Spectroscopic Investigation of the Interaction of Silicate Ions with Lead Carbonates Under Drinking Water Conditions

Abstract

The presence of lead has been identified as a critical health risk in drinking water systems serviced by Pb-bearing plumbing. Among several corrosion control strategies, the use of sodium silicates has attracted interest due to the advantages it offers compared to other approaches, such as phosphate dosage. However, the interaction of silicate ions with lead corrosion scales and other ubiquitous dissolved species such as Al ions in drinking water is not well understood. In this work, surface and bulk spectroscopic analysis of the solid scale is combined with quantitative analysis of the aqueous phase. A detailed spectroscopic probing of the transformations taking place on the solid phase enables us to develop a mechanistic framework for reports published in the last four years in the open literature, suggesting that silicates may not be an adequate corrosion control option in drinking water systems rich in solid lead carbonates. The spectroscopic data obtained demonstrate that in the presence of chlorine residual, silicates inhibit Pb(II) carbonates from oxidizing into less soluble Pb(IV) oxides thus, negatively impacting water quality. Furthermore, aluminum ions interact with silicates resulting in the formation of solid allophane phase over the lead scale surface, extending into the bulk. However, the formation of this new solid allophane phase does not protect against lead dissolution.

Graphical abstract

This is a visual representation of the abstract.

Keywords

Drinking water silicates lead corrosion ultraviolet–visible diffuse reflection spectroscopy UV–Vis DRS Raman X-ray photoelectron spectroscopy XPS energy dispersive X-ray EDX

Introduction

Health concerns linked to lead poisoning in adults and children resulted in the ban of lead use in paints, gasoline, drinking water service lines, and connections including solder and brass in 1986.^1,2 Exposure to lead in drinking water can lead to high blood lead levels, which have been shown to have severe health effects including increased renal dysfunction and cardiovascular disease in adults and reduced cognition and neurological development issues in infants and children.^2,3 Health Canada has set the maximum acceptable concentration (MAC) for lead in drinking water at 5 parts per billion (ppb), as exposure to even low-lead-levels poses very serious health threats.² The U.S. Centers for Disease Control (CDC) states that there is no safe level of lead exposure.⁴ There are, however, numerous municipalities across the continent with homes and schools with lead levels above the MAC. Replacing all lead service lines, although ideal, is a very complex solution to implement. Thus, municipalities that have drinking water distribution networks with lead service lines must treat their water with a form of corrosion control to keep lead levels below the MAC. Commonly used forms of corrosion control include pH adjustment and orthophosphate treatment.

Sodium silicate has been proposed as an alternative in lead corrosion control in drinking water systems. Sodium silicate solutions consist of mono- and polymeric silicate species and are attractive for use as corrosion inhibitors due to their low environmental impact. It has been hypothesized that silicates inhibit lead release by forming a protective film on the lead surface.^5–7 Past research has demonstrated the effectiveness in both pilot and full-scale studies when pH was not controlled.^7–9 However, more recently, the effectiveness of sodium silicates as a corrosion control option for lead in drinking water distribution systems has been questioned.^10,11 These studies suggest that silicates may not be an effective method of corrosion control compared to concurrent pH adjustment and orthophosphate treatment.¹⁰

For instance, Li and collaborators compared the effectiveness of sodium silicates to orthophosphate.¹⁰ Compared to orthophosphate, sodium silicate treatment was not effective at decreasing dissolved lead levels. The authors hypothesized the formation of a nanometer thick silicate coating in the silicate-treated system and identified lead carbonate as the dominant scale phase. However, they found no evidence that this silicate coating inhibited lead release. Aghasedeghi and collaborators¹¹ conducted a pipe loop study in a lead carbonate rich system. At the two levels of silicate dosage used, dissolved lead levels were higher compared to pH adjustment alone. The increase in dissolved lead with silicate treatment was attributed to formation of colloids containing lead mobilizing into the aqueous phase. They observed Si and Al accumulated in the same scale layers suggesting possible interaction between these two ions; however, they reported that the presence of Al increased Pb release.¹¹ Another study evaluated the effect of silicates at consistent pH on corrosion control on a pipe loop system and observed silicates effectively reduce lead levels. The corrosion scales in the lead service lines used consisted of hydrocerussite, PbO, and SiO₂ as well as oxides of silicon and aluminum.¹² They observed a decrease in equilibrium dissolved lead concentrations in the pipes treated with 16 and 20 mg/L sodium silicates at pH 7.7. However, analysis of the scale from the treated pipes using scanning electron microscopy energy dispersive spectrometry (SEM-EDS) showed only a marginal uptake in silicon (0.7–1.5%). The authors proposed a uniform uptake of silica throughout the top layer of the scale and hypothesized that in the presence of silicates the system is forming Si- and Al-containing precipitates that are either adsorbing or depositing as an undetectable amorphous solid phase. This body of work clearly indicates that to develop the most effective corrosion control strategy for a given water quality, a thorough understanding of the mechanisms governing these phenomena is required. There is a knowledge gap specifically surrounding the mechanism of lead dissolution control by sodium silicates. We have attempted to provide a more fundamental insight on these processes by carrying spectroscopic studies on a pure phase lead compound commonly found in corrosion scale, cerussite^13–16 and its transformation in chlorinated drinking water at consistent pH in the presence of sodium silicate and aluminum. We comprehensively evaluated the solid precipitate formed in these systems after exposure to silicate-dosed synthetic drinking water using ultraviolet–visible diffuse reflection spectroscopy (UV–Vis DRS), Raman, X-ray photoelectron spectroscopy (XPS), and energy dispersive X-ray (EDX) analysis to probe the mechanistic details describing the relationship between lead corrosion scale transformation and the presence of sodium silicate in the aqueous phase. We found that, in the presence of chlorine residuals, sodium silicates prevent the oxidation of Pb(II) to Pb(IV) solids, interfering with the formation of less soluble lead oxides in the solid phase, consequently resulting in elevated aqueous lead levels. Furthermore, we confirmed previously reported electron microscopy results, as we observed that the presence of aqueous aluminum results in the formation of solid allophane on the surface of the lead solid phase, which extends into the bulk.¹² The presence of this solid aluminosilicate, however, did not result in lower dissolved lead levels.

Experimental

Materials and Methods

Cerussite (PbCO₃) (Sigma-Aldrich, American Chemical Society (ACS) reagent grade) was used as the source of lead. Sodium hypochlorite (NaOCl; Sigma-Aldrich, 10–15% available chlorine, reagent grade) was used as the source of free chlorine. Sodium sulfate (Na₂SO₄; Sigma-Aldrich, >99.0%, ACS reagent grade), sodium chloride (NaCl; Sigma-Aldrich, >99.0%, ACS reagent grade), calcium hydroxide (CaOH; Sigma-Aldrich, >95.0%, ACS reagent grade), and sodium bicarbonate (NaHCO₃; Sigma-Aldrich, >99.7%, ACS reagent grade) were used as sources of sulfate, chloride, hardness, and alkalinity, respectively. All precursor salts were dissolved in megapure water (resistivity of 3 MΩ) to prepare the synthetic drinking water. Necessary pH adjustments were performed with nitric acid (HNO₃; Sigma-Aldrich, 70%, purified by redistillation, 99.999% trace metals basis). Aluminum nitrate nonahydrate (Al(NO₃)₃·9H₂O) was used as the source of aluminum (Sigma-Aldrich, >98.0%, ACS reagent grade). N-sodium silicate (National Silicates, Na : Si ratio 3.22) was used as the source of silicates. Synthetic drinking water was prepared according to the parameters shown in the Supplemental Material (Table S1). A mother solution of synthetic drinking water was prepared before each series of experiments. Four different sets of batch reactor experiments were prepared to evaluate the role of different parameters in the interaction of cerussite with silicate. The set of experimental conditions used is summarized in Table S2 (Supplemental Material).

Two different reactor sizes were used, either 10 or 500 mL high-density polyethylene (HDPE) vials. The selection of the reactor size was a function of the amount of sample required for solid phase characterization. At the beginning of the experiment either 10 or 500 mL of synthetic water was measured using a graduated cylinder and poured into the 10 or 500 mL HDPE reactor, together with either 0.5 g (10 and 500 mL reactors) or 5 g (500 mL reactors) of cerussite. The reactors were capped, placed on a shaker (ThermoScientific MaxQ 2000), and homogenized at 170 rpm. For the case of tests conducted using 10 g/L of cerussite, the reactors were first placed on a magnetic stirrer for 10 days, then taken off the stirrer for the remainder of the experiment.

Spectroscopic Analysis of the Solid Phase

Ultraviolet–visible spectroscopy was performed on the dried solid phase samples using a UV-3600 Shimadzu spectrophotometer equipped with a Harrick Praying Mantis cell for diffuse reflectance. The wavelength range used was 200 to 800 nm at a sampling interval of 1 nm. The spectra were obtained in diffuse reflectance mode and transformed to pseudo absorption using the Kubelka–Munk function. Raman analysis was performed using a Renishaw micro-Raman 2000 system. The spectra were obtained using a He–Ne laser with a 632.8 nm excitation wavelength, 0–2000 cm^–1 range in micro mode. The acquisition time was 20 s at 2 mW for each spectrum. XPS was performed using a Kratos Axis Supra spectrometer with a monochromatic AI Kα source (15 mA, 15 kV). The Kratos charge neutralizer system was used on all samples. Survey scans were carried out with an analysis area of 300 × 700 μm and a pass energy of 160 eV, and high-resolution analysis was carried out with the same analysis area but a pass energy of 20 eV. Spectra were charge corrected to the main line of the carbon 1 s (adventitious carbon) spectrum set to 284.8 eV. Scanning electron microscopy (SEM) and EDX were performed with the Zeiss 1540XB field emission SEM, Oxford Instruments X-Max X-ray detector, and Inca analysis software. EDX analysis was performed at 20 keV beam energy, and SEM imaging was performed at 1 keV beam energy. The XRD data were obtained on a Rigaku RPT 300 RC diffractometer using Co Kα (λ = 1.78890 Å) radiation over the range of 10–70° 2θ with a 0.02° step size; analysis of results was performed using Jade software and compared to the RRUFF Database.

Aqueous phase analysis: For dissolved lead measurements, 10 mL aliquots were taken, filtered through a 0.2 μm polyethersulfone (PES) filter and preserved in HDPE vials with 1 mL 70% nitric acid. Free chlorine and pH were monitored consistently and adjusted, as necessary. The last set of experiments, carried out using varying amounts of aluminum, were run by equilibrating the solid/aqueous mixture using a magnetic stirrer for two days, then taken off the stirrer for the remainder of the first iteration. After 21 days, the experiments were spiked with an additional 1 mg/L of Al³⁺, then again after seven days, and then at shorter intervals.

Chlorine levels were monitored with 10 mL sacrificial samples using the DPD free chlorine Hach colorimetric method 8021. Sacrificial samples (10 mL) were also collected to measure silicate as mg/L SiO₂ using the silicomolybdate Hach colorimetric method 8185. Silica and chlorine measurements were both performed using a Hach DR900 colorimeter. The pH was measured using a Thermo Scientific Orion Star A111 pH meter with a non-glass, ion-sensitive field-effect transistor (ISFET) probe. Then 5.6 mL sacrificial samples were taken to measure alkalinity using the Hach drop count titration AL–AP method with sulfuric acid. Following completion of the experiment, vacuum filtration with a 0.45 μm membrane filter was used to separate the solid for further analysis. The solid sample was rinsed twice with megapure water prior to drying at room temperature. Dissolved lead concentrations were measured from the acid preserved 10 mL aliquots taken throughout the experiment using inductively coupled plasma mass spectroscopy (ICP-MS) using the U.S. Environmental Protection Agency (EPA) method 6020B.

Results and Discussion

Aqueous free chlorine monitoring and adjustments were carried out periodically in all experiments. In a batch system, chlorine depletes relatively quickly in the presence of solid lead phases.^17,18 Specifically, for the case of lead carbonates, it is well known that they are typically oxidized in the presence of free chlorine under typical drinking water conditions:¹⁹

{PbCO}_{3} + {OCl}^{-} + 2 {OH}^{-} \to {PbO}_{2} + {Cl}^{-} + {CO}_{3}^{- 2} + H_{2} O

(1)

In our experiments, the chlorine residual depleted, and it was necessarily to adjust it back to initial levels every four to eight days (Figure S1, Supplemental Material). However, a slower rate of chlorine depletion was observed in the experiments carried out in the presence of silicates. This effect was insensitive to the presence of aluminum in the system. As time progressed, a change in color of the solid phase was observed for the samples where silicate was not present. Indeed, for the samples run in the absence of silicate (PbCO_3(s)/Al³⁺_(aq) and PbCO_3(s)) the solid phase changed from the characteristic white of lead carbonates to an orange-brown (Figure S2, Supplemental Material). This observation hints at a change in the composition of the solid lead phase. XRD analysis of these samples confirmed the transformation of cerussite to hydrocerussite after both dissolution experiments (Figure S2), but a defined crystalline phase associated with lead oxides was not immediately apparent. However, both cerussite and hydrocerussite are white, whereas higher oxidized Pb species display color. Based on thermodynamic data and previously reported work,⁷ a phase change could be hypothesized from PbCO₃ to Pb₃O₄ and/or PbO₂ through oxidation by free chlorine.¹⁴ PbO₂ occurs as red-brown or black in color. Minium has been identified as an intermediate phase in the oxidation of Pb(II) lead carbonates to Pb(IV) lead oxides and displays a bright orange color.^14,18 The formation of these two phases could explain the color change that occurred over the course of the experiment in the absence of silicates. In the presence of dissolved silicates, no obvious color change was observed. Lytle and collaborators reported a similar finding in their study of the effects of orthophosphate on lead phase transformation from hydrocerussite to lead oxides. Throughout their experiment in the absence of orthophosphate, they observed white hydrocerussite darkening in color to shades of dark red and gray, as well as a decrease in lead solubility.²⁰

To obtain a more quantitative description of this observation we carried a mass balance, assuming that chlorine was only consumed through a direct redox pathway where Pb(II) is oxidized to Pb(IV) in the system (Eq. 1). In this scenario, free chlorine is consumed in a 1 : 1 molar ratio to form Pb(IV) species from either aqueous or solid Pb(II). Using this equation, we calculated a hypothetical percentage of Pb(IV) formed using the total lead present in the system (aqueous and solid) as a benchmark. The results obtained from this calculation are presented in Figure 1 (left panel) as a function of time. The assumption of a direct redox reaction between chlorine and cerussite represents the upper limit of Pb(IV) generated. As expected, the amount of Pb(IV) in the system steadily increases over time. However, the presence of silicates dramatically slows down this process. There is almost an order of magnitude difference in final amount of Pb(IV) formed with and without silicates. Moreover, the presence of aluminum ions does not affect the observed trend in the case of the silicate-free system (circles versus triangle data points); however it seems to have a small (though minor) effect for the case where silicate is present.

Figure 1.

Left panel: Nominal Pb(IV) formation resulting from chlorine consumption as a function of time for cerussite exposed to silicate free (circles and triangles) and silicate (squares and diamonds) treated water. Right panel: Observed dissolved levels obtained on the same experiments, dashed bars: silicate free system, solid bars: silicate 20 mg/L (as SiO₂) treated water. Experiments were carried at pH = 8. The error bar depicts the standard error of the mean based on the results obtained from a minimum of three different sampled aliquots.

The results obtained for the ICP analysis of aqueous dissolved lead levels in these experiments are also shown in Figure 1. Average dissolved lead levels fall within the region corresponding to the equilibrium concentration of aqueous dissolved lead from cerussite, 25–250 μg/L using K_sp values previously reported.^18,21 Results indicate that, regardless of the presence of silicates, lead levels are regulated entirely by the thermodynamic solubility of lead carbonate.

Compared to carbonates, lead (IV) oxides are insoluble in drinking water.^16,20 Plattnerite is found in water distribution systems that have maintained high free chlorine residuals due to free chlorine oxidizing Pb species to the insoluble Pb(IV) product PbO₂.^15,17,22 Water that has been treated with other disinfectants, such as monochloramine, has a much higher presence of solid Pb(II) species in the scale. Chloramines are unable to provide the thermodynamic driving force needed to produce or keep Pb(IV) species, and as a result some systems that switched to monochloramine have, in the past, experienced dissolved lead concentrations above action levels.^23,24 The differences observed in the dissolved lead levels in Figure 1 are consistent with these studies. In the absence of silicates, a faster change to higher valence lead oxides is taking place, which in turn results in the formation of a less soluble lead oxide phase in the solid (Figure 1, right panel, dashed bars). When silicates are present, aqueous chlorine is unable to rapidly trigger the formation of Pb(IV) species, and as a consequence slightly larger amounts of dissolved lead are generated (Figure 1, right panel, solid bars). This result is also consistent with reports of higher dissolved lead levels in the presence of silicates in a pH-controlled pipe loop study.¹¹ In that case the authors attributed this observation to the presence of colloids containing lead species. Although the intensities of the XRD peaks characteristic for plattnerite were similar among all pipes characterized on that study, the information reported for the solid phase composition by the authors (and previous studies by other researchers²⁵) clearly indicates that the solid phase present in the scale is rich in lead carbonates. Hence, the results on higher lead levels reported by the authors could be attributed to the inherent solubility of these carbonate species.

Aqueous levels of silicate were monitored through all these experiments. Aqueous silicate measurements were taken every three to seven days (Figure S3, Supplemental Material). The aqueous silica levels decreased in the presence of aluminum. The decrease in silica levels in the presence of aluminum hints at the formation of a solid aluminosilicate phase, potentially on the solid phase lead surface, as previously reported for the case of the outermost scale layers harvested from lead service lines.^13,14,26 Some studies have proposed aluminosilicate formation on pipe surfaces may have a protective effect on metal ion dissolution, and many have investigated the prevalence of aluminosilicates in corrosion scale harvested from municipalities across North America.²⁶ In phosphate-treated systems, it has been shown that the presence of aluminum increased dissolved lead levels due to interference with formation of protective lead-phosphate compounds on the pipe surface.²⁷ Other studies have also proposed that the presence of aluminum may increase dissolved lead levels in the absence of orthophosphate.¹¹

Allophane and imogolite are two most common aluminosilicate phases identified in drinking water systems. Studies have found that allophane is the most common phase formed under drinking water conditions with the Si/Al ratio in the solid affected by pH.^13,28,29 The solubility of allophane is not only controlled by pH, but also the speciation of silicates. Moreover, in contrast with other aluminum bearing phases such as gibbsite, the K_sp (solubility product constant) seems to be a function of the degree crystallinity (or properly stated, lack thereof) of the allophane solid phase.³⁰ Some reports indicate Si-rich allophanes with an Al/Si ratio near 0.6 forms naturally at pH above 5 to near neutral and to have a low solubility at pH below 7.3,^13,31 though other studies suggest a pH closer to 8 before significant dissolution of this phase is observed.³²

Allophane typically has the composition Al₂O₃(SiO₂)_1.3−2·(H₂O)_2.5−3 and has a positive surface charge.¹³ The Al/Si ratio in the solid depends on pH and affects the solid surface properties.^13,33 Pb(II) has been shown to sorb selectively to allophane and imogolite in soils. Previous studies have shown that the surface of cerussite and hydrocerussite under drinking water conditions have negative zeta potentials, indicating that the positively charged aluminosilicates could adsorb to the surface of the solid lead carbonate.^13,33 Furthermore, the zeta potential of sodium silicate-treated iron oxide has been shown to be substantially more negative than untreated iron oxides; this result is also expected for other transition metals, such as lead.³⁴ Sodium silicate-treated lead carbonate or oxide could have a more negative surface charge which would result in enhanced deposition of aluminosilicates to the lead surface.

We tested the possibility that an aluminosilicate phase is forming on the solid lead phase by designing a separate set of experiments with a larger amount of solid phase (10 g/L PbCO₃) under the hypothesis that a higher amount of solid surface area would increase the number of sites available for adsorption–precipitation. The underlying hypothesis assumes that aluminosilicate precipitation is occurring at the surface of the solid lead phase, thus Pb-phase mediated, and not freely in the aqueous phase through direct interaction of silicate and aluminum ions reaching the thermodynamic solubility product. The aluminum dosage was consistently spiked with 1mg/L over time to attempt to trigger decreases in aqueous silica levels using cerussite at both pH 7 and 8. The results obtained for aqueous silica levels at pH 8 in these experiments are shown in Figure 2. Consistently, after each aluminum spike of 1 mg/L, a decrease in aqueous silicate levels is observed. The amount of silicate precipitating, however, depends on the system's pH. For the case of the experiment run at pH 8, a consistent decrease of 2 mg/L as SiO₂ (Figure 2) is observed after a 1 mg/L Al spike. At pH 7, each aluminum spike of 1 mg/L led to a decrease in aqueous silica levels of only 1 mg/L (Figure S4, Supplemental Material). When aluminum was not present, varying levels of aqueous SiO₂ were observed with no consistent clear trend over time. As discussed above, at this pH the surface of both lead carbonates and oxides is negative; therefore both would attract the positively charged aluminosilicates towards surface precipitation.

Figure 2.

Variation in the concentration of aqueous silicate concentration over 120 days of solid cerussite exposed to chlorinated drinking water at pH 8 with increasing dosing of aqueous aluminum ions.

It is well stablished that in aqueous solutions at neutral pH, aluminum ions are forming hexa-aqueous complexes of the form (Al(OH₂)_x(OH)_y)^+3–y.³⁵ At pH 7, the main species present are (Al(OH₂)₄(OH)₂)⁺¹, (Al(OH₂)₃(OH)₃)⁰, and (Al(OH₂)₂(OH)₄)^–1.³⁶ We could propose that positively charged aluminum complexes act as a bridge between the negative solid surface and the negatively charged silicate ions. Given similar atomic weights of Si (28 g/mol) and Al (27 g/mol), it is remarkable that the results indicate that, at pH 7, per each atom of aluminum introduced to the system, one atom of silicate is lost from the aqueous phase into the solid surface and, moreover, this stoichiometry is pH dependent. Indeed, when the pH is raised to 8, twice as much (2 mg/L as SiO₂) aqueous silicate is lost to the solid phase after each 1 mg/L Al spike to the system. These values are consistent with those previously reported for allophane formation.^14,33 We could further rationalize this observation once again assuming the positively charged aqueous aluminum complexes are acting as an interface between the negatively charged solid lead surface and the negatively charged aqueous silicate ions. At higher pH, three effects take place: the surface of the solid lead bearing phase becomes more negatively charged (if no total charge saturation has taken place yet), the coordination sphere around the silicate ions might also become more negative, and the speciation of the aqueous aluminum complexes changes due to hydroxyl ligand substitution in the inner coordination sphere,³⁷ resulting in a larger amount of (Al(OH₂)₂(OH)₄)^–1 species at the expense of the positively charged and neutral aluminum complexes. Hence, the effectiveness of this “interface” effect would be very sensitive to aqueous pH levels, potentially explaining reports linking the efficacy of silicate treatment to pH.¹²

The final solids obtained after the experiments depicted in Figure 2 and those obtained from two blank experiments (cerussite before dissolution and only Al dosing during dissolution) were characterized by SEM. Figure 3 shows the results obtained at two different magnifications. The results indicate that the morphology of all samples is different from that of the cerussite reference sample. The crystal morphology obtained for the sample after exposure to synthetic drinking water suggests the presence of both hydrocerussite and plattnerite on the surface of the solid phase.³⁸ Some hydrocerussite on the reference cerussite sample surface is also observed (indicated by a solid black circle) which is attributed to a reaction between the cerussite and moisture in the air. Clusters with morphology similar to PbO₂ in the PbCO_3(s)/Al³⁺_(aq) sample that are much larger and more prevalent than in the samples exposed to aqueous silicate (PbCO_3(s)/SiO_x(aq) or PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq)) were observed (shown with white circles), once again suggesting the inhibition of PbO₂ formation from PbCO₃ in the presence of silicates. A fine precipitate with a distinct morphology appears visible on the PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq) sample (shown by dashed black circles). At higher magnification, the presence of this distinct fine precipitate is more obvious. With a relatively small amount of a precipitate visible on the smooth surface of hydrocerussite in the sample exposed to silicates in the absence of Al⁺³ (PbCO_3(s)/SiO_x(aq) sample, circled in blue), and a much larger amount on the surface of the sample that was exposed to both Al and Si aqueous ions (PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq)). Moreover, even at higher magnification this precipitate is only visible on the samples exposed to aqueous silicate. We could hypothesize that this phase is composed of mainly silicates, depositing on the sample surfaces and more silicate is precipitating out in the presence of aluminum as previously inferred from the analysis of dissolved silicate in the aqueous phase.

Figure 3.

SEM images obtained in harvested solid cerussite after 120 days dissolution in chlorinated synthetic drinking water at pH 8, with and without exposure to aqueous Si and Al ions. Black solid circles show hydrocerussite, white circles show plattnerite and dashed black circles show aluminosilicate formation on the lead carbonates surface.

Energy dispersive X-ray analysis was performed on the solid samples exposed for 45 days to aqueous silicate and aqueous silicate + aluminum (experiment depicted in Figure S3, Supplemental Material) to determine bulk phase composition and to quantify the precipitation of a silicate phase. XPS analysis was carried on the same samples as well, to evaluate the possibility of a previously reported bulk migration of silicate into the solid lead phase.¹² These results are shown in Tables I and II, respectively. The EDX results quantitatively confirm that the highest amount of Si in the solid is observed for the sample that was exposed to both aqueous silicate and aluminum ions (PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq), 1.0 at% of solid Si content in the sample), whereas only a small amount is present in PbCO_3(s)/SiO_x(aq) at 0.2 at%. Indeed, in the solid, the atomic ratio Si : Pb is five times higher when aqueous Al is present. This observation is consistent with the observation of the fine precipitate observed in the high magnification SEM images and with aluminum playing a role in the interaction between the aqueous silicate and the solid lead phase. Specifically, the presence of aluminum ions promotes the precipitation of silicates from the aqueous phase in drinking water.

Table I.

Bulk composition of cerussite treated in chlorinated water at pH 8 for 45 days with and without silicate presence as obtained by EDX analysis.

Element	PbCO₃ before dissolution	PbCO_3(s)/Al³⁺_(aq) After 45 days	PbCO_3(s)/SiO_x(aq) After 45 days	PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq) After 45 days
	Atomic (%)
Pb	8.5	9.4	8.1	9.4
C	43.2	49.4	55.6	44.4
O	48.3	39.5	36.2	44.8
Al	Not detected	1.7	Not detected	0.4
Si	Not detected	Not detected	0.2	1.0

Table II.

Surface composition of Cerussite treated in chlorinated water at pH 8 for 45 days with and without silicate presence as obtained by XPS analysis.

Element	PbCO_3(s)/Al^3 +_(aq)After 45 days	PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq)After 45 days
Element	Atomic (%)
Pb	17.8	13.1
C	32.5	23.1
O	47.0	52.4
Ca	1.6	1.3
Al	0.8	3.5
Si	Below detection limit	6.6

The XPS survey scan (Table II) also confirmed the formation of a silicate precipitate in the presence of aluminum. The results presented in Tables I and II are summarized in Figure 4 in terms of atomic ratios for Si : Pb and Si : Al, together with the values for these two ratios present in the overall system (aqueous + solid phase, value represented by the dashed horizontal line). To begin, without Al present, the Si : Pb ratio in the solid reflects that of the overall system, i.e., there is no clear silicate enrichment in the solid, while when Al is present in the aqueous phase, the solid gets enriched in silicon. However, a higher Si : Pb ratio is observed in the surface (∼0.5 in the XPS) than in the bulk of the samples (∼0.1 in the EDX), and these values are an order of magnitude above of that in the overall system (solid + aqueous) clearly indicating that the solid is indeed enriched with silicon. The Si : Al ratios obtained, however, are similar for both the solid surface (XPS) and the solid bulk (EDX). This result suggests the silica–alumina precipitate formed extends well into the bulk phase with a similar atomic ratio of that in surface layers and that silicon ions are indeed migrating into the solid bulk, perhaps in the form of an aluminosilicate. This result is consistent with two reports, where SEM-EDX indicated an increase of 0.7–1.5 wt% in pipe loop scale Si concentration following sodium silicate treatment and the author's proposed hypothesis of silicate migrating into the bulk of the corrosion layer.^11,12 The results portrayed in Figure 4 together with the SEM micrographs clearly show that after aqueous silicate exposure the sample is heterogeneous with a silica–alumina rich outer layer.

Figure 4.

Si:Pb (top) and Si:Al (bottom) atomic ratios as measured using EDX (left) and XPS (right) obtained in harvested solid cerussite after dissolution in chlorinated synthetic drinking water at pH 8 for 45 days. The black horizontal dashed line represents the total (solid + aqueous) ratio in the system (Si : Pb = 0.02, Si : Al = 10).

While all these results confirm the formation of a solid aluminosilicate deposit on the surface of the solid lead phase, they do not explain the observed increase in dissolved lead levels in presence of aqueous silicates. Raman and UV–Vis spectroscopy were used to identify any potential changes in the surface state of the lead solid phase, which might have been triggered upon exposure to the aqueous phase. The spectra obtained are presented in Figure 5. UV–Vis reference spectra for cerussite, minium (Pb₃O₄), and plattnerite are presented in Figure S5 (Supplemental Material). The band centered around 200–300 nm present is in all samples characteristic of lead carbonate species.¹⁸ Figure 5 shows this band is present in all samples after exposure to the aqueous phase, regardless of the presence of dissolved silicate. The broad band from 300 to 600 nm is characteristic of lead oxides. As Figure S5 shows, the absorption spectrum of minium extends to ∼550 nm; while that of plattnerite extends to the NIR region. An inspection of Figure 5 indicates that the samples exposed to chlorinated water in the presence of aqueous silicate do not experience a phase transformation; the spectral features observed (band at 200–300 nm) are only those attributed to cerussite. However, the spectra obtained for the samples where silicate was not present (PbCO_3(s) and PbCO_3(s)/Al³⁺_(aq)) display a second absorption band starting at ∼300 nm that seem to extend to the NIR region. This feature indicates the presence of a lead oxide phase. Based on the UV results, this lead oxide phase present in the samples obtained after exposure to water in the absence of silicate is likely a combination of minium and plattnerite. These results are consistent with the color change observed over the course of the 45-day experiment, as well as the presence of Pb(IV) oxide clusters observed in the SEM images. When combined, these results clearly show that after exposure to drinking water conditions in the absence of silicate ions, a Pb(IV) oxide phase is formed on the cerussite surface (likely the result of the interaction of chlorine residual with the lead bearing solid surface). In contrast, if silicate is present in the aqueous phase, this process does not take place, and a lead oxide phase is not formed.

Figure 5.

UV–Vis (left) and Raman (right) spectra obtained in the solids harvested 45 days after cerussite dissolution experiments at pH 8 using different conditions; from top to bottom: aqueous sodium silicate and aluminum, aqueous silicate, aqueous aluminum, and silicate and aluminum free aqueous phase. The inset on the top right shows a magnified area of the spectral Raman region.

Raman spectroscopy was performed on the same samples as well, results shown in Figure 5. It is evident from a quick inspection of the Raman spectra that the samples obtained in the presence of silicate exhibit different spectral features than those obtained in the absence of aqueous silicate. Specifically, the samples obtained in the presence of aqueous silicate (PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq) and PbCO_3(s)/SiO_x(aq)) show a sharp peak at 1045 cm^–1, as well as other distinct peaks at 1355, 410, and 110 cm^–1. On the other hand, the samples obtained without exposure to aqueous silicate (PbCO_3(s) and PbCO_3(s)/Al^3 +_(aq)) show main distinct peaks only at 515, 260, and 130 cm^–1. The main peak for cerussite occurs at 1050 cm^–1, and for hydrocerussite at 1040 cm^–1.¹⁴ These features are not present in the spectra obtained on the samples treated in the absence of aqueous silicates, whereas this feature dominates the spectra of the silicate-treated samples, thus confirming the UV–Vis results: dominant presence of lead carbonates in the sample surface when silicates are present in the aqueous phase. The peaks observed in the 600–100 cm^–1 region can be attributed to lead oxides or lead carbonate, while the one at 1355 cm^–1 is characteristic of carbonates.^14,39 Two small peaks, located at 1076 and 960 cm^–1, are present in the spectra obtained on the silicate-treated samples (PbCO_3(s)/Al³⁺_(aq)/SiO_x(aq) and PbCO_3(s)/SiO_x(aq)) that are not observed in lead oxides or carbonates. It is well established that Si–O–Si peaks are located at ∼1070 and ∼950 cm^–1.⁴⁰ Therefore, we can attribute these specific features to Si–O–Si or Si–OH moieties. Assignation of these last two features to allophane or imogolite would require evaluation of O–H stretchings in the infrared region. The XRD results (see below) did not show features associated with these aluminum silicate phases, but this could be the result of their lack of crystallinity.⁴¹

Both the UV–Vis and Raman results strongly indicate that in the absence of silicate, cerussite undergoes a phase transformation, likely triggered by chlorine, to less soluble lead oxides. We carried an XRD analysis aiming at identifying a crystalline lead oxide phase, but although the results (Figure S6) hint at the presence of PbO₂ (2θ = 29.4, 37.3°) they are inconclusive. This is likely because lead oxide phases formed in the surface have not reached the crystallinity needed (particle size) to generate well defined diffractograms. We therefore relied on the XPS features to carry this assessment. These data obtained for Pb is shown in Figure S7 (Supplemental Material), together with the high-resolution scans obtained in the Si 2p and Al 2p XPS regions.

The Pb binding energy (4f_7/2) of cerussite is ∼138.0 eV. While lead oxides containing Pb(IV) species display lower binding energies, Pb₃O₄ typically showing values between 137.4 and 137.7 eV and PbO₂ between 136.8 and 137.6 eV.^42,43 It is not possible to differentiate between the lead oxide binding energies because they are too similar. However, it is technically possible to differentiate between oxides and carbonates which can give an indication as to whether the surface layers contain Pb(II), Pb(IV), or a combination of both species. Figure S7 shows the intensity normalized survey scan obtained in the Pb 4f region for two of the samples exposed to synthetic chlorinated drinking water. One sample was exposed to aqueous silicate, while the other was not. As indicated by the arrow, the hint of a shoulder is observed on the Pb 4f peak for the sample obtained in the absence of aqueous silicate (PbCO_3(s)/Al³⁺_(aq)). This shoulder appears at a binding energy ∼137.5 eV and suggests that in the absence of silicates a highly oxidized Pb(IV) phase is formed. This XPS feature agrees with our previous observations: in the absence of silicates the solid consists of two lead phases, Pb(II) from PbCO₃ (as observed in the UV–Vis and Raman analysis) and Pb(IV) from either Pb₃O₄ or PbO₂. The UV–Vis, Raman, and XPS results thus confirm that Pb(IV) is not formed when the aqueous phase contains silicates.

The values obtained for the Si 2p (101.7 eV for PbCO_3(s)/SiO_x(aq)) and the Al 2p features (74.4 eV) on the sample exposed to the presence of both aluminum and silicate ions are consistent with the formation of a solid allophane phase.^13,27,44 In the absence of silicates, the Al 2p feature is located at 73.9 eV, suggesting the aluminum hydroxide phase gibbsite is present instead, which is hypothesized to form in aluminum-rich drinking water systems.^27,29 Therefore, it is likely that in the absence of silicates, gibbsite forms on the lead surface, while in the presence of silicates, aluminum resides in allophane as the main phase.

Conclusion

The implications of these results are significant and reinforce the importance of developing a thorough fundamental understanding of corrosion control processes to mitigate lead dissolution in drinking water systems. Our results show that silicate interferes with the oxidation of lead carbonates by free chlorine, this being insensitive to the formation of a solid silicate phase. The presence of aqueous silicate thus prevents the solid phase from interacting with the aqueous phase; as a result, under our experimental conditions, the dominant solid lead phase remained Pb(II), and dissolved lead levels increased. Under pH, DIC and redox potential conditions where lead dissolution is regulated by carbonates, the presence of silicates could slow down the transformation of cerussite/hydrocerussite into less soluble PbO₂-bearing phases, potentially resulting in serious consequences for municipalities considering the use of silicates as a corrosion inhibitor. Although, within a limited range of water quality, the results of our work suggest that silicates should be cautiously evaluated as a corrosion inhibitor alternative in water systems that have higher amounts of Pb(II) solids in the corrosion scale or where chlorine residual is the only parameter regulating passivation of lead bearing plumbing. Further long-term studies are required to investigate whether the aluminosilicate solids forming on the lead surface would eventually function as a passivation layer and provide protection against lead dissolution.

Supplemental Material

sj-docx-1-asp-10.1177_00037028241291072 - Supplemental material for Spectroscopic Investigation of the Interaction of Silicate Ions with Lead Carbonates Under Drinking Water Conditions

Supplemental material, sj-docx-1-asp-10.1177_00037028241291072 for Spectroscopic Investigation of the Interaction of Silicate Ions with Lead Carbonates Under Drinking Water Conditions by Hailey Holmes and José E. Herrera in Applied Spectroscopy

Footnotes

Acknowledgments

This research project was partially funded by the National Silicates and the Natural Sciences and Engineering Research Council of Canada. We would like to express our gratitude to Dr. Daoping Guo and Vicky Sidorkiewicz for useful insights and discussions. H.H. acknowledges the additional funding provided through an Ontario Graduate Scholarship. Some of the spectroscopic studies were carried using instrumentation purchased through the Canadian Foundation for Innovation.

Declaration of Conflicting Interests

The authors declared no potential conflicts of interest with respect to the research, authorship, and/or publication of this article.

Funding

The authors disclosed receipt of the following financial support for the research, authorship, and/or publication of this article: This work was supported by the Natural Sciences and Engineering Research Council of Canada, (grant number RGPIN-2019-06633).

ORCID iD

José E. Herrera

Supplemental Material

All supplemental material mentioned in the text is available in the online version of the journal.

References

O’Connor

Hou

Zhang

Y.S.

Song

, et al. “Lead-Based Paint Remains a Major Public Health Concern: A Critical Review of Global Production, Trade, Use, Exposure, Health Risk, and Implications”. Environ. Int. 2018. 121(Pt. 1): 85–101. https://doi.org/10.1016/J.Envint.2018.08.052

Health Canada. “Health Canada Sets New Guideline for Lead in Drinking Water”. https://www.canada.ca/en/health-canada/news/2019/03/health-canada-sets-new-guideline-for-lead-in-drinking-water-latest-in-series-of-government-actions-to-protect-canadians-from-exposure-to-lead.html [accessed Sep 24 2024].

Hanna-Attisha

Lachance

Sadler

R.C.

Champney Schnepp

. “Elevated Blood Lead Levels in Children Associated with the Flint Drinking Water Crisis: A Spatial Analysis of Risk and Public Health Response”. Am. J. Pub. Health. 2016. 106(2): 283–290. https://doi.org/10.2105/AJPH.2015.303003

U.S. Centers for Disease Control and Prevention (CDC). “Lead in the Workplace”. 2018. https://www.cdc.gov/niosh/lead/ [accessed Sep 26 2024].

Stericker

. “Protection of Small Water Systems from Corrosion”. Ind. Eng. Chem. 1945. 37(8): 716–720. https://doi.org/10.1021/ie50428a015

Larose Thompson

Scheetz

B.E.

Schock

M.R.

Lytle

D.A.

Delaney

P.J.

. “Sodium Silicate Corrosion Inhibitors: Issues of Effectiveness and Mechanism”. https://studylib.net/doc/18217123/sodium-silicate-corrosion-inhibitors–issues-of [accessed Sep 24 2024].

Schock

M.R.

Lytle

D.A.

Sandvig

A.M.

Clement

Harmon

S.M.

. “Replacing Polyphosphate with Silicate to Solve Lead, Copper, and Source Water Iron Problems”. J. Am. Water Works Assoc. 2005. 97(11): 84–93. https://doi.org/10.1002/j.1551-8833.2005.tb07521.x

Zhou

Payne

S.J.O.

Hofmann

Andrews

R.C.

. “Factors Affecting Lead Release in Sodium Silicate-Treated Partial Lead Service Line Replacements”. J. Environ. Sci. Health, Part A: Toxic/Hazard. Subst. Environ. Eng. 2015. 50(9): 922–930. https://doi.org/10.1080/10934529.2015.1030283

Lintereur

P.A.

Duranceau

S.J.

Taylor

J.S.

Stone

E.D.

. “Sodium Silicate Impacts on Lead Release in a Blended Potable Water Distribution System”. Desalination Water Treat. 2010. 16(1–3): 427–438. https://doi.org/10.5004/dwt.2010.1477

10.

Trueman

B.F.

Munoz

Locsin

J.M.

Gagnon

G.A.

. “Impact of Sodium Silicate on Lead Release and Colloid Size Distributions in Drinking Water”. Water Res. 2021. 190(15): 116709. https://doi.org/10.1016/j.watres.2020.116709

11.

Aghasadeghi

Peldszus

Trueman

B.F.

Mishrra

, et al. “Pilot-Scale Comparison of Sodium Silicates, Orthophosphate and pH Adjustment to Reduce Lead Release from Lead Service Lines”. Water Res. 2021. 195(1): 116955. https://doi.org/10.1016/j.watres.2021.116955

12.

Mishrra

Wang

Sidorkiewicz

Giammar

D.E.

. “Effect of Sodium Silicate on Lead Release from Lead Service Lines”. Water Res. 2021. 188(1): 116485. https://doi.org/10.1016/j.watres.2020.116485

13.

Guo

. Corrosion Scale Characteristics and Lead Oxide Dissolution in Chloraminated Water. [Doctor of Philosophy in Chemical and Biochemical Engineering Dissertation]. London, Ontario, Canada: University of Western Ontario, 2018.

14.

Kim

E.J.

Herrera

J.E.

. “Characteristics of Lead Corrosion Scales Formed During Drinking Water Distribution and Their Potential Influence on the Release of Lead and Other Contaminants”. Environ. Sci. Technol. 2010. 44(16): 6054–6061. https://doi.org/10.1021/es101328u

15.

Xie

Wang

Singhal

Giammar

D.E.

. “Effects of pH and Carbonate Concentration on Dissolution Rates of the Lead Corrosion Product PbO₂”. Environ. Sci. Technol. 2010. 44(3): 1093–1099. https://doi.org/10.1021/es9026198

16.

Lytle

D.A.

Schock

M.R.

. “Formation of Pb(IV) Oxides in Chlorinated Water”. J. Am. Water Works Assoc. 2005. 97(11): 102–114. https://doi.org/10.1002/j.1551-8833.2005.tb07523.x

17.

Guo

Robinson

Herrera

J.E.

. “Role of Pb(II) Defects in the Mechanism of Dissolution of Plattnerite (β-PbO₂) in Water Under Depleting Chlorine Conditions”. Environ. Sci. Technol. 2014. 48(21): 12525–12532. https://doi.org/10.1021/es502133k

18.

Guo

Robinson

Herrera

J.E.

. “Mechanism of Dissolution of Minium (Pb₃O₄) in Water Under Depleting Chlorine Conditions”. Corros. Sci. 2016. 103: 42–49. https://doi.org/10.1016/j.corsci.2015.10.042

19.

Schock

M.R.

Scheckel

K.G.

Desantis

M.K.

Gerke

. ““Mode of Occurrence, Treatment and Monitoring Significance of Tetravalent Lead”. Paper presented at: AWWA Water Quality Technology Conference. Quebec City”. Canada. November 2005. 6–10.

20.

Lytle

D.A.

Schock

M.R.

Scheckel

. “The Inhibition of Pb(IV) Oxide Formation in Chlorinated Water by Orthophosphate”. Environ. Sci. Technol. 2009. 43(17): 6624–6631. https://doi.org/10.1021/es900399m

21.

Bilinski

Schindler

. “Solubility and Equilibrium Constants of Lead in Carbonate Solutions (25°C, I = 0.3 mol dm⁻³)”. Geochim. Cosmochim. Acta. 1982. 46(6): 921–928. https://doi.org/10.1016/0016-7037(82)90048-5

22.

Xie

Wang

Giammar

D.E.

. “Impact of Chlorine Disinfectants on Dissolution of the Lead Corrosion Product PbO₂”. Environ. Sci. Technol. 2010. 44(18): 7082–7088. https://doi.org/10.1021/es1016763

23.

Switzer

J.A.

Rajasekharan

V.V.

Boonsalee

Kulp

E.A.

Bohannan

E.W.

. “Evidence That Monochloramine Disinfectant Could Lead to Elevated Pb Levels in Drinking Water”. Environ. Sci. Technol. 2006. 40(10): 3384–3387. https://doi.org/10.1021/es052411r

24.

Lin

Y.P.

Valentine

R.L.

. “Release of Pb(II) from Monochloramine-Mediated Reduction of Lead Oxide (PbO₂)”. Environ. Sci. Technol. 2008. 42(24): 9137–9143. https://doi.org/10.1021/es801037n

25.

Kim

E.J.

Herrera

J.E.

Huggins

Braam

Koshowski

. “Effect of pH on the Concentrations of Lead and Trace Contaminants in Drinking Water: A Combined Batch, Pipe Loop, and Sentinel Home Study”. Water Res. 2011. 45(9): 2763–2774. https://doi.org/10.1016/j.watres.2011.02.023

26.

Snoeyink

V.L.

Schock

M.R.

Sarin

Wang

, et al. “Aluminium-Containing Scales in Water Distribution Systems: Prevalence and Composition”. J. Water Supp.: Res. Technol.–AQUA. 2003. 52(7): 455–474. https://doi.org/10.2166/aqua.2003.0042

27.

Bae

Mishrra

Shi

Giammar

D.E.

. “Effect of Aluminum on Lead Release to Drinking Water from Scales of Corrosion Products”. Environ. Sci. Technol. 2020. 54(10): 6142–6151. https://doi.org/10.1021/acs.est.0c00738

28.

Arai

Sparks

D.L.

Davis

J.A.

. “Arsenate Adsorption Mechanisms at the Allophane–Water Interface”. Environ. Sci. Technol. 2005. 39(8): 2537–2544. https://doi.org/10.1021/es0486770

29.

Opiso

Sato

Yoneda

. “Adsorption and Co-Precipitation Behavior of Arsenate, Chromate, Selenate, and Boric Acid with Synthetic Allophane-Like Materials”. J. Hazard. Mater. 2009. 170(1): 79–86. https://doi.org/10.1016/j.jhazmat.2009.05.001

30.

Farmer

V.C.

. “The Solubility Characteristics of Organic Al Precipitates in Allophanic Bs Horizons Could Mimic Those of Al Hydroxides and Indicate the Solubility of Associated Allophanes”. Eur. J. Soil Sci. 1999. 50(4): 713–715. https://doi.org/10.1046/j.1365-2389.1999.00271.x

31.

Parfitt

R.L.

Kimble

J.M.

. “Conditions for Formation of Allophane in Soils”. Soil Sci. Am. J. 1989. 53(3): 971–977. https://doi.org/10.2136/sssaj1989.03615995005300030057x

32.

Berggren

Mulder

. “The Role of Organic Matter in Controlling Aluminum Solubility in Acidic Mineral Soil Horizons”. Geochim. Cosmochim. Acta. 1995. 59(20): 4167–4180. https://doi.org/10.1016/0016-7037(95)94443-j

33.

Harsh

Chorover

Nizeyimana

. “Allophane and Imogolite”. In: Dixon

J.B.

Wang

M.K.

Kampf

Schulze

D.G.

, editors. Soil Mineralogy with Environmental Applications. Madison, Wisconsin: Soil Science Society of America, 2018. Pp. 291–322.

34.

Trueman

B.F.

Rahman

M.S.

Gao

, et al. “Understanding the Impacts of Sodium Silicate on Water Quality and Iron Oxide Particles”. Environ. Sci.: Water Res. Technol. 2019. 5(8): 1360–1370. https://doi.org/10.1039/C9EW00257J

35.

Rudolph

W.W.

Mason

Pye

C.C.

. “Aluminium(III) Hydration in Aqueous Solution. a Raman Spectroscopic Investigation and an Ab Initio Molecular Orbital Study of Aluminium(III) Water Clusters”. Phys. Chem. Chem. Phys. 2000. 2(22): 5030–5040. https://doi.org/10.1039/b003764h

36.

Bertsch

P.M.

Parker

D.R.

. The Environmental Chemistry of Aluminum. Boca Raton, Florida: Lewis Publishers, 1996.

37.

Huang

P.M.

Wang

M.K.

Kämpf

Schulze

D.G.

. “Aluminum Hydroxides”. In: Dixon

J.B.

Wang

M.K.

Kampf

Schulze

D.G.

, editors. Soil Mineralogy with Environmental Applications. Madison, Wisconsin: Soil Science Society of America, 2018. Pp. 261–289.

38.

Lui

Korshin

G.V.

Ferguson

J.F.

. “Investigation of the Kinetics and Mechanisms of the Oxidation of Cerussite and Hydrocerussite by Chlorine”. Environ. Sci. Technol. 2008. 42(9): 3241–3247. https://doi.org/10.1021/es7024406

39.

RRUFF Project. “Hydrocerussite R160062”. https://rruff.info/hydrocerussite/R160062 [accessed Sep 24 2024].

40.

Creton

Bougeard

Smirnov

K.S.

Guilment

Poncelet

. “Structural Model and Computer Modeling Study of Allophane”. J. Phys. Chem. C. 2008. 112(2): 358–364. https://doi.org/10.1021/jp0738412

41.

Theng

B.K.G.

. “The Clay Minerals”. Develop. Clay Sci. 2012. 4: 3–45. https://doi.org/10.1016/B978-0-444-53354-8.00001-3

42.

National Institute of Standards and Technology (NIST). “NIST X-ray Photoelectron Spectroscopy Database”. 2012. https://srdata.nist.gov/xps/ [accessed Sep 24 2024].

43.

Thomas

J.M.

Tricker

M.J.

. “Electronic Structure of the Oxides of Lead Part 2: An XPS Study of Bulk Rhombic PbO, Tetragonal PbO, β-PbO₂, and Pb₃O₄”. J. Chem. Soc., Faraday Trans. 1974. 2(71): 329–336. https://doi.org/10.1039/F29757100329

44.

Childs

C.W.

Inoue

Seyama

Soma

, et al. “X-ray Photoelectron Spectroscopic Characterization of Silica Springs Allophane”. Clay Miner. 1997. 32(4): 565–572. https://doi.org/10.1180/claymin.1997.032.4.07

Supplementary Material

Please find the following supplemental material available below.

For Open Access articles published under a Creative Commons License, all supplemental material carries the same license as the article it is associated with.

For non-Open Access articles published, all supplemental material carries a non-exclusive license, and permission requests for re-use of supplemental material or any part of supplemental material shall be sent directly to the copyright owner as specified in the copyright notice associated with the article.

0.00 MB

0.96 MB