Environ. Sci. Technol. 2010, 44, 3895–3900
Reductive Dissolution of Lead Dioxide (PbO2) in Acidic Bromide Solution Y I - P I N L I N * ,† A N D RICHARD L VALENTINE‡ Division of Environmental Science and Engineering, Faculty of Engineering, National University of Singapore, Singapore 117576, and Department of Civil and Environmental Engineering, University of Iowa, Iowa City, Iowa 52242-1527
Received January 13, 2010. Revised manuscript received March 15, 2010. Accepted April 6, 2010.
Reductive dissolution of lead dioxide (PbO2(s)) has been attributed to a major route leading to elevated lead concentrations in drinking water. However, surface processes involved in this heterogeneous reaction have not been elucidated. In this study, the kinetics and mechanism of reductive dissolution of PbO2(s) in acidic bromide solutions were investigated to reveal the detailed surface reactions. The reduction of PbO2 by bromide can be expressed as PbO2(s) + 2Br- + 4H+ f Pb2+ + Br2 + 2H2O. The reaction kinetics was found to be proportional to the concentration of PbO2(s), and the reaction orders were 1.08 and 1.77 with respect to bromide and proton concentration, respectively. The observed kinetic data can be explained by the following reaction mechanism: adsorption of bromide on the PbO2(s) surface to form a precursor surface complex ≡PbIVBr, and two separate one-electron transfers from adsorbed Br- to structural PbIV, followed by the release of Pb2+ and Br2 into water. The adsorption of bromide ion on the PbO2 surface and the first one-electron transfer reaction were found to be important in regulating the overall rate of the reaction. It is expected that a similar reaction scheme can be applied to other reductive ions.
Introduction PbO2(s) is a strong oxidant possessing two different polymorphs including orthorhombic scrutinyite (R-PbO2(s)) and tetragonal plattnerite (β-PbO2(s)). Figure 1 shows the pe-pH diagram for lead (1). In neutral to acidic pH values (pH < 7.2), the oxidation potential of PbO2(s) is higher than that of the H2O/O2 boundary suggesting that the reduction of PbO2(s) by water is thermodynamically possible. In nature, PbO2(s) is found in the oxidized zone of hydrothermal lead-bearing ores (2). Recently, PbO2(s) was identified in drinking water distribution systems. Its presence has been attributed to the oxidation of leadcontaining plumbing material by free chlorine (3-5). Its stability in aqueous solution has drawn increasing interest because of its unique role in regulating lead concentration in drinking water. The solubility of PbO2(s) is extremely low in water if the redox potential is maintained at a highly * Corresponding author phone:
[email protected]. † National University of Singapore. ‡ University of Iowa. 10.1021/es100133n
(65)
6516-4729;
2010 American Chemical Society
Published on Web 04/26/2010
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oxidative level which favors its existence such as in the presence of free chlorine (6, 7). However, under the right conditions, it is subject to reductive dissolution which results in the release of soluble Pb(II) into water (8-14). For example, alarmingly high and potentially toxic levels of lead were recently observed in certain parts of the drinking water distribution system in Washington, DC (6, 15), which have resulted in elevated blood lead levels in children (16). These high levels of lead in drinking water were attributed to the likely dissolution of accumulated PbO2(s) or increased rate of corrosion of elemental lead bearing materials caused by a change of disinfectant from free chlorine to monochloramine that produced a more reductive environment which inhibited the formation of lead dioxide (6, 15, 17, 18). As a consequence of this incident, there has been a renewed interest in better understanding the processes that govern the stability of PbO2 in drinking water distribution systems. In the past, the reactions of PbO2(s) that have been studied are mostly pertinent to the chemistry of electrochemical cells due to its extensive use in lead acid batteries (19-24). The reactivity of PbO2(s) in aqueous solution has been much less studied. PbO2(s) is subject to reductive dissolution involving a number of species found in drinking water, but the whole picture is not yet complete enough to allow estimation of the overall importance of these reactions. In this study, bromide, which is occurring in many drinking water sources, is used as a redox-active probe to investigate the reduction of PbO2(s) in an acidic environment (pH 3.5-4.5). The use of acidic solutions was based on our preliminary studies showing that the reaction proceeds at a rate that allows its accurate quantification. The kinetics and mechanism of PbO2(s) reduction by bromide were examined to explore the surface processes involved in this heterogeneous redox reaction. The role of this reaction in increasing soluble lead levels in drinking water is also evaluated.
Materials and Methods Chemicals. Reagent grade chemicals were used in this study. All solutions were prepared by deionized water obtained from a Barnstead ULTRO pure water system (Barnstead-Thermolyne Corp.). Commercially available PbO2(s) was used upon received without further purification (Fisher Scientific). The specific surface area determined by the 7-point N2-BET method is 4.1 m2/g. Images acquired by the scanning electron microscopy (SEM) indicated that the size of each individual particle ranges from 0.1 to 0.5 µm but they tend to agglomerate to form bigger aggregates as shown in Figure 2. X-ray diffraction analysis (XRD) showed that this PbO2(s) is plattnerite (β-PbO2). Potassium bromide (Fisher Scientific) was used for the source of bromide. Kinetic Experiments. Kinetic experiments were conducted using 250 mL polypropylene bottles at 25 °C. The effects of PbO2(s) concentration (12-48 mg/L), bromide concentration (0-5 mM), and acidic solution pH (3.5-4.5) on the rate of PbO2(s) reduction by bromide were determined. All solutions were buffered by 10 mM acetate. The solution pH value was adjusted by 1 N NaOH and HCl. The solution was mixed by a magnetic star bar at 400 rpm during the course of the experiment. The concentration of Pb2+ was measured for up to 74 h. The initial rate of PbO2(s) reduction was derived from the initial slope of the Pb2+ concentration versus time curve. Duplicate experiments were conducted in selected conditions. The data points shown in the duplicate experiments are the average of the duplicates, and the error VOL. 44, NO. 10, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 1. pe-pH diagram of lead. The following conditions are considered for the construction of the figure: total inorganic carbon concentration ) 1 mM, total soluble Pb concentration ) 10-6 M, ionic strength ∼ 0, T ) 25 °C. Thermodynamic data were obtained from Pankow (1).
FIGURE 3. Effects of PbO2(s) concentration on the reduction of PbO2(s) by bromide. (a) Pb2+ concentration as a function of time at different PbO2(s) concentration. (b) Initial Pb2+ formation rate as a function of PbO2(s) concentration. pH ) 4.0, Br- ) 2.5 mM, acetate buffer ) 10 mM, and T )25 °C. background absorbance considering our experimental conditions. Thus, only Br3- was used qualitatively as an indicator for the bromide oxidation Analytical Methods. A Nano-Band Explorer II anodic stripping voltammetry instrument (TraceDetect, Seattle) was used for measuring Pb2+ concentration resulting from the reduction of PbO2(s) by bromide. The presence of PbO2(s) particles does not cause interference in this method. For each measurement, 0.1 M acetate was added to serve as the background electrolyte and a standard addition approach was employed. The UV-vis absorbance spectrum was acquired by a UV-vis spectrometer (UV1601, Shimadzu). SEM images were acquired by the Hitachi S-4000 SEM. A MiniFlex II desktop X-ray diffractometer (Rigaku Americas) was used to determine the XRD pattern of the PbO2(s) employed. The solution pH value was measured by a pH meter (A15, Fisher Scientific) coupled with a calomel combination pH electrode pre-equilibrated by pH 4 and 7 standard buffers. FIGURE 2. SEM images of the PbO2(s) employed in this study. The size of each individual particle ranges from 0.1 to 0.5 µm. The specific surface area is 4.1 m2/g.
bar represents the range of the measurements. Selected samples were collected for the UV-vis absorbance measurement to qualitatively show the oxidation of bromide. There are two oxidation products of bromide: Br2 and Br3- which possess molar absorptivity of 173 M-1 cm-1 at 390 nm and 35 000 M-1 cm-1 at 266 nm, respectively (25). Although Br2 was found to be the dominant species, its small molar absorptivity resulted in insignificant absorbance above 3896
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Results and Discussion Influence of PbO2(s) Concentration, Bromide Concentration, and pH Value on Reaction Kinetics. The influence of PbO2(s) concentration, bromide concentration, and pH on the rate of PbO2(s) reduction by bromide is shown in Figures 3-5, respectively. The reduction of PbO2(s) by bromide can be expressed by the following reaction: PbO2(s) + 2Br- + 4H+ f Pb2+ + Br2 + 2H2O
(1)
A small fraction of Br2 can react with excess Br- to form Br3- based on the following equilibrium reaction (25)
KBr
Br2 + Br- 798 Br-3 3
(2)
where KBr3 ) 16.6 M-1 (25). Considering the bromide concentration (25) employed in this study, the ratio of [Br3-]/ [Br2] ranges from 0.017 to 0.083. The rate of PbO2(s) reduction to form Pb2+ is proportional to the PbO2(s) concentration (Figure 3) from which a surface area normalized reaction rate can be determined. The rate of PbO2(s) reduction increases with increasing bromide concentration and decreasing solution pH. The apparent reaction orders with respect to bromide and proton concentration are 1.08 (Figure 4) and 1.77 (Figure 5), respectively. It should be noted that in control experiments without bromide, a small amount of Pb2+ was still formed (Figure 4a, Pb2+ ) 0.75 µM at the end of the experiment). This observation is consistent with the pe-pH diagram (Figure 1) and a previous study showing that water can serve as a reductant for PbO2(s) (9). However, the rate of this water-induced PbO2(s) reduction is insignificant compared to that resulting from bromide and does not contribute significantly to the overall rate of PbO2(s) reduction. Br3- possesses a peak UV absorbance at 266 nm with a molar absorbance of 35 000 cm-1 M-1 (25). The formation of Br3- is indicated by the increase of UV absorbance at 266 nm as a function of time for a selected experimental condition shown in Figure 6 (PbO2 ) 12 mg/L, Br- ) 2.5 mM, pH ) 4). Br3- can be slowly lost due to the evaporation of Br2 in to the air. Because of the long period (up to 74 h) and continuous stirring employed in our experiments that may enhance Br3- loss, its formation is only used for qualitative purpose to verify the oxidation of bromide by PbO2(s).
FIGURE 5. Effects of pH on the reduction of PbO2(s) by bromide. (a) Pb2+ concentration as a function of time at different pH values. The dashed line represents the maximum possible Pb2+ from added PbO2 concentration. (b) Pb2+ formation rate as a function of [H+]. Br- ) 2.5 mM, PbO2(s) ) 12 mg/L, acetate buffer ) 10 mM, T ) 25 °C.
FIGURE 6. Formation of Br3- indicated by the increase of UV 266 nm as a function of time. Br- ) 2.5 mM, pH ) 3.5, PbO2(s) ) 12 mg/L, acetate buffer ) 10 mM, T ) 25 °C. We initially assumed that the overall rate could be expressed as a single term R ) kobs(SA)[Br-]a[H+]b
FIGURE 4. Effects of bromide concentration on the reduction of PbO2(s) by bromide. (a) Pb2+ concentration as a function of time at different Br- concentration. (b) Pb2+ formation rate as a function of Br- concentration. pH ) 4.0, PbO2 ) 12 mg/L, acetate buffer ) 10 mM, and T ) 25 °C.
(3)
where R is the rate of PbO2(s) reduction (M/h), kobs is the observed rate constant, SA is the available surface area of PbO2(s) (m2), and a and b are apparent reaction orders with respect to bromide and proton concentrations, respectively. Using the experimental data from Figures 3-5, the following least-squares regression can be obtained: VOL. 44, NO. 10, 2010 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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R ) [3.93 × 106](SA)[Br-]1.08[H+]1.77 -1
-2
k1
≡PbIVOH + Br- + H+ a ≡ PbIVBr + H2O k-1
k2
≡PbIVBr 98 ≡ PbIIIBr k3
≡PbIIIBr + H2O 98 ≡ PbIIIOH + Br·+H+
≡PbIIIOH + Br- + H+ 98 ≡ PbIIIBr + H2O k5
k6
≡PbIIBr + H2O 98 ≡ PbIIOH + Br·+H+
(6)
(7)
(8)
(9)
(10)
(3) Br · radicals dimerize to form Br2 k7
2Br· 98 Br2
(11)
and Br2 can further react with aqueous Br- to form Br3according to eq 2. (4) Release of Pb2+ ion into the solution and exposure of the underneath structural Pb(IV) forms a new surface site (below the surface structural Pb(IV) in the reactant side is not shown): 3898
9
(12)
On the basis of the proposed reaction scheme and with the assumption that a steady state is achieved for each of the surface species, the following rate expression for the rate of PbO2(s) reduction or Pb2+ formation can be obtained R)
k1k2 {≡PbIVOH}[Br-][H+] k2 + k-1
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(13)
where the braces denote surface concentration. The rate expression indicates that the adsorption of bromide ion on the PbO2(s) surface and the first one-electron transfer reaction are both important in regulating the overall rate of the reaction. It should be noted that {≡PbIVOH} is pH-dependent and can be expressed in terms of the total surface PbIV concentration according to the following acid-base reaction: Ksurf
≡PbIVOH 798 ≡ PbIVO- + H+
{≡PbIVOH} )
{≡PbIV}tot[H+] Ksurf + [H+]
(14)
(15)
By substituting eq 15 into eq 13, the rate equation can be expressed as R)
k1k2 (k2 + k-1)(Ksurf + [H+])
{≡Pb(IV)}tot[Br-][H+]2 (16)
(5)
(2) Two one-electron transfer reactions occur subsequently, each of which is followed by the release of the bromine radical (Br · ) into the solution:
≡PbIIIBr 98 ≡ PbIIBr
≡PbIIOH 98 ≡ PbIVOH + Pb2+
-1.85
The unit of the rate constant is h m M . The noninteger power to which hydrogen ion is raised indicates a potentially complex overall mechanism. Reaction Mechanism. Reductive dissolution of higher oxidation state metal oxides, such as manganese and iron oxide/hydroxide, have been shown to involve several reaction steps, including adsorption of the reductive ion on the oxide surface, electron transfer between the adsorbed ion and the structural metal, release of the reduced metal ion and the oxidized foreign ion from the surface, and possibly the formation of highly reactive intermediate resulting from the electron transfer (26-31). Inaone-electrontransferreaction,theLangmuir-Hinshelwood kinetic law has been successfully used to explain the redox kinetics if it is assumed that the adsorption of the reductive ion and the electron transfer are rate determining steps, the rate of back electron transfer is negligible, and the surface coverage of reductive ions on the oxide surface achieves a steady state (30). It is anticipated that the oxidation of PbO2(s) by bromide would follow a broad surface reaction scheme because of the involvement of PbO2(s) solid phase although two electrons are involved in the overall reaction. It is postulated that the two-electron transfer occurs through two separate one-electron transfer reactions because a one-step two-electron transfer reaction is not likely to happen. The following reaction steps considering Pb2+ and Br3- as the final products are proposed to describe the reduction of PbO2(s) by bromide: (1) Adsorption of bromide on the PbO2(s) surface forms a precursor surface complex (≡ denotes surface species):
k4
k8
(4)
Assuming that k2 is much greater than k-1 due to the fast electron transfer and slow desorption, eq 16 can be simplified as R)
k1 (Ksurf + [H+])
{≡Pb(IV)}tot[Br-][H+]2
(17)
This overall rate expression suggests that the reaction orders with respect to Br- and PbO2(s) concentration (or surface area) are both 1 which are consistent with the observed apparent reaction orders (eq 4). Equation 17 also indicates that the reaction order with respect to [H+] can range from 1 if Ksurf , [H+] to 2 if Ksurf . [H+]. When neither of these assumptions is true, the “apparent” order with respect to proton concentration will appear to be a noninteger between these two limits. The apparent reaction order with respect to proton concentration determined in our experiments is 1.77 which is in the suggested range. The oxidation of nitrite by monochloramine was empirically described by a similar single term rate expression with a noninteger dependence on proton concentration (32), which was later found to be consistent with a much more complicated mechanism over the relatively limited range in reaction conditions used in its derivation (33). It should be pointed out that, in our case, the apparent order of 1.77 with respect to proton concentration would likely change if a much larger pH range had been studied. By employing the nonlinear regression of our experimental data to eq 17 and using PbO2(s) surface area to represent PbO2(s) concentration, the values of k1 and Ksurf determined from the least-squares fitting were 9.61 × 103 h-1m-2 M-1 and 3.63 × 10-4 M, respectively. Equation 17 can subsequently be revised as
R)
9.61 × 103 (SA)[Br-][H+]2 (3.63 × 10-4 + [H+])
(18)
where (SA) is the total surface area of PbO2(s) particles. Figure 7 shows the relationship between the experimentally determined initial rates and those calculated using eq 18. The close correspondence indicates that the proposed mechanism provides a good explanation for the experimental observations. It should be noted that, in our previous study on the reductive dissolution of PbO2 by iodide, we observed reaction orders of 0.66 and 1.68 for proton and iodide, respectively, using phosphate-buffered solutions (12). The discrepancies of the reaction orders determined from that study from the proposed mechanism are due to the strong adsorption of phosphate on the PbO2 surface that alter the adsorption of iodide and subsequent electron transfer reactions. Environmental Implications. Kinetic studies at low pH showed that the reaction kinetics of the reduction of PbO2(s) by bromide can be described by a series of surface reactions including the formation of a surface complex, and two subsequent one-electron transfer reactions between adsorbed Br- and structural Pb(IV), followed by the release of Pb2+ and formation of Br2. Among these reactions, the adsorption of bromide ion on the PbO2(s) surface and the first one-electron transfer reaction were shown to be the rate-determining steps. An extrapolation of the rate expression to higher pH values suggests that the reaction is very slow at neutral and alkaline pH values. For example, at pH 7, PbO2 ) 5 g/L, and Br- ) 0.025 mM (2 mg/L), the rate of lead release is only 63 µg/L/year. This will not likely cause significant concerns on soluble lead contamination considering that even lower bromide concentrations (
