Replacing polyphosphate with silicate to solve lead, copper, and source water iron problems

In 199, the town of Hopkinton, Mass., instituted a polyphosphate sequestrant feed in two BY MICHAEL R. SCHOCK, DARREN A. LYTLE, ANNE M. SANDVIG, JONATHAN CLEMENT, AND STEPHEN M. HARMON of its five wells to address colored water complaints caused by source water iron and manganese. Failure to meet the Lead and Copper Rule (LCR) action levels and continuing intermittent red water problems prompted the water utility to seek an alternative treatment strategy. After analyses of the water quality of the different wells, the scales on some lead service lines, the operational patterns, the location of lead service lines, and the distribution system configuration, a combination of sodium silicate (with chlorination) and sodium hydroxide addition at different wells was investigated. At study monitoring sites, an initial silicate dose of 2 3 mg/l elevated the ph from 6.3 to 7.1 and immediately resulted in a % reduction in lead levels and an 87% reduction in copper levels. An increase to a silicate dose of 4 mg/l elevated the ph to 7. and produced greater reductions in lead and copper. The treatment change reduced 9th percentile lead and copper levels by at least 9%, enabling compliance with the LCR. The aesthetic quality of the drinking water after treatment was equal or superior to the quality before treatment. The study showed that for many small and medium water systems with multiple wells and entry points, simultaneous overall corrosion control and sequestration of iron and manganese is possible, and it avoids the adverse drinking water/wastewater consequences associated with polyphosphate addition. Replacing polyphosphate with silicate to solve lead, copper, and source water iron problems H opkinton, Mass., has a medium-size water system as defined in the Lead and Copper Rule (LCR) classification system (USEPA, 1994, 1992, 1991a, 1991b). It is located about 3 mi (48 km) west of Boston and serves approximately 7, individuals. Water is obtained from five gravel-packed production wells in an unconsolidated sand and gravel aquifer. The wells have individual pumping stations and are near a surface water source (Whitehall Reservoir). At the time this article was written, however, the wells had not been classified as being under the influence of surface water. Figure 1 is a schematic of the water system layout, showing well locations. Wells 4 and combine into a single pipe before entering the distribution system and produce approximately 3% of the system s water (Clement et al, 1998). 84 NOVEMBER 2 JOURNAL AWWA 97:11 PEER-REVIEWED SCHOCK ET AL

Wells 1 3 account for the remainder of the water demand, although well 3 is rarely used and supplies a very small percentage of the produced water (<6%). Cement-lined ductile iron comprises approximately 8% of the 68 mi (19 km) of pipeline in the distribution system, with the remainder being either unlined cast iron or asbestos cement. There are believed to be between 2 and 3 lead (Pb) service lines in the system, located primarily in the area fed by wells 1 and 2. Historically, comprehensive and systematic source water quality monitoring has not been done. Table 1, however, shows data from some monitoring conducted in 199 (Harmon, 1998). Results of subsequent monitoring of a subset of key water chemistry parameters during the course of this research are consistent with the characterizations seen in Table 1, with expected small fluctuations in ph and alkalinity. During the course of the study, total alkalinity was generally closer to 2 mg/l calcium carbonate (CaCO 3 ), corresponding to a computed total inorganic carbon concentration of approximately 13 18 mg/l carbon (C). In 199, Hopkinton instituted a polyphosphate sequestrant 1 feed in wells 4 and to try to remedy colored water complaints resulting from elevated source water iron (Fe) and manganese (Mn). The target dose was approximately 2 mg/l as product. The utility also attempted to institute a flushing program to prevent accumulation of precipitated iron in the mains. Well 4 (high Fe and Mn) Well (high Fe and Mn) Frequent complaints about red water continued, and in October 1992, LCR monitoring revealed high lead and copper (Cu) levels at many taps, resulting in reported 9th percentile concentrations of.77 mg/l Pb and.87 mg/l Cu. Treatment options were considered limited. Simple elevation of ph would further degrade sequestration performance and reduce the inhibition of copper corrosion by the small amount of orthophosphate present from polyphosphate reversion (Schock et al, 199a, 199b). Therefore, Hopkinton utility managers agreed FIGURE 1 Schematic of well locations and treatment zones for Hopkinton, Mass., system Fe iron, Mn manganese TABLE 1 Sequestrationtreated section General water quality data for Hopkinton, Mass., wells* Well # Well 1 Well 2 (high Mn) Well 3 Parameter 1 2 3 4 Ammonia <. <. <. <..3 Calcium 9. 7. 9..1 6.7 Carbon 32 18 36 Copper <.4 <.4 <.4 <.4 <.4 Iron <.1 <.1 <.1 2.19 3.34 Lead <.1 <.1 <.1.3 <.1 Magnesium 2.2 1.6 2.1 1. 1.8 Manganese <.1.8 <.1.14.11 Nitrogen dioxide 1.99.73 1.81.68 <.2 ph 6.3 6.1 6.2 6. 6.1 Potassium 1.6 1.6 1.6 1. 1.1 Silicate 14 14 14 13 1 Sodium 16 1 1 Sulfate 9 9 9 2 Total alkalinity mg/l as CaCO 3 34 34 34 34 1 *Data collected November 9, 199 All chemicals measured in mg/l to participate in a US Environmental Protection Agency- (USEPA-) sponsored two-year study of innovative corrosion control approaches for smaller water systems. The study was conducted by Black & Veatch Corp. engineers under the guidance of the New England Water Works Association (NEWWA; Clement et al, 1998). STUDY DESIGN After analyzing the water quality of the different wells and the scales on some lead service lines and reviewing SCHOCK ET AL PEER-REVIEWED 97:11 JOURNAL AWWA NOVEMBER 2 8

FIGURE 2 Effect of different silica-to-sodium-oxide ratios in the silicate chemical formulation on ph SiO 2 mg/l 12 96 72 48 24 Ratio of Silicate to Sodium Oxide 3.22 2.8 2. 1.8 1. DIC = 13 mg/l C 6 7 8 9 1 Desired ph C carbon, DIC dissolved inorganic carbon, SiO 2 silicate Type N is most common, with a ratio of 3.22. This study used a product with a ratio of 1.6, the lowest that is widely available on a commercial basis. operational patterns, the distribution system configuration, and the location of lead service lines, the project team decided to investigate a combination of chlorination and sodium silicate (Na 2 SIO 3 ) addition in those wells containing elevated levels of iron and manganese (wells 2, 4, and ). The sodium silicate solution would provide sequestration for the iron (and to a lesser extent, the manganese) and would also increase the ph. The copper and lead problems would therefore be simultaneously addressed by either the direct or indirect action of the silicate addition. Well 1 contained neither significant manganese nor iron, so only chlorine disinfection and addition of sodium hydroxide (NaOH; caustic) for ph adjustment was configured for that location. The dosage of sodium hydroxide was adjusted to make sure the ph of well 1 matched that of the wells receiving sodium silicate treatment (wells 2, 4, and ). The work of Robinson and colleagues (1987) suggests that silicate chemicals sequester iron and manganese by a colloidal dispersion mechanism. Deng (1997) found that Si(IV) adsorbed to Fe(III) hydroxide at an early stage of precipitation. Si(IV) surface complexes on Fe(III) oxyhydroxides can form (Stumm & Morgan, 1981; Sigg & Stumm, 198), and as a result, silicate surface species have been shown to inhibit crystallization of iron solids (Schwertmann & Thalmann, 1976). USEPA research (unpublished) shows that adsorbed silicate also decreases the surface charge of iron colloids under some conditions, which produces stable colloidal iron suspensions. The presence of major multivalent cations such as calcium, however, can destabilize charged colloids. This could cause problems in sequestering iron (Robinson et al, 1987). The role of silicate in lead and copper corrosion control has been and still is somewhat uncertain. Although some have suggested that a thin silicate coating forms and acts as a protective diffusion barrier (Stericker, 194), no strong body of analytical evidence to support this claim exists. Because sodium silicate chemicals are basic, they do increase ph in sufficient dosages, which is usually beneficial in terms of reducing lead and copper solubility. Regardless of the exact mechanism of sequestration and corrosion inhibition that the silicate provides, several unpublished studies have documented improvements in lead, copper, and iron release, singly or in combination. Because the silicate sequestration operates on the oxidized forms of the metals (LaRosa-Thompson et al, 1997; Robinson et al, 1992; Robinson et al, 1987; Dart & Foley, 1972, 197; Lehrman & Shuldener, 191), the proposed treatment would also be compatible with maintaining a disinfectant residual in the water and concurrently inhibiting copper(ii) corrosion (Schock et al, 199a, 199b). To gather accurate baseline data and to gain a good understanding of the responses of different parts of the system to the treatments under investigation, the project team monitored 22 tap sites. The sites were geographically distributed so that 11 sites were located in a section predominantly receiving water from wells 1 and 2, and another 11 were located in the part of the system fed by wells 4 and. Water quality sampling and operational analysis confirmed that there were essentially two separate hydraulic zones in the system. Water system personnel collected monthly samples at these 22 sites. Four of the sampling sites in the section served by wells 1 and 2 were from locations containing lead service lines. The remaining service lines were copper. For the part of the system fed by wells 4 and, five months of data were collected before sodium silicate treatment was implemented, and 1 months of data were collected after treatment. Almost nine months of data were collected in the distribution system zone fed by wells 1 and 2 before treatment was implemented. After 86 NOVEMBER 2 JOURNAL AWWA 97:11 PEER-REVIEWED SCHOCK ET AL

treatment, 12 months of data were collected in this zone. SAMPLING AND ANALYTICAL METHODOLOGY Samples collected at the distribution system entry points included volumes of chemicals used, water quality data, operational problems or interruptions, volume of water pumped, and times of pumping cycles for each well. Water samples collected at the 22 distribution system sites included consecutive 1-L and -ml samples collected after the water had been allowed to stagnate for 6 1 h. The 1-L sample was preserved for metals analyses (lead, copper, iron, manganese, zinc, calcium, and magnesium), and the -ml sample was analyzed for other water quality parameters (alkalinity, total phosphate and orthophosphate, silica, turbidity, and chlorine residual). At the four sites with lead service lines, a third consecutive 1- L sample was collected, which was acid preserved for metals analyses. The project team conducted analyses of ph, temperature, and conductivity onsite by allowing the sample water to flow into an airtight chamber. The probe was inserted into a hole in the top of the chamber. The chamber was rinsed thoroughly with the water to be sampled before measurement. Immediately after the samples were collected, alkalinity was analyzed. Metals analysis samples were collected in bottles precleaned with nitric acid, and those samples were preserved by adding an appropriate volume of concentrated nitric acid to achieve a ph of <2. All analytical procedures followed specifications in Standard Methods (1992) or USEPA methods manuals (199, 1983). Calibrations were performed immediately before each field ph and conductivity analysis, and they were rechecked immediately after each analysis to verify the absence of errors or instrumental drift. For each ph calibration, buffers of ph 4, 7, and 1 were used. Selection of silicate formulation. The initial distribution system ph target ranged from 7.4 to 7.6. The target dose of sodium silicate was 2 mg/l, on the basis of the amount needed for proper FIGURE 3 Solubility diagram for lead showing different theoretical metastable solid-phase stability fields Pb mg/l 1 1.1.1.1.1 PbCO 3 Pb 3 (CO 3 ) 2 (OH) 2 Pb 3 (PO 4 ) 2 Pb (PO 4 ) 3 CI Pb (PO 4 ) 3 OH 6 7 8 9 1 ph Pb lead, PbCO 3 lead carbonate, Pb 3 (CO 3 ) 2 (OH) 2 hydrocerrusite, Pb 3 (PO4) 2 lead phosphate, Pb (PO 4 ) 3 OH hydroxypyromorphite, Pb (PO 4 ) 3 Cl chloropyromorphite Diagram assumes 2 C, an ionic strength of., and dissolved inorganic carbonate, orthophosphate, sulfate, and chloride concentrations of 1 mg/l C, 1 mg/l PO 4, and mg/l sulfate and chloride, respectively. Polyphosphate complexation not included. FIGURE 4 X-ray diffraction pattern comparison for the outer scale layer of copper pipe specimens before and after silicate treatment Intensity counts 1, 7 2 1, 7 2 Presilicate Postsilicate 1-399> Malachite, synthetic - Cu 2 +2 CO 3 (OH) 2 1 2 3 4 6 7 8 2 θ degrees SCHOCK ET AL PEER-REVIEWED 97:11 JOURNAL AWWA NOVEMBER 2 87

This scanning electron microscopy secondary emission micrograph of a Hopkinton, Mass., lead service line specimen shows a fibrous mat of chloropyromorphite intermixed with amorphous iron, overlaying a surface composed predominantly of lead carbonate. sequestration per Robinson and colleagues (1992, 1987). Commercially available sodium silicate solutions are available with SiO 2 :Na 2 O ratios ranging from 1.6 to 3.22. Figure 2 shows the effects of various dosages of these mixtures on the ph of the Hopkinton water. For this study, a product with the ratio of 1.6 was chosen. 2 Lead and copper pipe analyses. To attempt to understand the mechanism producing the high lead levels, the project team analyzed two pieces of lead service line at the beginning of the study. The exact locations of origin and ages of the pipe specimens were not known, but they were from an area of the system exposed to the phosphate sequestrant chemical for approximately two years. One pipe section had an inside diameter of 1 in. (2 mm); another was.7 in. (2 mm). The pieces of pipe were cut and subjected to a variety of chemical and mineralogical analyses, including optical microscopy, polarized and reflected light microscopy, x-ray diffraction (XRD), electron microprobe analysis, and scanning electron microscopy (SEM). Details of sample preparation and analysis are given elsewhere (Harmon, 1998). A layered structure was apparent on both specimens (Harmon, 1998). On the outermost layer exposed to the water was a combination of a discontinuous reddish amorphous phase and the very insoluble lead phosphate mineral chloropyromorphite [Pb (PO 4 ) 3 Cl]. Beneath this outermost layer was a continuous white layer consisting almost entirely of the mineral lead carbonate or cerussite (PbCO 3 ). Below the cerussite layer was a gray-white layer containing a substantial quantity of litharge (PbO). The presence of this highly soluble PbO phase indicates that this layer exists below the depth penetrated by flowing water. The presence of the mineral chloropyromorphite was very surprising because it had not been previously reported in lead pipe scales to the knowledge of the authors at the time of the study. Further, many researchers have believed that chloropyromorphite is of such low solubility as to preclude the levels of dissolved lead reported in systems dosed with orthophosphate (Schock et al, 1996; Schock, 1989; Schock & Wagner, 198; Hunt & Creasey, 198). For example, Figure 3 shows a solubility diagram for Pb(II) with dissolved inorganic carbon (DIC) and orthophosphate residual conditions similar to Hopkinton water, illustrating the solubilities of different possible solubility-controlling minerals at various ph values. The solubilities shown are not precise quantitative estimates. Lead levels predicted are likely lower than true solubility because aqueous polyphosphate species form complexes with lead that are not accommodated in the calculations (species and equilibrium constants cannot be identified). However, lead solids become less soluble over time because of increased crystallinity and structure ordering, making the quantitative prediction of the solubility of a lead mineral at different development times quite problematic. In other studies, polyphosphate has been shown to be an aggressive agent in lead control treatment (Edwards & McNeill, 22; Edwards et al, 21; Cantor et al, 2; Schock et al, 1996; Dodrill & Edwards, 199). Figure 3, however, clearly shows that if chloropyromorphite comprised the passivating film, lead solubility would be considerably lower than the solubility observed in numerous water systems in which other lead orthophosphates of higher solubility such as hydroxypyromorphite [Pb (PO 4 ) 3 OH] and lead phosphate [Pb 3 (PO 4 ) 2 ] have been identified as major scale components (Schock et al, 1996). FIGURE Monthly average ph values for the section fed by wells 1 and 2 ph 9 8 7 6 Before treatment 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 Month 88 NOVEMBER 2 JOURNAL AWWA 97:11 PEER-REVIEWED SCHOCK ET AL

A combination of SEM and electron microprobe analysis clarified the interpretation of the XRD data, although this technique is not spatially precise (Harmon, 1998). The photo on page 88 shows an example SEM secondary emission micrograph of one of many areas examined on the surface of both pipe specimens. Elemental analysis with the electron microprobe verified that the needlelike and stringlike crystals are the chloropyromorphite associated with an amorphous iron oxide material. The primary surface of the pipe is covered with a uniform coating of lead carbonate underneath fibrous clumps of iron hydroxide and chloropyromorphite. The presence of lead carbonate is consistent with solubility models given the ph and DIC of the water to which the pipe was originally exposed. The morphology of the chloropyromorphite is clearly not protective and does not inhibit lead release (Harmon, 1998). Under the other water chemistry conditions (ph, alkalinity), the solubility of lead is high enough to promote substantial lead release. The occurrence of the discontinuous amorphous iron deposits are also consistent with the observation of incomplete iron sequestration during the period when the polyphosphate was added. This absence of a continuous coating of iron oxides or oxyhydroxides is consistent with the general solubilizing or sequestration ability of polyphosphate ligands. This chloropyromorphite/iron occurrence is strikingly similar to observations of the growth of the mineral chloropyromorphite onto existing ferric oxide (goethite) surfaces reported in studies of lead immobilization in soils by calcium orthophosphate (hydroxyapatite) treatment (Zhang et al, 1997). One domestic.-in. (1-mm) inside diameter copper pipe specimen removed before the study and one removed after the study were available for examination. On the pipe specimen removed before the silicate addition, XRD analysis revealed both a loose surficial deposit and harder inner scale composed almost entirely of malachite (Cu 2 (OH) 2 CO 3 ) with a lesser fraction of cuprite (Cu 2 O) as the only detectable crystalline solids. Peak broadening of the malachite component indicates a solid of small crystallite size and large surface area. The levels of soluble copper observed in the study drastically exceed levels that would be in equilibrium with crystalline malachite in a pipe specimen with aged scale (Schock et al, 199a, 199b). Elemental analysis of the outer scale deposit by energy dispersive x-ray analysis (EDXA) on the SEM showed that the scale was somewhat inhomogeneous, with iron approximately 4 17% by weight, phosphorous approximately.6 1.3%, and silica approximately.6 1.6%. Calcium and sulfur were each less than.2%. The rest of the scale was composed of copper, oxygen, and carbon, as would be expected, and hydrogen (which would be present in malachite) is undetectable by EDXA. Coulometric inorganic carbon analysis (ASTM, 1997) showed 16% as carbon dioxide (CO 2 ), a little lower than the theoretical value for malachite of 19.8%, suggesting FIGURE 6 Monthly average ph values from the part of the system fed by wells 4 and ph 9 8 7 6 Before treatment 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 7 6 4 3 2 1 Before treatment Action level Month Silicate feed rates 2 3 mg/l 4 mg/l FIGURE 7 Monthly lead and copper concentrations in the part of the system fed by wells 1 and 2 9th Percentile Lead mg/l 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 9th Percentile Copper mg/l 8 7 6 4 3 2 Before treatment Action level Month SCHOCK ET AL PEER-REVIEWED 97:11 JOURNAL AWWA NOVEMBER 2 89

FIGURE 8 Monthly lead and copper concentrations for the part of the system fed by wells 4 and 9th Percentile Lead mg/l 9th Percentile Copper mg/l 7 6 4 3 2 1 8 7 6 4 3 2 1 Before treatment 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 1/93 2/94 4/94 6/94 8/94 1/94 2/9 4/9 6/9 Lead mg/l Before treatment Silicate feed rates 2 3 mg/l 4 mg/l Silicate feed rates Action level 2 3 mg/l 4 mg/l Month Action level FIGURE 9 Regulatory monitoring data for lead before (round 1) and after (remaining rounds) silicate and ph adjustment treatment Round 1, 8/92 Round 3, 8/99 12 11 1 9 8 7 6 4 3 2 1 Round 2, 3/98 Round 4, 6/ 1 2 3 4 6 7 8 9 1 Sites % some additional amorphous material or minor concentrations of noncarbonate minerals not obvious in the XRD patterns. The XRD results after the silicate treatment were remarkable (Figure 4). Coulometric inorganic carbon analysis of two subsamples of the posttreatment scale yielded.2 and.3% as CO 2. After silicate treatment, malachite was no longer present in the scale. The XRD and total inorganic carbon data therefore suggest that the silicate facilitated the conversion of malachite to what is most likely an amorphous cupric silicate phase. Effects of initial polyphosphate treatment. From 199 through 1993, polyphosphate treatment was implemented at Hopkinton. Water quality data were collected in the distribution system monthly for three months before the implementation of silicate treatment, while polyphosphate was being used. Polyphosphate concentrations were estimated as the difference between the analytical values for total phosphate and orthophosphate. Polyphosphate concentrations varied between 1.2 and.8 mg/l PO 4 and averaged 3.4 mg/l PO 4. Orthophosphate concentrations ranged from.3 to 2.8 mg/l PO 4, with an average of approximately 1.2 mg/l PO 4. The initial distribution system ph averaged 6.1 with 8% of the values being between.9 and 6.4. Detailed examination of the first-draw copper data showed that they followed the expected trends for some inhibition of copper by orthophosphate at the lower ph range (Schock et al, 199a, 199b), particularly when the orthophosphate concentration exceeded 1 mg/l PO 4. Elevated lead levels apparently resulted from the sequestration of the polyphosphate and the ineffectiveness of orthophosphate at the low ph. Effects of silicate and ph adjustment. Figures and 6 show the effect of the silicate and caustic treatment on the ph throughout the system. Figure represents samples taken in the portion of the distribution system served by wells 1 (caustic) and 2 (silicate), and Figure 6 represents samples taken in the portion of the distribution system served by wells 4 and (silicate). Notably, and consistent with many other studies, the distribution system ph took several months to stabilize. This is at least partly the result of buffering by existing scale minerals throughout the distribution system. The silicate dosage was increased at wells 4 and in December 1994 (from 2 3 mg/l SiO 2 to approximately 4 mg/l SiO 2 ), to try to get closer to the desired final ph to further reduce copper and lead levels. Figures 7 and 8 show the responses of copper and lead levels to the treatment during the course of the study. Typical of treatment changes, spikes of elevated lead concentrations were observed initially after the silicate addition was initiated. Taken as a whole, these data appear consistent with the slow formation of protective films, which is the postulated mechanism of corrosion control with silicates (LaRosa-Thompson et al, 1997; AwwaRF and DVGW Forschungsstelle-TZW, 1996). 9 NOVEMBER 2 JOURNAL AWWA 97:11 PEER-REVIEWED SCHOCK ET AL

Although the regulatory monitoring sites do not correspond exactly to the study monitoring sites, the overall effectiveness of the treatment program can be seen in the results of the posttreatment regulatory monitoring. Figures 9 and 1 show the regulatory monitoring data for lead and copper, in which overall and 9th percentile levels have stabilized or continued to decrease since the mixed silicate and complementary ph adjustment treatment became operational. For lead, the 9th percentile level declined from.77 to.2 mg/l, a 97% reduction. Compared with the initial 1992 monitoring, the 2 data show that the treatment has lowered the 9th percentile copper level from.87 to.27 mg/l a reduction of 9%. Aesthetic and distribution system effects of treatment. The change from polyphosphate to the moderate silicate dose (~2 3 mg/l) caused higher ph and possibly lower turbidity levels. Figure 11 shows the monthly average iron, silica, and turbidity levels measured in the portion of the distribution system fed by wells 4 and. Color was not systematically monitored throughout the study, but only two samples exceeded 1 cu after silicate dosage began. Iron concentrations increased following silicate addition. One interpretation is that sequestering was improved, with more iron remaining complexed with silicate rather than precipitating out of solution to create turbidity and color. However, the increase in iron may also result from changes in the use of a particular well or other seasonal FIGURE 1 Regulatory monitoring data for copper before (round 1) and after (remaining rounds) silicate and ph adjustment treatment Copper mg/l 1 9 8 7 6 4 3 2 1 Round 1, 8/92 Round 3, 8/99 Round 2, 3/98 Round 4, 6/ Action level 1 2 3 4 6 7 8 9 1 Site % FIGURE 11 Relationships among total silica, iron, and turbidity levels for the high-iron portion of the Hopkinton, Mass., distribution system (fed by wells 4 and ) Iron mg/l and Turbidity ntu 3. 2. 2. 1. 1.. Iron Turbidity Silica Before treatment. 1/93 11/93 1/94 2/94 3/94 4/94 /94 6/94 7/94 8/94 9/94 Month 1/94 11/94 1/9 2/9 3/9 4/9 /9 6/9 7/9 effects. A seasonal trend is suggested in Figure 11, and because of the limited baseline time period, no clear way to evaluate that possibility exists. The treatment produced acceptable water quality (i.e., color <1 cu and turbidity <1 ntu) at the higher ph conditions necessary for copper and lead control and produced more consistent sequestration performance. Manganese can contribute to discoloration and staining; however, the sequestration ability of silicate for manganese is known to be somewhat poor relative to its ability toward iron. Research has indicated that chlorine oxidation of manganese is slower at ph conditions below 8. In addition, chlorine doses well above the stoichiometric quantities (greater than 3 mg/l for Hopkinton wells 4 and ) would be needed for complete oxidation, assuming all original manganese was soluble (Knocke et al, 1987). Investigation of relationships between color and both iron and manganese concentrations in this study 7 6 4 3 2 1 Silica mg/l SCHOCK ET AL PEER-REVIEWED 97:11 JOURNAL AWWA NOVEMBER 2 91

showed little relationship with manganese. As a result of these factors, no problem with manganese was observed or is expected with this water supply. In addition to sampling throughout the distribution system, samples were also collected from hot and cold water taps in five homes to evaluate the effect of temperature on the effectiveness of silicate as a sequestering agent. Monthly samples were collected four times and analyzed for ph, turbidity, iron, color, and manganese. Comparisons of results from the hot and cold taps did not reveal any significant differences in the averages of any measured parameter. However, a paired t-test showed that at 9% confidence (p <.) the color levels and iron concentrations in the hot and cold water samples were consistently different. The hot water samples had slightly higher color but lower iron levels than the cold water samples. These differences did not indicate a general trend that the elevated temperatures adversely affected the efficacy of silicate to sequester the iron adequately. With only four samples, however, there were insufficient data, and more research into temperature effects on sequestration is clearly warranted. CONCLUSIONS This article shows that simultaneous overall corrosion control and sequestration of iron and small concentrations of manganese are possible for many small water systems with multiple wells and entry points. The aesthetic quality of the drinking water after silicate treatment was equal or superior to the water s quality before the treatment. The higher ph and silicate combination is also more compatible with protecting asbestos cement and cement-lined pipes than polyphosphate treatment at a lower ph. As a result of the treatment change, the utility was able to meet the regulatory action level for both lead and copper and move into a reduced monitoring category. The trends in the data also show slow and continual long-term decreases in corrosion and metal release from plumbing materials, which helps to support observations that many short-term corrosion testing experiments using new coupons or pipe sections may give inadequate prediction of long-term performance of silicate treatment in real distribution systems. This investigation of sodium silicate and sodium hydroxide treatment was completed in the summer of 199. From 199 through 21, the town of Hopkinton continued to implement this treatment and was able to maintain compliance with the LCR. In early 22, the town switched to a purchased groundwater supply that was treated at an iron and manganese removal plant. At the time this article was written, the wells described in this article were used as an emergency supply only. ACKNOWLEDGMENT The authors thank Ray Raposa of the NEWWA for his assistance and cooperation with the project and for REFERENCES ASTM (American Society for Testing and Materials), 1997. Standard Test Method for Total and Dissolved Carbon Dioxide in Water. Volume 11.1, D 13-92 (Reapproved 1996). ASTM International, West Conshohocken, Pa. AwwaRF (AWWA Research Foundation) and DVGW Forschungsstelle- TZW, 1996 (2nd ed.). Internal Corrosion of Water Distribution Systems. AwwaRF, Denver. Cantor, A.F. et al, 2. Use of Polyphosphate in Corrosion Control. Jour. AWWA, 92:2:9. Clement, J. et al, 1998. Evaluation of Lead and Copper Corrosion Control Treatment Strategies in Small Systems. Internal Report, National Risk Management Research Laboratory, Office of Research and Development, Cincinnati, Ohio. Dart, F.J. & Foley, P.D., 1972. Silicate as Fe, Mn Deposition Preventative in Distribution Systems. Jour. AWWA, 64:4:244. Dart, F.J. & Foley, P.D., 197. 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Using Alternate Oxidants to Remove Dissolved Manganese From Waters Laden With Organics. Jour. AWWA, 79:3:7. LaRosa-Thompson, J. et al, 1997. Sodium Silicate Corrosion Inhibitors: Issues of Effectiveness and Mechanism, Proc. AWWA WQTC, Denver. Lehrman, L. & Shuldener, H.L., 191. The Role of Sodium Silicate in Inhibiting Corrosion by Film Formation on Water Piping. Jour. AWWA, 43:17. Robinson, R.B.; Reed, G.D.; & Frazier, B., 1992. Iron and Manganese Sequestration Facilities Using Sodium Silicate. Jour. AWWA, 84:2:77. Robinson, R.B.; Minear, R.A.; & Holden, J.M., 1987. Effects of Several Ions on Iron Treatment by Sodium Silicate and Hypochlorite. Jour. AWWA, 79:7:116. 92 NOVEMBER 2 JOURNAL AWWA 97:11 PEER-REVIEWED SCHOCK ET AL

sharing the results. Some of the data interpretation cited in this article resulted from work by other members of the report team, including Richard Protasowicki and Bruce Pohlman from Black & Veatch and Michelle Frey of PureSense Inc. (formerly of Black & Veatch Corp.). Karen Eager, formerly of NEWWA, uncovered some additional water quality data and hydrologic information. Any opinions expressed in this paper are those of the author(s) and do not necessarily reflect the official positions and policies of the USEPA, NEWWA, HDR/EES, or Black & Veatch Corp. Any mention of products or trade names does not constitute recommendation for use by these groups. ABOUT THE AUTHORS Michael R. Schock (to whom correspondence should be addressed) is a research chemist with the USEPA Water Supply and Water Resources Division (WSWRD), National Risk Management Research Laboratory (NRMRL), Office of Research and Development, 26 Martin Luther King Dr., Cincinnati, OH 4268; e-mail schock.michael@epa.gov. He has 27 years of experience in drinking water and groundwater chemistry research, specializing in corrosion control and distribution system inorganic chemistry. An AWWA member, he also belongs to the International Water Association and the American Chemical Society. Schock has a BS degree in geology from Wright State University in Dayton, Ohio, and an MS degree in geology from Michigan State University in East Lansing, Mich. He publishes in numerous professional journals and has been awarded USEPA Gold and Bronze medals. Darren Lytle is an enviornmental engineer and Steve Harmon is a quality assurance manager with the Treatment Technology Evaluation Branch in the WSWRD. Anne M. Sandvig is a senior engineer with HDR/EES in Custer, S.D., and Jonathan Clement is a project manager with Black & Veatch Corp. in Amsterdam, the Netherlands. FOOTNOTES 1 SHAN-NO-CORR, Shannon Chemical Co., Malvern, Pa. 2 BW-, PQ Corp., Valley Forge, Pa. If you have a comment about this article, please contact us at journal@awwa.org. Schock, M.R.; Wagner, I.; & Oliphant, R., 1996 (2nd ed.). The Corrosion and Solubility of Lead in Drinking Water. Internal Corrosion of Water Distribution Systems. AwwaRF and DVGW Forschungsstelle-TZW, Denver. Schock, M.R.; Lytle, D.A.; & Clement, J.A., 199a. Effect of ph, DIC, Orthophosphate, and Sulfate on Drinking Water Cuprosolvency. EPA/6/R-9/8, USEPA Office of Research and Development, Cincinnati, Ohio. http://www.epa.gov/ord/webpubs/effect. Schock, M.R.; Lytle, D.A.; & Clement, J.A., 199b. Effects of ph, Carbonate, Orthophosphate and Redox Potential on Cuprosolvency. Proc. NACE Corrosion/9, Orlando, Fla. NACE Intl., Houston. Schock, M.R., 1989. Understanding Corrosion Control Strategies for Lead. Jour. AWWA, 81:7:88. Schock, M.R. & Wagner, I., 198. The Corrosion and Solubility of Lead in Drinking Water. Internal Corrosion of Water Distribution Systems. AwwaRF and DVGW Forschungsstelle-TZW, Denver. Schwertmann, U. & Thalmann, H., 1976. The Influence of [Fe(II)], [Si], and ph on the Formation of Lepidocrocite and Ferrihydrite During the Oxidation of FeCl 2 Solutions. Clay Minerals, 11:189. Sigg, L. & Stumm, W., 198. The Interaction of Anions and Weak Acids With the Hydrous Goethite Surface. Colloid Surface, 2:11. Standard Methods for the Examination of Water and Wastewater, 1992 (18th ed.). APHA, AWWA, and WEF, Washington. Stericker, W., 194. Protection of Small Water Systems From Corrosion. Industrial Engrg. Chem., 37:716. Stumm, W. & Morgan, J.J., 1981. Aquatic Chemistry. Wiley-Interscience, New York. USEPA (US Environmental Protection Agency), 1994. Drinking Water Regulations; Maximum Contaminant Level Goals and National Primary Drinking Water Regulations for Lead and Copper. 4 CFR Parts 141 and 142, 9:12:3386. USEPA, 1992. Drinking Water Regulations: Maximum Contaminant Level Goals and National Primary Drinking Water Regulations for Lead and Copper. 4 CFR Parts 141 and 142, 7:12:2878. USEPA, 1991a. Maximum Contaminant Level Goals and National Primary Drinking Water Regulations for Lead and Copper. 4 CFR Parts 141 and 142, 6:11:2646. USEPA, 1991b. Drinking Water Regulations; Maximum Contaminant Level Goals and National Primary Drinking Water Regulations for Lead and Copper. 4 CFR Parts 141 and 142, 6:13:32112. USEPA, 199. Method 2.7: Determination of Metals and Trace Elements in Water and Wastes by ICP Atomic Emission Spectroscopy, Revision 3.2. EPA/6/R-94/111. Environmental Monitoring Systems Laboratory, Office of Research and Development, Cincinnati, Ohio. USEPA, 1983. Methods of Chemical Analyses of Water and Wastes. EPA-6/4-79-2. Environmental Monitoring Systems Laboratory, Office of Research and Development, Cincinnati, Ohio. Zhang, P.; Ryan, J.A.; & Bryndzia, L.T., 1997. Pyromorphite Formation From Goethite Adsorbed Lead. Envir. Sci. & Technol., 31:2673. SCHOCK ET AL PEER-REVIEWED 97:11 JOURNAL AWWA NOVEMBER 2 93