E112 measurement. Relevance t o Environmental Systems Our results show that phenolic compounds reduce Cr(V1) in aqueous solution under moderately acidic conditions. I t is useful to examine the conditions under which these reactions occur in the field and thereby gain a sense of the environmental relevance and applications of these experiments. As discussed in the Introduction, Cr(V1)and phenols are common industrial chemicals, and both are found at a number of contamination sites. Many chromate-contaminated sites have high concentrations of contaminants and a very acid pH localized within one or more plume areas. For example, at the United Chrome Products Superfund site near Corvallis, OR, a leaking chrome plating bath created a subsurface plume environment with solution concentrations of Cr(V1) in excess of 0.2 M and a near-source pH of approximately 2 (55). However, the plume of high acidity is attenuated by the exchange capacity of the subsoil within tens of meters
from the source whereas the Cr(V1) plume extends over 100 m downgradient and thereby encompasses a range from pH 2 to pH 7. At many sites, Cr(V1) and phenols stem from spatially distinct industrial processes so the sources are likewise spatially separated and not initially mixed. During subsequent subsurface transport, Cr(V1) and phenol plumes can mix in varying proportions and over a wide range of pH. The empirical rate law developed here (eq 5) is robust with respect to [Cr(VI)l and [ArOHl because of the unit order with respect to these reactants over a wide range of concentrations. In an environment in which the pH is well buffered and nearly constant, eq 5 yields reasonable estimates of Cr(V1)-phenol reaction rates for a wide range of solution concentrations. However, because the reaction order with respect to [H+l varies sharply with pH, predictions of reaction rates across a large pH gradient may exhibit a greater degree of uncertainty. We can use eq 5 and the experimentally determined rate constants to quantify the characteristic time scales of Cr(V1) reduction by phenols and place them in the context of time scales of other environmental processes. For example, for concentrations of Cr(V1) and 4-methylphenol in the range of 0.1 M, eq 5 predicts reaction halflives of 2.7 h at pH 2 and 25 days at pH 5. For millimolar concentrations of the same components, resultant halflives increase to approximately 11days and 7 years at pH 2 and pH 5, respectively. Reaction rates also depend strongly on phenol structure. For identical concentrations of mono-, di-, and tri-substituted phenols, rates of reactions with a given concentration of Cr(V1) vary by as much as 6 orders of magnitude, depending on the composition and positioning of ring substituents (Table 2). Figure 7 compares characteristic time scales of selected reactions with other environmental processes. Time scales for the reduction of Cr(V1) by substituted phenols span a large range that encompasses characteristic transport times of both surface and subsurface waters. Re-oxidation of Cr(II1) to Cr(V1) is also possible within the time domain of reduction and transport processes. Note also that only homogeneous reactions were examined here. Future studies should explore the potential role of solid mineral phases in Cr(V1)-organic interactions. Conclusions The reduction of Cr(V1) by phenolic compounds in aqueous solution follows a general rate law of the form -d[Cr(VI)l/dt = ~A~oH[HCI’O~-] [ArOHl, where [HCr04-I denotes the total concentration of bichromate, [ArOHl is the total concentration of phenol, and kArOH is a function of pH. Although we were unable to quantitatively determine a proton dependence generally applicable to all phenols, we found evidence that the proton affects the rate by an interaction with the phenol or a reaction intermediately prior to, or probably during, the rate limiting step. Cr(II1) does not inhibit the reaction, which indicates it has no significant involvement in the ratelimiting step. Ringsubstituents strongly affect the kinetics of Cr(V1) reduction, causing over 5 orders of magnitude difference in reaction rates. In general, electron-donating substituents (methyl and methoxy groups) enhance the reactivity compared to the unsubstituted phenol, while electron-withdrawing groups (chloro, nitro, and aldehyde) diminish the reactivity of the phenol. The rates of Cr(V1) reduction correlated well with the half-wave potential Environ. Scl. Technol., Vol. 28,
No. 12, 1994 2167
Transport Processes
Time (SI
Other Relevant Reactions
Phenol Reaction with Cr(VI) (PH and Cones.)
1
lo 1 1 min
2,d-dimethoxyphenol +pH2,
T
Cr(ll)/Co(lll) (electron transfer)
1 rnM
l h
i d
2,6-dirnethoxyphenol +pH2,
4-methylphenol
1 Yr
Cr(lll)lH20 (exchange)
1 UM
r-
Cr(lll)
+
H202
= Cr(VI)
Fe(ll) + 02 = Fe(lll) Cr(lll) + 02 = Cr(Vi)
+p H 5 , 1 mM Mn(ll)
100 yr
10
_I
+
02
= Mn(lll/lV)
4-rnelhylphenol 4-pH5,
1 uM
Schematicdiagram for characteristictime scales of Cr(V1) reduction by substituted phenols, major environmental transport processes, and related envlronmental reactions. Figure adapted from ref 56. Flaure 7.
( E l p ) of the phenol reductant. This structure-activity relationship allows a semi-quantitative estimate of Cr(V1) reduction rates by other phenol compounds. Our findings of the importance of substituent position and composition may be a useful starting point for detailed investigations of the active Cr(V1) reductant moieties in natural organic matter. For example, such materials almost certainly contain phenol moieties substituted with methyl, methoxy, and aldehyde groups (57,581. Finally, by applying our empirical rate expression to a variety of solution conditions typical of sites contaminated with Cr(V1) and phenols, we established characteristicreaction time scales. Cr(V1)-phenol interactions can be expected to occur on the time scales comparable to transport time scales in surface and subsurface waters.
Acknowledgments
We thank James K. Hurst, Paul G. Tratnyek, Carl D. Palmer, A. Lynn Roberts, and two anonymous reviewers for their insightful comments and contributions to this research. We also thank Lorne Isabelle for assistance with GC-MS and other analyses and Cheryl Martin and Rachel Edelstein Martin for assistance with many experiments. Support was provided by the U.S. EPA, Office of Exploratory Research Grant R-814136-01-0. The findings, opinions, and conclusions expressed in this paper are those of the authors and do not necessarily reflect those of the U.S. EPA. Notations kexp
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pseudo-first-orderCr(V1)reduction rate coefficient Envlron. Scl. Technol., Vol. 28, No. 12, 1994
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Received for review March 8, 1994. Revised manuscript received August 1, 1994. Accepted August 5, 1994.'
* Abstract published in Advance ACS Abstracts, September 1, 1994.
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