Environ. Sci. Technol. 2008, 42, 2919–2924
Reduction of Lead Oxide (PbO2) by Iodide and Formation of Iodoform in the PbO2/I-/NOM System Y I - P I N L I N , * ,† M I C H A E L P . W A S H B U R N , AND RICHARD L. VALENTINE* Department of Civil and Environmental Engineering, University of Iowa, Iowa City, Iowa 52242-1527
Received November 7, 2007. Revised manuscript received January 3, 2008. Accepted January 22, 2008.
Lead oxide (PbO2) can be an important form of lead mineral scale occurring in some water distribution systems. It is believed to be formed by the oxidation of lead-containing plumbing materials by free chlorine. Its reactivity in water, however, has not been well studied. Iodide is also found in source drinking waters, albeit at low concentrations. Consideration of thermodynamics suggests that iodide can be oxidized by PbO2. In this investigation, iodide ion was used as a probe compound to study the reduction of PbO2 and the formation of iodoform, which has been predicted to be a carcinogen, in the presence of natural organic matter (NOM). The reduction of PbO2 by iodide can be expressed as PbO2 + 3I- + 4H+ f Pb2+ + I3- + 2H2O, and the reaction kinetics has been determined in this study. In the presence of NOM, I3- reacts with NOM to form iodoform and its concentration is proportional to the NOM concentration. Our results indicate that PbO2 is a very powerful oxidant and can possibly serve as an oxidant reservoir for the formation of iodinated disinfection byproduct through a novel reaction pathway.
Introduction Lead is a toxic metal. Although the major source of human exposure to lead comes from lead-contaminated urban soils and lead-based paint (1–3), release of lead from lead-bearing plumbing materials in drinking water poses another great risk to public health. To control lead in drinking water, the U.S. EPA proposed an action level of 15 ppb in the 1991 Lead and Copper Rule (LCR) (4), and a revision of this rule was proposed in 2006 (5). In the past, it was generally thought that soluble lead in drinking water was regulated by the solubility of Pb(II) solid phases, such as cerrusite (PbCO3) and hydrocerrusite (Pb3(CO3)2(OH)2) (6–9). Most lead control strategies were proposed based upon their solubility (8, 10) or the use of phosphate to form a protective passivating film on leadbearing materials (11). However, recent studies suggest that PbO2, a Pb(IV) solid phase, may exist in some systems with lead service lines or lead-containing plumbing materials, especially those that use free chlorine as the disinfectant (11–14). PbO2 is almost insoluble and this property may * Address correspondence to either author. Phone: (65) 65164729(Y.P.L.), (319) 335-5653(R.L.V.); e-mail:
[email protected] (Y.P.L.),
[email protected] (R.L.V.). † Current address: Division of Environmental Science and Engineering, National University of Singapore. 10.1021/es702797b CCC: $40.75
Published on Web 03/06/2008
2008 American Chemical Society
attribute to the lower soluble lead concentration than that predicted from the solubility of Pb(II) carbonate solid phases (14). Thermodynamically, PbO2 is a strong oxidant and may only be stable in solutions with high oxidation potentials such as water containing free chlorine (14). A reduction of the solution oxidation potential could result in the release of soluble lead from PbO2 due to reductive dissolution. Switching disinfectant from free chlorine to chloramines, a relatively weaker oxidant than free chlorine, was blamed as one of the major factors for the high concentration of lead observed (4800 µg/L) in the Washington, DC drinking water in 2003 (12, 13). Supporting the conjecture of the importance of the reductive dissolution of PbO2 is the recent observation that PbO2 is stable in chlorinated water but not in chloraminated water (12, 13), suggesting that PbO2 is a weaker oxidant than free chlorine but may be a stronger one than chloramines. However, due to the lack of published literature, the reactivity of PbO2 in water is not certain and research is needed to fill the knowledge gap. Thermodynamic considerations indicate that iodide may be oxidized by lead oxide (Table 1) and possibly lead to the formation of iodated disinfection byproduct in the presence of NOM. In this research, iodide is used as a probe reductant to investigate the reactivity of PbO2 in water, as well as a reactant to form iodoform, which has been predicted to be a carcinogen (15) and identified in water that contains iodide and NOM and disinfected by free chlorine, monochloramine, or ozone (16–21). The objectives of this study were to (1) study the reactivity of PbO2 in water with iodide, along with the effects of buffer type on this process and (2) determine if oxidation of iodide by PbO2 in the presence of NOM can result in the formation of iodoform.
Material 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.). The PbO2 particles (Fisher Scientific) have a specific surface area of 4.14 m2/g as determined by the 7-point N2-BET method, and a size of 0.1–0.3 µm as determined by scanning electron microscopy. These particles tend to form aggregates as large as >5 µm. X-ray diffraction analysis indicated that this PbO2 is plattnerite, a tetragonal polymorph of PbO2. Plattnerite and scrutinyite, a orthorhombic polymorph of PbO2, can both be formed by the oxidation of Pb2+ by free chlorine (14) and present on the inside of pipes in the distribution system (23). A phosphate buffer stock (0.283M, pH 6.8) was prepared with NaH2PO4 and K2HPO4 (Fisher Scientific). NaHCO3 (Fisher Scientific) was used in the carbonate-buffered experiments. Potassium iodide (Fisher Scientific) was used for the source of iodide. NOM Sample. The Iowa River reverse osmosis (RO) concentrate was used as the NOM source for iodoform formation experiments. A RealSoft PROS/2S reverse osmosis unit (Stone Mountain, GA) was used to process 700 L of Iowa River water to produce a 40-L concentrate. Before RO concentration, the Iowa River water had a DOC and SUVA254 of 3.4 mg/L and 2.3 L/mg-m, respectively. In the RO concentrate, the DOC and SUVA254 were 50.5 mg/L and 2.21 L/mg-m, respectively. This RO concentrate has been shown to exhibit the same reactivity in terms of monochloramine consumption and NDMA formation kinetics when normalized to the DOC content of the original water (24, 25). Kitis et al. (26) also observed that the reactivity of NOM with respect VOL. 42, NO. 8, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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TABLE 1. Thermodynamics of PbO2 Reduction and IOxidationa EH° (volt) 4H+
2e-
PbO2 reduction: PbO2 + + Iodide oxidation: 3I- ) I3- + 2e-
)
Pb2+
+ 2H2O
1.45 -0.62
a Calculations were based on thermodynamic data from ref (22).
FIGURE 2. Effects of phosphate buffer concentration on the reduction of PbO2 by iodide. I3- concentration was measured as a function of time at different phosphate buffer concentrations. pH ) 7.0, I- ) 1.2 mM, PbO2 ) 8.0 mg/L, temp ) 25 °C, IS NaCl ) 0.1 M. The initial I3- formation rates with 95% confidence level in the presence of different phosphate buffer concentrations are shown in the inset.
FIGURE 1. Typical UV absorbance spectrum generated from the kinetic experiment. Experimental condition: pH ) 7.0, I- ) 1.2 mM, PbO2 ) 8 mg/L, phosphate buffer ) 10 mM, temp ) 25 °C, IS NaCl ) 0.1 M, reaction time 8 min. to formation of selected halogenated DPBs is also conserved when concentrated by this method. Kinetic Experiments of Iodide Oxidation by PbO2. Kinetic experiments were conducted using 125-mL amber bottles at 25 °C. The effects of phosphate buffer concentration, PbO2 concentration (or surface area), iodide concentration, and pH on the kinetics were investigated. Additional kinetic experiments with 10 mM NaHCO3 and without buffer addition were conducted at different pH values to determine the effects of buffer type on the oxidation of iodide by PbO2. Detailed experimental conditions for this kinetic study are presented in Table S1 in the Supporting Information. Ionic strength of these solutions was adjusted by adding 0.1 M NaCl, and pH value was adjusted by 1 N NaOH and HCl. The solution was mixed by a magnetic stir bar at 400 rpm during the course of the experiment. Based on preliminary experiments, the rate of I3- formation gradually decreased after 3 min of reaction, probably due to the consumption of most reactive sites on the PbO2 surface and a possible change of the PbO2 surface area. Thus, the reaction kinetics were tracked by measuring the concentration of I3- for 8 min after the reaction was commenced and the initial I3- formation rate was derived from the initial slope of the I3- concentration vs time curve. Iodoform Formation Experiments. Experiments were conducted using 125-mL amber glass bottles at 25 °C. Each sample, which was buffered by 2 mM NaHCO3, contained 100 mg/L PbO2 and 10 mg/L potassium iodide. NOM concentration of 0, 1, 3, and 5 mg/L as C were added from the Iowa River RO concentrate. The pH value of each sample was adjusted to 7.0 by 1 N NaOH and HCl. All samples were kept headspace free, covered with aluminum foil, and placed on a shaker table rotating at 200 rpm for 48 h before the measurement of iodoform. Analytical Methods. The concentration of I3- was measured using a UV–vis spectrometer (UV-1601, Shimadzu), with the molar absorption coefficient of 25700 M-1cm-1 at 351 nm (27). I3- has another signature peak at 288 nm which was not used due to the interference caused by the high 2920
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absorbance below 270 nm of our working solutions. A typical UV spectrum obtained from the kinetic experiment is shown in Figure 1. All UV measurements were corrected for the absorbance of the PbO2 particles, which was measured using the same solution composition but without iodide. The redox half-reaction of I- and PbO2 can be described by eqs 1 and 2: PbO2 reduction: 4H+ + PbO2 + 2e-f Pb2 + 2H2O (1) I- oxidation: I- + H2O f HOI + 2e- + H+
(2)
Considering the pH values employed in our experiments, and the proton dissociation constant of HOI (pKa ) 10.4 ( 0.1 at 25 °C (28)), HOI is the dominant species in our experiments. In the presence of excess I-, HOI is quantitatively transformed to I3- according to the following rapid equilibrium reactions: I- + HOI + H+ T I2 + H2O I
-
+ I2 T I3
(3) (4)
The equilibrium constants of eqs (3) and (4) are K3 ) 1.84 × 1012 and K4 ) 724, respectively (29). Given our experimental conditions, more than 95% of the HOI was converted to I3-. The overall redox reaction can be expressed by eq (5). PbO2 + 3I- + 4H+ f Pb2+ + I3 + 2H2O
(5)
Iodoform was analyzed by a gas chromatograph coupled to an electron capture detector (O.I. Analytical) with liquid/ liquid extraction by pentane in accordance with EPA Method 551.1 for trihalomethanes (30). An iodoform standard solution (25 mg/L) was prepared by adding reagent grade iodoform (Sigma Aldrich) to methanol. This standard was used for preparing the calibration curve and calibration check before iodoform measurements. Chromatographic separation was performed on a DB-5 capillary column (Agilent Technologies). The temperature program was set as follows: 80 °C initial, ramped to 230 at 15 °C/min, hold at 230 °C for 10 min. Solution pH values were measured by a pH meter (A15, Fisher Scientific) coupled with a calomel combination pH electrode previously calibrated with standard buffers.
Results and Discussion Effects of Solution Composition on the Reaction Kinetics. The effects of phosphate buffer concentration, PbO2 concentration, iodide concentration, and pH on the reduction of PbO2 by iodide are shown in Figures 2-5.
FIGURE 3. Effects of PbO2 concentration on the reduction of PbO2 by iodide: (a) I3- concentration as a function of time at different PbO2 concentrations; (b) I3- formation rate as a function of PbO2 concentration. pH ) 7.0, I- ) 1.2 mM, phosphate buffer ) 10 mM, temp ) 25 °C, IS NaCl ) 0.1 M.
FIGURE 4. Effects of I- concentration on the reduction of PbO2 by iodide: (a) I3- concentration as a function of time at different Iconcentrations; (b) I3- formation rate as a function of Iconcentration. pH ) 7.0, PbO2 ) 8.0 mg/L, phosphate buffer ) 10 mM, temp ) 25 °C, IS NaCl ) 0.1 M.
As shown in Figure 2, the initial rate of I3- formation did not show significant difference for phosphate buffer concentration ranging from 5 to 30 mM when other experimental parameters were kept the same. Although phosphate is a strong complexing agent and, based on our solution compositions, capable of forming Pb2+-phosphate soluble complex, a surface complex and pyromorphites [Pb5(PO4)3X, X ) OH-, Cl-] precipitates, none of these processes possess significant inhibition impacts on the reduction of PbO2 by iodide over the concentration range employed. Based on the octahedral molecular geometry of PbO2 and the bond length between Pb and O of approximately 2.5 Å (31), the site densities of surface Pb and O are estimated to be 5 and 10 per nm2, respectively, which would result in a molar ratio of phosphate to total surface site ranging from 4.76 × 103 to 2.85 × 104 for the phosphate concentrations employed in this study (5-30 mM). The adsorption constant for phosphate onto PbO2 surface is not available in the literature and an accurate determination of the phosphate surface coverage is not possible. However, if we assume a value that is similar to those for iron oxide and alumina oxide (60–337 L/mmol at pH 5 (32)), a nearly complete coverage of phosphate on the PbO2 surface sites can be deduced. The observation of no effect of increasing phosphate concentration could be due to the existence of complete surface coverage at 5 mM and above phosphate concentrations. The kinetics of I3- formation can be expressed by the following equation based on eq 5:
where R is the rate of I3- formation (M/min), kobs is the observed rate constant, and a and b are apparent reaction orders with respect to iodide and proton concentrations. The units of [I-] and [H+] are mol/L and that of (PbO2) is mg/L. The rate of I3- formation is proportional to the concentration of PbO2 present in the solution (Figure 3b). In addition, I3- formation rate increases with decreasing solution pH and increasing iodide concentration, with the apparent reaction orders of 1.68 (Figure 4b) and 0.66 (Figure 5b) with respect to [H+] and [I-], respectively. Combining all regression results and if PbO2 surface area (SA, unit: m2) is used instead of its concentration, the following rate expression with 95% of confidence in the rate constant was derived:
R ) kobs(PbO2) [I-]a[H+]b
(6)
R ) [1.79 ((0.53) × 105](SA)[I-]1.68[H+]0.66
(7)
-1m-2M-1.34.
where the unit for the rate constant is min Additional experiments using buffer-free and carbonatebuffered solutions were conducted to determine whether the presence of different types of buffer possesses an effect on the kinetics of iodide oxidation by PbO2. Figure 6 shows the formation rate of I3- as a function of pH in the absence of buffer and in the presence of 10 mM phosphate and carbonate. The kinetics in phosphate-buffered solution was found to be the fastest, followed by the carbonate-buffered, and buffer-free conditions. The general trends of I3- formation rate for each buffer showed that the differences in I3formation rate increase with decreasing pH. The catalytic effects of carbonate and phosphate were surprising because it was anticipated that they would compete with iodide for the adsorption on the PbO2 surface and result in reduced VOL. 42, NO. 8, 2008 / ENVIRONMENTAL SCIENCE & TECHNOLOGY
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FIGURE 7. Formation of iodoform as a function of NOM concentration. PbO2 ) 100 mg/L, I- ) 10 mg/L, pH ) 7.0, and reaction time ) 48 h.
FIGURE 5. Effects of pH on the reduction of PbO2 by iodide. (a) I3- concentration as a function of time at different pH values. (b) I3- formation rate as a function of pH. The insert in 3(b) shows the I3- formation rate as a function of [H+]. I- ) 1.2 mM, PbO2 ) 8.0 mg/L, phosphate buffer ) 10 mM, temp ) 25 °C, IS NaCl ) 0.1 M.
FIGURE 6. Effects of buffer type on the reduction of PbO2 by iodide. I- ) 1.2 mM, PbO2 ) 8.0 mg/L, temp ) 25 °C, IS NaCl ) 0.1 M. For the buffer-free experiments, pH values were measured at the end of experiments and the following increases in final pH were observed: pH 4.0 f 4.4; pH 5.0 f 5.5; pH 6.0 f 6.3; pH 7.0 f 7.1. iodide oxidation rates. It is not clear about the details regarding the faster kinetics in the presence of phosphate and carbonate. Perhaps a general-acid-assisted mechanism similar to that involved in iodide and bromide oxidation by hypochlorite and chloramines occurs at the water-PbO2 interface and promotes the reaction kinetics (33, 34). It is also possible that the oxidation of iodide by PbO2 might involve a transient complex of Pb(II) and an iodine-containing intermediate. The suppression of free lead concentration via the formation of Pb(II) complexes with carbonate and 2922
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phosphate might accelerate the breakdown of this transient and subsequently increase the rate of iodide oxidation. In buffer-free solutions, final pH values were measured and increases in pH value were observed (Figure 6). These increases in pH value were expected because protons are consumed in the reaction as suggested by eq 5. Our results suggested that the oxidation of iodide by PbO2 is a buffercatalyzed reaction and the rate expression, such as eq 7, is buffer-specific. Studies of reductive dissolution for other higher oxidation state metal oxides, such as manganese and iron oxide/ hydroxide, have shown that several reaction steps can occur during its reductive dissolution (or the oxidation of reductive ions), including adsorption of reductive ions on the oxide surface, electron transfer between adsorbed reductive ions and structural metal, release of reduced metal ions and oxidized reductive ions, and possibly the formation of highly reactive intermediates during the electron transfer (35–42). In a one-electron transfer reaction, if assuming that the adsorption of reductive ions and electron transfer are rate determining steps, back electron transfer is negligible, and a steady-state surface coverage of reductive ion on the oxide surface is achieved, the reaction kinetics can be described by a Langmuir–Hinshelwood or Michaelis–Menten kinetic law (35). In this rate expression, the reaction rate initially follows a first-order kinetics with respective to the reductive ion concentration (rate determined by adsorption) and reaches a zero-order kinetics after the ion concentration is greater than a critical level that saturates all the reactive sites on the oxide surface (rate determined by electron transfer), i.e., the reaction order with respective to reductive ion concentration ranges from 0 to 1. We speculate that the oxidation of PbO2 by iodide should follow a broad surface reaction scheme because of the involvement of PbO2 solid phase. However, a much more complicated mechanism may be involved because a reaction order greater than one with respective to iodide concentration was observed and twoelectrons transfer is involved in the reaction (either a twoelectron transfer reaction or two subsequent one-electron transfer reactions). Certainly, the unusual effect of phosphate may also alter the reaction mechanism described above and result in the fractional reaction order with respective to iodide concentration. Formation of Iodoform in the PbO2/I-/NOM System. NOM was added to the PbO2/I- solutions to explore the formation of iodoform using 10 mg/L iodide, 100 mg/L PbO2, and 0–5 mg/L as C of NOM at pH 7.0. The resulting chromatograms from the iodoform formation experiments are shown in Figure S1 in the Supporting Information and the relationship between NOM concentration and iodoform concentration is presented in Figure 7. In the controlled
experiment without NOM, no iodoform formation was observed. In the presence of NOM, significant amounts of iodoform were detected and the concentration increased linearly with increasing NOM concentration under our experimental conditions, with the highest of approximate 200 µg/L formed in the presence of 5 mg/L NOM. Our results clearly demonstrate that I3-, resulting from the oxidation of iodide by PbO2 (eq 5), reacts with NOM to generate iodoform. It is likely that other iodoorganic compounds, such as iodoacetic acid (19, 21), also formed in our system. Iodoform and other iodoorganic compounds have been identified as an emerging class of disinfection byproducts (DBPs) (16–21, 43). Recent toxicological studies also showed that these iodinated DBPs are much more toxic than their respective chlorinated and brominated ones (43). The formation of iodoform in the PbO2/I-/NOM system suggests that the PbO2 can serve as an oxidant reservoir and lead to the formation of halogenated DBPs. Environmental Implications. Although the experimental conditions used in this study are quite different from normal water quality with the iodide concentration approximately 100–1000 times higher than those observed in natural source waters, several important implications on PbO2 stability in drinking water can be derived from our findings. First, the capability of PbO2 to oxidize I- to form I3- reveals the strong oxidizing ability of PbO2. Release of lead through reductive dissolution of PbO2 is highly possible if PbO2 reacts with other water constituents with relatively low oxidation potentials. Changes of treatment processes that can alter the oxidation potential of water in the distribution system need to be carefully evaluated for lead release from PbO2 if leadbearing materials and free chlorine are historically used and PbO2 may already exist in the distribution system. Second, the formation of iodoform in the PbO2/I-/NOM system suggests that PbO2 can provide an oxidant reservoir of DBP formation potential in the distribution system. This represents a novel pathway for the formation of iodinated DBPs even in the absence of disinfectant. Lastly, phosphate is commonly used as a corrosion inhibitor. Very high concentrations of phosphate (5–30 mM or 155–930 mg/L as P) employed in our experiments, however, had little effect on protecting PbO2 from reductive dissolution by iodide, at least in the initial phase of reaction. The rate of PbO2 reduction by I- is moderately increased by the presence of carbonate and phosphate at these high concentrations. Considering that phosphate is typically used at several mg P/L, it is expected that phosphate would not play a significant role in the reduction of lead oxide by iodide but might play a role in the formation of Pb(II) mineral, such as pyromorphite, that prevents the release of Pb2+. Additional studies on the influence of phosphate and carbonate on the reductive dissolution of lead oxide and subsequent release of soluble lead are warranted.
Acknowledgments This work was supported by a grant from the American Water Works Association Research Foundation (Project 3172). Additional financial support from the University of Iowa Center for Health Effects of Environmental Contamination (CHEEC) is acknowledged.
Supporting Information Available Additional table and figure of experimental conditions and chromatograph from the iodoform formation experiments. This material is available free of charge via the Internet at http://pubs.acs.org.
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ES702797B
