Environ. Sci. Technol. 1992, 26, 79-85
Weast, R. C., Ed. CRC Handbook of Chemistry and Physics, 61st ed.; The Chemical Rubber Co.: Cleveland, OH 1980. Kirchner, W. Dark Reactions of Nitrogen Oxides in a Smog Chamber, Measurements and Model Calculations. Master’s Thesis, University of Bonn, 1986. Zerbach, T. Organic Nitrates in the Atmosphere. Doctoral Thesis, University of Mainz, 1990. Atkinson, R.; Baulch, D. L.; Cox, R. A.; Hampson, R. F.;
Chromium( I I I ) Oxidation by 6-Mn0,.
Kerr, J. A.; Troe, J. J. J . Phys. Chem. Ref. Data 1989,18, 881-1097. (24) Roberts, J. M. Atmos. Environ. 1990, 24A, 243-287.
Received for review March 6,1991. Revised manuscript received July 19,1991. Accepted July 22,1991. The present study has received support from the Germany Federal Ministry of Research and Technology.
1. Characterization
Scott E. Fendorf” and Robert J. Zasoski
Department of Land, Air, and Water Resources, University of California, Davis, California 95616 ~~
The oxidation of Cr(II1) by a common naturally occurring form of manganese oxide, 6-MnOZ,was characterized over a range of Cr(II1) concentrations and pH values. Cr(II1) oxidation was limited as pH and Cr(II1) concentrations increased. Reaction products, Mn(1I) and Cr(VI), did not limit Cr(II1) oxidation. Initial Cr(II1) oxidation rates were very rapid at pH = 5, but were subsequently followed by a dramatic rate decline. Thermodynamic calculations using solution species indicated the reaction should proceed under conditions where the reaction had terminated. A surface alteration induced by Cr(II1) at pH > 3.5 appears to prohibit the extent of oxidation. Various mechanisms may account for the electrophoretic mobility and oxidation reactions; however, surface precipitation of chromium hydroxide seems to be the most plausible explanation based on our results. Chromium oxidation was dependent on Cr(II1) concentration, pH, initial surface area, and ionic strength.
Introduction The environment is sensitive to heavy metals due to their longevity and toxicity. Chromium is a heavy metal that has many uses in the metallurgic, refractory, chemical, and tannery industries. Chromium oxidation states range from -2 to +6, but only the +3 and +6 states are stable under most surficial conditions. Cr(II1) is considered to be less toxic than Cr(VI), and levels greater than 1.7 ppm Cr(V1) in effluent water exceed regulatory limits (I). Although Cr(II1) compounds have no established mammalian toxicity and pose a low health risk, Cr(V1) compounds are more toxic. Cr(V1) is a carcinogen, an irritant, and a corrosive, which can be absorbed by ingestion, through the skin, and by inhalation (2). Cr(V1) is mobile in most soil and water systems, while Cr(II1) is rather immobile due to its limited solubility, sorption by negatively charged surfaces prevalent in soils and sediments, and complexation by insoluble organic material. Thus, disposal of Cr(II1) is considered to be less problematic than Cr(V1); however, the hazard of Cr(II1) is tantamount to Cr(V1) if oxidation occurs. Cr(VI) can be reduced in many soil systems by reaction with organic matter, Fe(II), and sulfide compounds. Although a variety of compounds are capable of reducing Cr(VI), the only common naturally occurring oxidizing agents of Cr(II1) include molecular oxygen and manganese oxides (3). Small amounts of Cr are oxidized by O2 at pH values greater than 9, but these conditions are uncommon in soil or water environments. However, Cr(II1) oxidation in soils was only hypothesized, and first studies ( 4 ) indi-
* Present address: Dept. of Plant and Soil Sciences, University of Delaware, Newark, DE 19717-1303. 0013-936X/92/0926-0079$03.00/0
cated Cr(II1) was not oxidized in soils. Bartlett and James (5)demonstrated that Cr(II1) could be oxidized using fresh rather than dried aged soil samples. It appears that manganese oxides are the principle oxidizing agents for Cr(II1) in soil systems. Bartlett and James (5)and Amacher and Baker (6)listed possible half-reactionsfor the oxidation-reduction reaction involving Cr(II1) and manganese oxides. Using half-reactions which were assumed to be representative of many environmental conditions, Amacher and Baker (6) proposed an overall reaction involving Mn02 and Cr3+,forming HCr04-: Cr3+ + 1.56-Mno2 + HzO = HCr04- 1.5Mn2+ H+ (1) The forward reaction is not thermodynamically favorable under standard conditions [AGO = 2.48 kJ mol-l using thermodynamic values from Garrels and Christ (7)]. However, for conditions reflective of surface environmental conditions, AG for the reaction can be negative. Amacher and Baker (6) examined surface effects on rates of Cr(II1) oxidation by 6-Mn02. At pH 2 5.5, rapid initial Cr(II1) oxidation rates were followed by decreasing rates and cessation of the reaction prior to complete utilization of either Cr(II1) or Mn02. They suggested that the decreased rates could only occur if the available reactive surface was “used up” as the reaction proceeded. Increasing Mn02levels retarded the onset of the reaction rate decrease, thus enhancing the extent of Cr(II1) oxidation. Amacher and Baker (6) postulated that competition for adsorption sites between Cr(II1) and Mn(I1) limited Cr(II1) oxidation. These authors assumed a weak association between Cr(II1) and 6-MnO2,followed by a rapid oxidation and subsequent repulsion of negatively charged Cr(V1). However, 6-MnO2has a demonstrated strong affinity for various cationic metals (8-12). Solution data indicated 1.5Mn(II) was released for each Cr(V1) formed, which supported the proposed stoichiometry. Therefore, since all the Mn(I1) expected to form was found in solution, it does not appear likely that adsorption of Mn(I1) inhibited Cr(II1) oxidation. Effects of pH on Cr(II1) oxidation are unresolved. The rate of Cr(II1) oxidation by 6-Mn02 increased with pH until pH = 5.5; thereafter, the rates were not altered by increased pH (6). In contrast, equilibrated soil systems have shown that the extent of Cr(II1) oxidation decreased with increased pH (5). The rate of Cr(II1) oxidation by P-Mn02 (pyrolusite) was found to increase with decreasing pH (3). Eary and Rai (3) developed a model for the rate of Cr(II1) oxidation by pyrolusite, which was based on Cr(V1) levels as the limiting parameter in the oxidation process. Pyrolusite would have a net positively charged surface throughout the pH range of their study. Eary and
0 1991 American Chemical Society
+
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79
Rai (3) postulated that Cr(V1) was readsorbing to the positive surface of the oxide and inhibiting further Cr oxidation. However, oxidation increased at lower pH values, contradicting an expected stronger adsorption of Cr(V1) at lower pH values where the surface was more positive. Chromium oxidation by manganese oxides has been demonstrated to occur under environmental surface conditions. However, the oxidation process does not proceed until the reactants are limiting. A surface phenomenon appears to limit the extent of Cr(II1) oxidation, but the limiting factors are not clear. Therefore, the intent of this research was to (i) characterize Cr(II1) oxidation by a commonly occurring form of manganese dioxide (6-Mn02) and (ii) determine factors (if any) which limit the oxidation process.
Methods and Materials 6-Mn02 was prepared by adding HC1 to an excess of KMn04 (13). The oxide produced was thoroughly washed with deionized water, dried at 60 "C, ground, and passed through an 80-mesh sieve. Specific surface area of the MnOz determined by the ethylene glycol monoethyl ether (EGME) technique (14) was 223 X lo3 m2 kg-l. Potentiometric titration (15) estimated a zero point of charge (ZPC) of 2.7. Random powder amounts were used for X-ray diffraction analysis. A Cu target and Ni filter were used according to the procedures of Whittig and Allardice (16). Four broad peaks characteristic of 6-Mn02were observed; the most intense peak was at 0.73 nm with weaker signals at 0.36, 0.24, and 0.14 nm. These values are in good agreement with d spacing for 6-Mn02 found by other workers (17, 12). Static Studies. Static studies were performed using batch methods to obtain (quasi) steady-state data. Twenty-five milliliters of solution and 2.5 mg of Mn02were placed in 50-mL polyethylene tubes to yield an initial solid concentration of 0.1 g L-l. Each run consisted of dispensing the oxide in the reaction vessels, adding supporting electrolyte, hydrating for 24 h, adjusting pH, and finally adding C T ( N O ~solutions, )~ which brought the final volume to 25 mL. Preliminary studies indicated the reaction reached a steady state within 24 h. After a 24-h shaking period, the suspensions were filtered through a 0.45-pm pore membrane filter and the filtrates analyzed. The supporting electrolyte was purged of C02 with N2(g)prior to the experiment, and the systems were maintained in a N2 atmosphere for the duration of the experiments at laboratory temperature (25 f 4 "C). The Cr(N0J3 stock solution was made from reagent grade material, and stored at pH = 2 for no longer than 1 week. The static experiments were carried out in 0.1 and 0.001 M NaN03. The higher ionic strength was used to acquire the proper conductivity and current flow for the electrophoretic cell and to elucidate ionic strength effects on the oxidation process. To be consistent with the electrokinetic data, steady-state data are reported from 0.1 M NaNO, matrices except where otherwise indicated. Data are not reported in terms of solution to solid ratios because Mn02 is consumed in the oxidation-reduction reaction with Cr(II1). Thus, adsorption is reported in terms of micromoles of metal removed from solution. Hexavalent Cr was determined by a modified s-diphenylcarbazide procedure (5). The reagent was made by mixing 120 mL of 85% H3P04 diluted with 280 mL of water and adding this solution to 0.2 g of s-diphenylcarbazide dissolved in 100 mL of 95% ethanol. One milliliter of this reagent was added for each 10 mL of 80
Environ. Sci. Technol.. Vol. 26,No. 1, 1992
Cr(V1) solution to be analyzed. The color was allowed to develop for 20 min and was read at 540 pm on a Hitachi U-2000 UV-vis spectrophotometer. Solution Mn was determined by atomic absorption using a Perkin-Elmer 508 spectrophotometer with an air-acetylene flame. Total solution Cr was determined using a sequential ICP-optical emission spectrophotometer. Electrokinetic Investigation. The electrophoreticc mobility of the MnOz particles were measured using conventional procedures (18,19) and a Rank Brothers apparatus. The stationary layer was located by initially focusing on the theoretical stationary plane 293 pm from the inner wall of the electrophoretic cell. The plane of focus was then adjusted to the actual stationary layer using human erythrocytes as a standard of known mobility, 1.27 (pm s-l)/(V cm-l) in 0.067 M phosphate buffer at pH = 7.4 at 25 "C (18). The actual stationary layer was located a nominal distance of 305 pm from the inner wall of the electrophoretic cell. Electrophoretic mobility of the Mn02 was measured at 25 "C, with a current flow of 1.0-1.3 mA and an applied voltage of 48-72 V. The rate of particle movement was found by timing the movement over a fixed number of graticule divisions selected to give readings near 10 s at an applied potential gradient of 3-5 V cm-l. Current direction was reversed halfway through each reading to minimize polarization, and the mobility values were based on 10 individual particles. All treatments were run in triplicate using separate preparations. The samples were prepared by suspending Mn02 in a 0.01 M acetate buffer of the desired pH with metals added as determined by treatment. Constant ionic strength and adequate current flow were established by using a matrix of 0.1 M NaN0,. When a steady state was reached, the electrophoretic mobilities of the colloidal particles were measured. The samples were then analyzed as described earlier. The 0.01 M acetate buffer did not significantly alter measured steady-state values. The electrophoretic mobility was measured directly and can be related to the {potential by the Helmholtz-Smoluchowski equation * l =( 4 w 4 / € (2) where @{ is the zeta potential (mV), 7 is the viscosity of water, t is the dielectric constant of bulk water, and p is the electrophoretic mobility [ (pm s-l)/(V cm-l)]. Kinetic Studies. Reaction kinetics were evaluated using a stirred-flow procedure modified from that of Walker et al. (20). A syringe drive pump was replaced with a peristaltic pump and a flow cell colorimeter used in place of the UV-vis spectrophotometer. A reaction vessel with approximately 25 mL of 0.001 M NaN03 was maintained at constant pH with 0.01 M acetate buffer. The effluent solution from the reaction cell was mixed with a color development reagent, diluted to the proper range for measurement by a colorimeter, and read at 570 pm. Color development was as previously described for Cr(VI), but the chromophore was diluted by half to obtain the current concentration range for measurement. The influent solution contained a Cr(II1) concentration equal to that of the initial solution present in the reaction cell; thus, a constant inflow concentration of reacting solution was maintained. Each run consisted of bringing the solution reactants to the desired volume in the reaction cell. The system was then allowed to flow at a rate of 5 mL min-'. When a steady base line was read, the reaction was initiated by injecting 1mL (2.5 mg) of Mn02 suspension into the reaction cell.
s
Table I. AG,,, for the Oxidation of Cr(II1) at Highest and Lowest pH and Initial Cr(II1) Concentrations Used in This Study"
400-
3.
Y
3 !a
300-
0
El
PI
F 1
0
100 -
200
C
:: s
'r'
400
"
"
800
600
u Y
HCr04-eqb
pH
2.4 1.9 254 635
33.1 32.5 364 80.1
3 5 3 5
-
U
\
AG,,,,
kJ
-43.97 -53.56 -39.39 -57.17
a Values were calculated from equilibrium activities and AGO,, of eq 1. bThe values listed represent the initial concentrations. Values are reported as activities based on the Davies equation. Activities are based on concentration units of micromolar.
Table 11. Speciation of Solution Components Present in the Cr-Mn02 System as a Function of pH, Calculated by the Computer Program MINTEQAZ -8
80 -
60-
Cr(1IUe;
33.5 770
200-
"
Cr(III)in: pM
component
PH
% species
Mn(I1) Cr(II1)
3-5 3
100 Mnt2 92.3 Crt3 7.7 Cr(OH)+2 53.9 Crt3 45.1 Cr(OH)+2 8.9 Crt3 74.8 Cr(OH)+2 45.1 Cr(OH)2t 99.7 HCr0; 99.4 HCr0496.5 HCrOL 3.5 CrOA2-
\
4 5
C
40-
a"
Cr(V1)
3
4 5
PH Figure 1. Cr(II1) oxidized by 6-Mn0, (initial concentration of 0.1 g L-I): (a) as a function of added Cr(II1) concentration at various pH values, and (b) as a function of pH at various Cr(II1) concentrations.
Results and Discussion Steady-State Cr-MnOz Systems. The 6-MnOZused in this study contained 48% Mn by weight with an [O]/ [Mn] > 1.9; thus, based on a 1.51 stoichiometry, 2.5 mg 6-Mn02should be capable of oxidizing approximately 14 pmol of Cr(II1). At a volume of 25 mL, this corresponds to a concentration of 460 pM. Figure 1 illustrates the Cr(II1) oxidation with varying initial Cr(II1) concentrations over the pH range of 3-5 in a matrix of 0.1 M NaN0,. At pH values less than 3.5, Cr(V1) production produces the maximum potential oxidizing capacity of the MnO,; however, as the solution pH and Cr(II1) concentration was increased, the extent of oxidation progressively decreased (Figure lb). At pH 1 4 and initial [Cr(III)] > 77 pM, reactants remained upon reaching a steady state. The reaction of Cr(II1) and 6-Mn02 clearly does not go to completion as pH and Cr(II1) concentrations are increased. Although a positive AGO, was calculated from eq 1for standard conditions, the formation of Cr(V1) indicates the reaction proceeds under initial conditions. Using a form of the Lewis equation AG,,, = AGO,, + RT In Q (3) the AG,, can be calculated (21),where Q is the reaction quotient and RT have their usual meanings. Table I lists the AG,, calculated for steady-state values generated by the addition of 33.5 or 770 pM Cr(II1) at pH 3 and 5. Equilibrium activities were calculated by the chemical equilibrium program and associated thermodynamic data Z using the Davies equation to calculate base M I ~ (32), activity coefficients. As illustrated in Table I, the reaction is favorable at the steady state for the conditions employed in this study. From a thermodynamic standpoint, the reaction should proceed to a greater extent at higher pH
values and at higher reactant concentrations. This was not observed in the oxidation process under experimental conditions of this study; therefore, assuming the reaction stoichiometry of eq 1is correct, the limitation in Cr(II1) oxidation must not be thermodynamically related. Except for Cr(III), the speciation of solution components present in the Cr-MnOz system was not altered greatly over the pH range studied. Table I1 gives the solution speciation, in percentage, according to the chemical equilibrium program MINTEQAP. The Cr(II1) fraction was dramatically affected in the pH range studied and may affect the extent of oxidation. However, the Cr(II1) species are in rapid equilibrium with each other, which allows even a minor species to react and be replinished for further reaction. At pH = 5, the largest fraction of solution Cr(II1) is present as the first hydrolysis product, &OH2+,and thus reaction 1may not be descriptive of the oxidation reaction. A possible reaction involving MnOz and &OH2+ (representative of pH = 5) would be CrOHz++ 1.56-Mn02 = HCr0,-
+ 1.5Mn2+
(4)
Thermodynamics calculations using this equation do not include H+ (although the species present are pH dependent) and thus would no longer predict a more spontaneous reaction with increasing pH. However, utilizing this reaction stoichiometry, thermodynamic calculations at pH = 5 still predict a favorable AG,,, after a steady state is reached (where reactants are not limiting). The bulk solution solubility of Cr(OH),(,, was approached as pH and Cr(II1) concentration were increased. At pH = 5, MINTpredicted that the Cr(II1) concentrations greater than 400 pM exceeded the solubility of &(OH),. If AG for the reaction of solid Cr(OH),(,, and MnOz is calculated, a positive value occurs at pH values greater than 3. This formulation 2H+ + Cr(OH), + 1.56-MnoZ= HCr0,-
+ 1.5Mnz++ 2H20 (5)
Environ. Sci. Technol., Vol. 26, No. 1, 1992
81
0 pH = 3.0 pH=3.5
y = - 0.25
+
y = - 37
1 . 7 6 ~ R”2 = 0.99
+
1 . 5 6 ~ R”2 = 0.98
5
5
15
25
35
45
55
Time [minutes] 0
100
200
300
500
400
Cr(VI) Produced [pMJ Figure 2. Stoichiometry of Mn(I1) formed for each Cr(V1) produced in the oxidation of Cr(II1) by 6-Mn0,.
_L
0.0
3.0
3.4
3.8
4.2
4.6
5.0
PH Figure 3. Sorption of Cr on 6-Mn0, in a matrix of 0.1 M NaNO,.
suggests an inverse relationship with pH and predicts the reaction to be thermodynamically less favorable with increasing pH. Therefore, the solution speciation does not appear to thermodynamically limit Cr(II1) oxidation unless the solid-phase Cr(OH), is formed. The stoichiometry of the reaction proposed by Amacher and Baker (6) indicated that 1.5Mn(II) should form for each Cr(V1) produced. Data were fit by linear regression, and the slope approximates the formation ratio of Mn(I1) to Cr(V1). With an [O]/[Mn] = 1.9 for the oxide, the theoretical slope would be equal to 1.66 (O/Mn less than 2 occurs from Mn(II), Mn(III), or a vacancy in place of Mn(V1)). The relationship of Cr(V1) to Mn(I1) generated was found to be in relatively good agreement with the predicted stoichiometry (Figure 2). Deviation from the theoretical value can be explained by experimental error and the oxide preparation procedure. Chromium sorption increased with increased pH and Cr(II1) concentrations (Figure 3). Chromium not found in solution was assumed to be sorbed by the oxide, but the form of the Cr species bound to the surface cannot be identified by the methods employed in this study. However, as shown in Figure 3, less than 1.5% of the added Cr(V1) was sorbed by unreacted 6-Mn02over the pH range of 3-5. In addition, it would be hypothesized that more anionic Cr(V1) would bind at lower pH values, yet the loss of Cr from solution increased with pH. This leads to the 82
Environ. Sci. Technol., Vol. 26, No. 1, 1992
Figure 4. Cumulative Cr(V1) formed over time by the reaction of Cr(II1) with 6-MnO, (0.1 g L-‘) in a stirred-flow reaction chamber.
conclusion that unless a surface alteration occurred during the oxidation process, most of the sorbed Cr was Cr(II1). In order to investigate the possibility of product formation inhibiting the oxidation process, Fendorf et d. (22) individually added Cr(V1) or Mn(I1) to MnOz suspensions prior to the addition of Cr(II1). Neither competing ion was found to significantly influence the extent of Cr(II1) oxidation. Therefore, on the basis of sorption data and direct competition studies, it is unlikely that the reaction products inhibit Cr(II1) oxidation. Rates of Cr(V1) Formation. A stirred-flow cell allows reactants to be mixed, significantly minimizing transport phenomena, while reactant concentrations are maintained and products removed (23). In this study, a kinetic method was used to supplement information on Cr(II1) oxidation derived from the static studies. Figure 4 depicts cumulative Cr(V1) produced during the oxidation reaction as a function of time. With the stirred-flow apparatus, the oxidation rate at pH = 5 was too fast to accurately interpret rates. The rates of reaction at pH = 3 were slower, but continued for prolonged times relative to pH = 5. Oxidation was extremely rapid at pH = 5 regardless of the Cr(II1) concentration. After an initial rapid oxidation, higher Cr(II1) concentrations caused a more extensive decrease in the oxidation rate at this pH. The rate of reaction under constant solution reactant influent at pH = 3 is dependent on the amount of surface consumed by the reaction. As the MnOz is consumed, the rate decreases, with the reaction continuing until the surface is diminished-the reaction rate was dependent on the level of Mn02. Very rapid initial rates of oxidation at pH = 5 are followed by a “shut down” in oxidation. Because products are removed in this flow cell, it is again apparent that at pH = 5 solution products do not limit the oxidation reaction. Thus, these rate studies provide additional information, suggesting that the reactant concentrations and their interactions with the surface are controlling oxidation. Electrokinetics. Electrophoretic mobilities of colloidal particles allows the determination of the electric potential at the shear plane (the { potential, *{). The { potential is related only to the mobile part of the electric double layer and gives no definitive information about surface potential, ion plane potentials, or charge densities. Yet, realistic assumptions allow surface potential and charge density to be calculated. The binding plane must be located at or inside the plane of shear, because ions could not be bound to the surface if they moved without the
-
-< E >
eE
sx
-5
4 '
0 0
Cr(WdpM Cr(JII)=38.5pM
A
Cr(III)=77.0pM
0
Cr(lII)=227pM
'A 2-
.
*
Cr(UI)=377pM Cr(Ll1)=596p~
I -
E .-
1-
.
0
c)
o!0
0.
c
P,
+
-Wu -'; b)
-2
Figure 6. Chromlum(V1) produced by the reaction of Cr(II1) with 6-Mn0, at pH = 3 and 5 in 0.001 and 0.1 M NaNO,.
Ionic Strength Effects. Indifferent electrolytes should not greatly alter the reactions under investigation, but ionic strength will affect activity coefficients and surface properties. When reactions involve a solid surface, the effects on surface charge and competitive adsorption from the counterions also arise. All experiments in this study were conducted at pH levels above the ZPC (pH > pHPC); thus, the surface maintained a negative charge which increased with pH. At constant pH, higher ionic strengths will increase surface charge. The effect of changing ionic strength, via different NaNO, concentrations, on the oxidation process is shown in Figure 6. At pH 3, the electrolyte concentration had very little effect on the oxidation process. Since the solution pH is near the ZPC, ionic strength should have little effect on the surface charge, yet increased ionic strength should still lower the activity of solution species. At pH 5, increased electrolyte concentration substantially increased Cr(II1) oxidation. Thus, although increases in ionic strength and pH both increase surface charge, these two parameters had opposing effects on Cr(II1) oxidation. Increases in pH resulted in a limitation in the oxidation process, while greater ionic strengths enhanced oxidation. Mechanisms of Cr(II1) Oxidation Inhibition. A limitation in Cr(II1) oxidation with increased pH and Cr(II1) concentration was observed. Reaction products do not appear to inhibit the reaction; thus, a surface alteration may explain the oxidation inhibition. The influence of Cr(II1) on the electrophoretic mobility of 6-Mn02appears to be related to the decrease in rate and extent of oxidation. In the Cr-MnO, system, only Cr(II1) or Mn(I1) would be capable of making the electrophoretic mobility more positive; however, the electrophoretic mobility of MnO,-Mn(II) is not similar to the Cr-MnO, system (22). Therefore, the electrophoretic mobility of the oxide must be influenced by Cr(II1). At pH values greater than 3.5, 6-Mn0, must oxidize Cr(II1) as well as sorbing it in a manner that does not allow the oxidation reaction to occur. At pH = 5, when Cr(II1) levels exceeded 77 pM the amount of Cr(V1) produced becomes relatively invariant to increases in added Cr(II1) (Figure 1). This suggests that a finite amount of Cr is oxidized prior to a surface alteration, which subsequently inhibits further oxidation. The speciation and increase in saturation indexes (SI = IAP/K,,) of Cr(II1) also correspond to the change in electrophoretic mobility and decrease in oxidation. It may be possible that Cr(II1) adsorbs to the surface without undergoing subsequent oxidation because of either site or Environ. Sci. Technol., Vol. 26, No. 1, 1992
83
ion characteristics influenced by pH. In addition, these sites would have to be strategically located, allowing the bound Cr(II1) to inhibit solution Cr(II1) from reaching surface sites capable of oxidation. These restrictions seem a bit unreasonable for this mechanism to apply. Surface precipitation rather than adsorption may explain the codependency of the oxidation limitation on Cr(II1) concentration and pH, and explain the effects of ionic strength. James and Healy (26) provided several explanations for electrophoretic mobility and adsorption isotherms of hydrolyzable ions on Ti02: (i) adsorption of the free aquo cation or the first hydrolysis product, (ii) surface binding of a polynuclear complex of discrete composition, or (iii) a surface polymer or precipitate of metal hydroxide formed or nucleated by the substrate. These authors dismissed these postulates and developed a model which predicted surface precipitation prior to bulk solution precipitation to be the most plausible explanation for sorption of hydrolyzable ions to oxide surfaces. Surface precipitation rather than adsorption has explained the sorption behavior of various metals to oxide surfaces (27-30). Tewari and Lee (30)used X-ray photoelectron spectroscopy (XPS), adsorption, and electrophoretic mobility data to show that Co formed a surface precipitate of CO(OH)~ on various oxides under conditions that did not cause precipitation in bulk solution. Furthermore, CO(OH)~ has been observed to form on 6-Mn02 when critical concentrations and pH were surpassed (28). However, extended X-ray absorption fine structure spectroscopy (EXAFS) has indicated that at 10% surface coverage Pb and Co formed multinuclear species on oxide surfaces (33). The multinuclear metal species observed at lower surface coverages may be a precursor to the formation of a surface precipitate. The mass action balance proposed by Bleam and McBride (34)to explain the progression from isolated-site binding to the formation of multinuclear sorbed clusters with increases in solution metal concentration may explain the formation of a surface precipitate at the even greater solution metaI ion concentrations. For the oxidation of low molecular weight organic compounds by MnO,, it is hypothesized that the organic molecule forms a surface complex prior to electron transfer (35,36). Hence, an initial oxidation of Cr(II1) would be terminated by subsequent precipitation of chromium hydroxide, and as previously calculated, at pH values greater than 3, oxidation of Cr(OH),,,, by MnOz would not be thermodynamically favorable. The hydroxide precipitate would thus not undergo oxidation by the Mn02 and would alter the surface properties of the oxide, as observed in the electrokineticstudies. Moreover, a finite amount of Cr(II1) would be oxidized prior to the formation of the surface precipitate, and after formation of the precipitate, oxidation would be inhibited. The primary effect of ionic strength would be to lower the activity of solution species; consequently, a lower saturation index would result in slowing precipitation and would allow oxidation to proceed to a greater extent before precipitation occurred. Conclusions
Chromium oxidation by 6-Mn02occurred over a range of pH values and Cr(II1) concentrations. The reaction stoichiometry of Mn(I1) to Cr(V1) appears to be 1.51. The reaction was limited as pH and Cr(II1) concentration increase; reactants remain in the system at the steady state. However, thermodynamic calculations indicate that the reaction should continue until reactants are limiting. Products of the oxidation-reduction reaction between Cr and Mn02 do not limit the oxidation process. It appears that Cr(II1) concentration, pH, and the amount of initial 84
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available surface are the parameters controlling the degree of oxidation. Cr(II1) caused the electrophoretic mobility of 6-Mn02 to become less negative and induced a charge reversal. The speciation and increase in saturation indexes of Cr(II1) followed trends exhibited by the electrophoretic mobility and decreases in oxidation. While increases in both pH and ionic strength increased surface charge, these two variables had opposing effects on the oxidation process. A surface alteration inhibiting further Cr oxidation seems a plausible explanation for the observed phenomena. Cr(II1) sorption to the oxide surface dominates the surface properties of the MnO, at higher pH values and inhibits the redox reaction. Possibly, multinuclear Cr(II1) complexes begin to form with a progression to complete surface precipitation at higher solution pH and Cr(II1) concentrations. The increase in electrophoretic mobility and decrease in oxidation would thus be due to the onset of multinucleated species and finally the formation of a surface precipitate. Therefore, surface precipitation may explain the observations of this study. Registry No. MnOz, 1313-13-9;Cr, 7440-47-3;Mn, 7439-96-5.
Literature Cited U.S. Environmental Protection Agency. Municipal sludge management: enuironmental factors; NTIS/EPA 430/977-004; Office of Water Program Operations: Washington, DC, 1977; p 19. National Research Council, Committee on Biologic Effects of Atmospheric Pollutants. Chromium. National Academy of Science: Washington, DC, 1974. Eary, L. E.; Rai, D. Enuiron. Sci. Technol. 1987, 21, 1187-1193. Bartlett, R. J.; Kimble, J. M. J. Enuiron. Qual. 1976, 5, 379-383. Bartlett, R. J.; James, B. J. Enuiron. Qual. 1979,8, 31-35. Amacher, M. C.; Baker, D. E. DE-AS08-77DP04515, Institute for Research on Land and Water Research, Pennsylvania State University, University Park, PA, 1982. Garrels, R. M.; Christ, C. L. Solutions, minerals, and equilibria;Freeman, Cooper, and Co.: San Francisco, CA, 1965; p 450. Morgan, J. J.; Stumm, W. J. Colloid Sci. 1964,19,347-359. Posselt, H. S.; Anderson, F. J.; Weber, W. J. Environ. Sci. Technol. 1968,2, 1087-1093. Murray, D. J.; Healy, T. W.; Fuerstenau, D. W. Adu. Chem. Ser. 1968, No. 79, 74-81. Loganathan, P.; Burau, R. G. Geochim. Cosmochim.Acta 1973, 37, 1277-1293. Zasoski, R. J.; Burau, R. G. Soil Sci. SOC.Am. J . 1988,52, 81-87. Buser, W.; Graf,P.; Feitknecht, W. Helu. Chim. Acta 1954, 37, 2322-2333. Heilman, M. C.; Carter, D. L.; Gonzalez, C. L. Soil Sci. 1965, 100,409-413. Blok, C.; De Brun, J. E. J . Colloid Interface Sci. 1970,32, 518-525. Whittig, L. D.; &dice, W. R. In Methods of soil analysis: Part 1. Physical and mineralogical methods, 2nd ed.; Klute, A., Ed.; 1986; Vol. 9, Chapter 12. Loganathan, P. Sorption of heavy metals on a hydrous manganese oxide. Ph.D. Thesis, University of California, Davis, 1971. Bangham, A. D.; Heard, D. H.; Flemans, R.; Seaman, G. V. Nature 1958, 182, 642-644. Akeson, M. A.; Munns, D. N.; Burau, R. G. Biochim. Biophys. Acta 1989, 986, 33-40. Walker, W. J.; Cronan, C. S.; Patterson, H. H. Geochim. Cosmochim. Acta 1988, 52, 55-62. Rock, P. A. Chemical thermodynamics;University Science Books: Mill Valley, CA, 1983; pp 372-373. Fendorf, S. E.; Zasoski, R. J.; Burau, R. G. Soil Sci. SOC. Am. J , submitted.
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Sparks, D. L. Kinetics of soil chemical processes; Academic Press: New York, 1989; p 47. Hansmann, D. D.; Anderson, M. A. Environ. Sci. Technol. 1985,19,544-551. Healy, T. W.; James, R. 0.;Cooper, R. Adv. Chem. Ser. 1968, NO. 79, 62-73. James, R. 0.;Healy, T. W. J. Colloid Interface Sci. 1972, 40, 53-65. James, R. 0.;Healy, T. W. J. Colloid Interface Sci. 1972, 41,6581. Crowthier, D. L.; Dillard, J. G.; Murray, J. W. Geochim. Cosmochim. Acta 1983,47, 1399-1403. Murray, J. W.; Dillard, J. G. Geochim. Cosmochim. Acta 1979, 43, 781-787. Tewari, P. H.; Lee, W. J. Colloid Interface Sci. 1975,52, 77-88.
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Received for review January 28, 1991. Revised manuscript received June 13, 1991. Accepted June 26, 1991.
Isotopic Composition of Sulfur in Mosses across Canada Jerome 0. Nrlagu" and Walter A. Glooschenko
National Water Research Institute, Box 5050, Burlington, Ontario L7R 4A6, Canada
rn The isotopic composition of sulfur in mosses is used as the basis for dividing the country into four source regions of atmospheric sulfur: (i) the Western region dominated by seasalt spray sulfur, (2) the Prairie region which is strongly influenced by sulfur from evaporitic dust particles, (3) Ontario and Quebec region where the sulfur is derived mostly from industrial sources, and (4) the Atlantic region with a mixture of sulfur from Atlantic seasalt sprays and the polluted air masses from continental North America. The major point sources such the smelters at Sudbury (Ontario), Rouyn-Noranda (Quebec), Thompson (Manitoba), and Flin Flon (Saskatchewan) can also be identified on the basis of their isotopic signature on the moss sulfur. The data suggest that the stable isotopic composition of sulfur in plants can be used as tracers in the study of regional and transboundary movement of pollutant sulfur.
Introduction Cryptogams (especially mosses and lichens) are very sensitive to sulfur oxides (SO,) and hence have been used in biomonitoring the degree of air pollution on both local and regional scales (1-5). Such plants are able to provide an integrated record of the intensity of air pollution because (a) they derive most of their mineral requirements along with the contaminants from the atmosphere and (b) they are efficient traps that can bioaccumulate many elements to levels well above their metabolic needs (1,6-8). In general, the relationship between the ambient concentration or deposition of SO, and the sulfur content of the moss or lichen is not straightforward. The accumulation of sulfur by any moss or lichen species can also be influenced by the duration and episodicity of sulfur fumigation, humidity and other weather conditions, rates and processes of wet/* deposition, rate and robustness of plant growth, general chemodynamics of the particular plant species, etc. These variables, by contrast, do not always induce a significant fractionation of sulfur isotopes, and the isotopic composition of sulfur in mosses should be similar to that of atmospheric sulfur oxides (9-12). Measurement of the isotopic composition of sulfur in mosses should thus represent a unique and powerful tool which has yet to be fully exploited in the study of sulfur pollution on a regional scale (13, 14). 0013-936X/92/0926-0085$03.00/0
This report presents the concentrations and isotopic composition of sulfur in mosses from bog ecosystems across Canada, with detailed sampling around two major SO, point sources (base metal smelters) at Sudbury, ON, and Rouyn-Noranda, PQ. The objectives of the study are 2-fold: (a) to demonstrate that the isotopic signature of sulfur in ambient air is similar to that in the moss samples and (b) to define any regional differences in the sources and intensity of sulfur pollution across the country.
Methodology The moss species, Sphagnum fuscum, is particularly suitable for this study. It grows on low hummocks in ombrotrophic bogs which receive all their nutrients and contaminant burdens from the atmosphere. It occurs throughout Canada and western Europe. The numerous studies using this species in biomonitoring the atmospheric deposition of metals and other contaminants show that it is an excellent accumulator of air pollutants (7). The Sphagnum spp. specimens were collected from bogs located along a transect running approximately northwest from Sudbury (7,15). In the Rouyn-Noranda region, all the samples were obtained within a 30-km radius of the copper smelter. Around these two major sources of pollutant sulfur in Canada, the occurrence of s.fuscum tends to be patchy, and samples of Sphagnum capillaceum (another moss species) and Chamaedaphne calyculata (leatherleaf)were also obtained to compare the interspecies differences in the sulfur isotope composition as well. The sampling of mosses from bogs across Canada (Figure 1) was not based on any particular grid but was determined by the occurrence of S. fuscum especially in boreal forest areas with no obvious local source of sulfur pollution. At each locality, at least four aliquot samples of the living or recently dead mosses were collected randomly from hummocks in the bog and then composited to make a volume of approximately 2.0 L. Any areas under trees or shrubs were avoided, and the collection was restricted to a depth of less than 5.0 cm. Unwanted plants and twigs were removed. About 100-200 g of the shrub leaves were also collected in the Sudbury and Rouyn-Noranda areas. Samples were kept cool in the field and stored in sealed containers upon return to the laboratory. Subsequently, the samples were dried in an oven at a temperature of 90
0 1991 American Chemical Society
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