Method to Rapidly Characterize Reduced Lead Corrosion in Phosphate Inhibited Drinking Water

Transcription

1 Method to Rapidly Characterize Reduced Lead Corrosion in Phosphate Inhibited Drinking Water L.K. Nalley,* V.N. Rafla,* R.J. Santucci,, * and J.R. Scully* The 21 Flint water crisis incited discussion on the legacy practice of utilizing lead in drinking water delivery pipe, and the critical need for sustained phosphate treatment in order to establish and maintain a protective lead film. However, there is a need to develop a method to quickly assess lead corrosion susceptibility and its mitigation by the addition of phosphate in drinking water. A potentiostatic testing procedure has been utilized to characterize the residual corrosion rate of lead in the presence of phosphate inhibitor more rapidly than existing long-term water sampling practices. This method involves electrochemical testing of freshly polished lead samples in representative potable waters simulating drinking water at the open-circuit potential, directly followed by a 6-day anodic potentiostatic hold in quiescent drinking water. The hold potential was selected based on a relatively common galvanic couple potential observed between lead and stainless steel and copper. This simulates common metal pairings in water delivery pipe and therefore represents a high anodic potential that is a near-worst-case scenario of corrosion in these systems. Results demonstrated that this is an effective method for determining the inhibitor-film-influenced anodic behavior of lead and provides a more rapid assessment of whether phosphate mitigates corrosion than long-term water sampling for dissolved lead. The ability to compare data from these tests with and without additions of phosphate enables rapid, and possibly timely, assessment of the effect of phosphate treatment on propensity for lead corrosion in stagnant drinking water. It was found that phosphate did inhibit the corrosion of lead, but not at short times and not linearly with increasing concentration of phosphate but monotonically. This particular set of experiments has a direct application in situations such as encountered in the Flint, Michigan water crisis and when considering new water sources in situations where information on drinking water corrosiveness and inhibitor efficiency is limited. This technique serves as a foundation for further modification as it can be coupled with Pb(II) water sampling and sample interface analysis to assess the complete fate of Pb(II) cations whether incorporated into scales or present in drinking water. KEY WORDS: corrosion, drinking water, lead, phosphate inhibitor, water quality, water treatment INTRODUCTION Lead is a known toxin that can enter the bloodstream through drinking water, and consequently can have serious health effects. Most importantly, consuming lead at an early age can have lasting impacts on a child s neural and cognitive development. 1- To prevent this exposure within the United States, the Environmental Protection Agency () has established a maximum contaminant level (MCL) of 15 parts per billion (ppb) of lead, 5-6 although no concentration of lead is safe for consumption, and all water utilities should strive for zero exposure. To achieve this, public utilities must balance competing interests in preventing public health risks while minimizing the cost of lead pipe replacement and/or the water treatment processes. 7-8 Due to the legacy use of lead in drinking water delivery pipe,9-11 in some countries such as the United States, the addition of phosphate inhibitor is an important aspect of water treatment in order to limit lead release. Phosphate has a scaling effect on lead and forms a stable film on its surface that prevents the leaching of lead into the water supply. 5-8,12 Moreover, the concentration of aqueous Pb(II) in equilibrium with lead phosphate is very low, as lead phosphate is very stable (Equation [1]). 13 In 21, the city of Flint, Michigan failed to add adequate phosphate inhibitor, citing economic struggles, 1 which raises critical questions concerning doseresponse, persistency, and the effects of many factors such as phosphate levels necessary to form and maintain a protective film on lead pipe. In public water supplies, phosphate is generally added as orthophosphate, whereas polyphosphates can be used but are deemed less effective in mitigation. 3Pb 2þ þ 2PO 3 =Pb 3ðPO Þ 2 logðkþ = 56. (1) Research on lead in drinking water has been continuously performed since 1991, 1- often reinforcing the use of phosphate inhibitor as a treatment method. This includes mineral studies on the precipitation of lead oxide in the presence of chlorine, 15 long duration pipe rig tests, 8,16-2 and on-site water sampling at utilities around the United States. 21 Each of these tests require months or years of testing, and acquisition and sampling of actual point-of-source water utility samples. Moreover, while Pb(II) concentrations may be assessed by direct water sampling, Submitted for publication: August 3, 218. Revised and accepted: November 8, 218. Preprint available online: November 8, 218, Corresponding author. rjsdv@virginia.edu. * Center for Electrochemical Science and Engineering, Department of Materials Science and Engineering, School of Engineering and Applied Science, University of Virginia, Charlottesville, VA 229. CORROSIONJOURNAL.ORG ISSN (print), X (online) 219 NACE International. Reproduction or redistribution of this article in any form is prohibited without express permission from the publisher. FEBRUARY 219 Vol. 75 Issue 2 17

2 information on interfacial reaction rates that release Pb(II) is lacking. Separation of variables and clear indications of controlling factors might be compromised due to the difficulty of maintaining constant conditions. In fact, in long-term flow exposure testing of lead coupled to copper in the presence of free chlorine disinfectant, the electrode potential of lead steadily increased to more noble potentials such that it surpassed the corrosion potential of copper and even approached that of silver. 22 In the same study, there was also a time-dependent response of lead-copper galvanic corrosion to phosphate, as phosphate increased the corrosion rate of lead for short times, consistent with other observations that the addition of phosphate can make galvanic corrosion of lead worse, even though it lowers the general corrosion of lead The utilization of electrochemical methods for direct assessment of scaling could shed insight on some of the questions above by enabling a more rapid determination of the efficacy of lead corrosion inhibitors using simulated drinking water solutions. If successful, this would permit a faster analysis of corrosion mitigation approaches, water corrosiveness, inhibitor persistency, time required for dose-desired response, and other operational factors under relatively fixed conditions. Electrochemical impedance spectroscopy (EIS) has been utilized to assess drinking water infrastructure, 2 but this has not been utilized to assess lead pipe corrosion in drinking waters. An anodic potentiostatic method was successfully deployed previously to assess the effect of chloride, sulfide, and alkalinity on the corrosion response of lead. 25 However, the method has not been applied to evaluate inhibitors. A method for determining the corrosion behavior of lead in to various phosphate inhibitor doses by implementing anodic potentiostatic testing is demonstrated here. The method might be particularly effective when combined with water sampling. EXPERIMENTAL PROCEDURES 2.1 Materials: Metallic Lead Coupons, Electrochemical Cell Design, and Solution Chemistries 99.99% purity lead samples were tested in a standard three-electrode electrochemical cell with a Pt-mesh counter electrode, a saturated calomel electrode (SCE) reference with a Luggin probe tip, and lead working electrode samples. Prior to testing, lead coupons were polished to an 8 grit finish with deionized water. The exposure area was.785 cm 2, the electrolyte volume was 3 ml, and the distance between the working and reference electrodes was approximately 2 mm. Two simulated drinking water chemistries, representative of two different drinking water sources, with varying concentrations of phosphate (Table 1) were utilized as the testing electrolytes in Trade name. (1) Estimation of a galvanic couple potential of V SCE was based on potentiodynamic testing of stainless steel and copper electrodes over a range of electrode sizes. The coupling potential of the smallest electrodes used for lead, Type 316 stainless steel, and copper (91, 8, and 25 cm 2 ) was chosen for the potentiostatic hold. These samples provided relief from ohmic resistance effects on the E-log(i) behavior due to a decreasing ohmic resistance contribution with decreasing electrode size. Experimental E-log(i) data were normalized for all tested electrode sizes to maintain a 1:1 area ratio within the galvanic pairs. It is recognized that anodecathode area ratio, distance between anodes and cathodes, flow rate, water chemistry, and the galvanic couple current and potential distribution all affect the potential experienced by lead. However, ohmic potential effects only decrease the anodic galvanic couple potential experienced by lead. Therefore, the approach taken here established an upper bound on galvanic couple potentials. order to isolate the effect of phosphate inhibitor. The primary test solution simulated water from the Detroit water utilities before the deleterious switch in water supply to the Flint River. 1,26 This synthetic drinking water solution was designed to have a low chloride-to-sulfate mass ratio of. in order to lower the corrosion rate. The ph of the simulated drinking water was adjusted to 7. with acetic acid or sodium hydroxide additions prior to each trial, and the conductivity of the solution was 1.85 ms (as measured by an Accumet conductivity meter, model AB3 ). The concentrations of phosphate, added as Na 2 HPO, were,.1, 1, 2, 5, 8, 1, and mg PO 3 =L (Table 1). Another simulated drinking water was created based on the supply in Washington, D.C., sourced from the Potomac River, 27 with phosphate concentrations of mg/l and 1 mg/l. This secondary solution was included to demonstrate the utility of the potentiostatic hold test to be capable of assessing inhibitor benefits across various simulated drinking water chemistries. All solutions were tested in open-atmosphere conditions with an assumed CO 2 partial pressure of atm. 2.2 Anodic Potentiostatic Hold Technique to Assess Electrolyte Corrosiveness as a Function of Inhibitor Concentration The open-circuit potential (OCP) of the lead sample in each electrolyte was measured for 3 min prior to testing. This was immediately followed by a 6 d (1 h) anodic potentiostatic hold at V SCE to quantitatively assess the relative corrosion susceptibility of lead in the selected waters with and without phosphate inhibitor. This potential was chosen based on the estimated galvanic couple potentials (1) of lead coupled with copper and stainless steel in these solutions, mimicking possible worst-case galvanic pairs that might be found in application 16-17,28 (Figure 1). Three replicates of each phosphate concentration were performed. RESULTS 3.1 Assessment of Open-Circuit Potential, Anodic Current, and Anodic Charge The OCP, current density, and accumulated charge density over 1 h in inhibitor-free drinking water and with the addition of 1 mg/l of phosphate is demonstrated in Figure 2. In Figure 2(a), the OCP of the 1 mg/l was initially much higher than the mg/l, but after 18 min (1,1 s), these values approach each other. At this point, the OCP of the solution containing no phosphate continued to increase, whereas the OCP of the phosphate solution decreased and leveled off. Figure 2(b) reports the current density measured over the course of the 6-d potentiostatic test. A similar pattern appears in the phosphate-containing solution, which initially has a higher current density over the first 15, s. At this point the current density increased more gradually compared to the mg/l case. Comparatively, the current density of the mg/l solution increased either linearly or accelerated with time. Figure 2(c) depicts the same long-term potentiostatic data but instead illustrates the results as a function of the total measured charge density. Figure 2(d) illustrates the results of the same potentiostatic tests performed in an alternate simulated drinking water solution at mg/l and 1 mg/l of phosphate. These results were reproducible between the three trials performed at each condition. The ability to distinguish between phosphate-treated waters and phosphate-free solutions is evident. 18 FEBRUARY 219 Vol. 75 Issue 2 CORROSIONJOURNAL.ORG

3 Table 1. Summary of the Electrolyte Compositions Used for Each Trial as Simulated River Waters Simulated Detroit-Sourced Drinking Water Chemistry, CSMR =., ph = 7. Simulated Potomac River-Sourced Drinking Water Chemistry, CSMR =.8, ph = 7.9 Chemical Species (A) Equivalent PO 3 concentrations of,.1, 1, 2, 5, 8, 1, and mg/l. Concentration (mm) NaCl 1.1 MgCl Na 2 SO 2.86 NaHCO 3.79 Na 2 HPO, 1, 11, 21, 53, 8,.15, 1 (A) CH 3 COOH 5 CaCl 2 2H 2 O 58 MgSO 2 NaHCO 3 9 Na 2 HPO,.15 Potential, E (V SCE ) Lead Stainless steel Copper Current Density, i (A /cm 2 ) FIGURE 1. Anodic potentiodynamic polarization behavior of lead and cathodic potentiodynamic polarization behavior of copper and stainless steel in simulated Detroit-sourced phosphate-free solution to simulate a 1:1 galvanic couple surface area ratio with limited ohmic resistance between anode and cathode. The galvanic couple potential thus determined was used in subsequent potentiostatic polarization tests. Experimental E-log(i) data were normalized to simulate the 1:1 surface area ratio, accounting for varying surface areas of the lead, Type 316 stainless steel, and copper electrodes (91, 8, and 25 cm 2 ). Figure 3 demonstrates the variation in total accumulated charge density over time with the addition of varying phosphate levels. The solution with mg/l of phosphate demonstrated the highest charge density, achieving a total charge density of 1.2 C/cm 2 over the course of the experiment. In comparison, the 1 mg/l and mg/l samples demonstrated the lowest and nearly the same accumulated anodic charge density, at a value of 7 C/cm 2. If all of this charge produced only aqueous lead species into the closed volume 3 ml flat cell, then the resulting concentration of lead would be 3, ppb in the absence of PO 3 and 76 ppb in the cases of 1 mg/l and mg/l [PO 3 ]. Figure 3 also reports the calculated soluble Pb(II) concentration in the case where 1% and 1% of the oxidized lead assessed potentiostatically is dissolved into the solution, assuming the remaining 9% and 99% contribute to the formation of Pb 3 (PO ) 2 which remains on the surface. DISCUSSION.1 Role of Phosphate Inhibitor in Mitigating Oxidative Discharge of Lead Each potentiostatic test provided key information about the corrosion and passivation rates of lead. The ability to ascertain the efficacy of the phosphate treatment was examined in two different simulated drinking water solutions based on water supplies in both Flint, Michigan and Washington, D.C. Each plot shown in Figure 2 demonstrates that the accumulated amount of anodic dissolution and anodic corrosion rates are lower in the presence of phosphate (as verified from anodic current/charge analysis). This is evident due to both the lower accumulated total charge density (Figure 3) and lower current density (Figure 2[b]) over the duration of 6 d. In no case did the net anodic polarization evolve into net cathodic polarization of lead such as would be possible if the OCP of lead with a phosphate scale rose above V SCE. Figures 2(a) and (b) confirm this point, with respect to the OCP and the consistently anodic current measured during the test. In future work, it would be informative to elucidate the state and amount of Pb(II) released as aqueous species into solution vs. insoluble lead corrosion products in the scale to determine the percentage of charge associated with each. As a first approximation, a comparison is included here where the soluble lead released into solution is calculated from the anodic charge measured during the potentiostatic hold experiment. The amount of charge which produced aqueous Pb(II) ions was assessed at 1%, 1%, and 1% as soluble lead released. In the theoretical case of 1% soluble lead release (Figure 3), the concentration of phosphate inhibitor is the crucial difference between passing and failing the minimum safety threshold of ingestible lead. It is noted that within the first 2 h, lead in the phosphatecontaining solution initially corrodes at a higher rate than the phosphate-free solution, and this is likely attributed to the trade-off between the higher electrolyte conductivity of the phosphate-treated solution which might accelerate corrosion and lead phosphate film formation, growth, and coverage which requires time for the formation of a protective scale that CORROSIONJOURNAL.ORG FEBRUARY 219 Vol. 75 Issue 2 19

4 (a) OCP (V SCE ) (c) mg/l PO 3 1 mg/l PO 3 5 1, 1,5 mg/l PO 3 1 mg/l PO 3 (b) Current Density,i (A /cm 2 ) (d) mg/l PO 3 1 mg/l PO mg/l PO 3 1 mg/l PO Detroit-Sourced Solution Potomac River-Sourced Solution FIGURE 2. Comparison of (a) OCP (V SCE ), (b) anodic current density, and total charge density collected from 6-d potentiostatic holds comparing mg/l and 1 mg/l concentrations of phosphate in simulated waters: (c) Detroit-sourced and (d) Potomac River-sourced solutions mg/l 1 mg/l mg/l 5, 1, 15,.1 mg/l 2, 5, 8 mg/l 1, mg/l Increasing [PO 3 ] mg/l 1% Released 1 mg/l 1% Released 3, 2, 1, 1% Released FIGURE 3. The effect of phosphate concentration in simulated drinking water solution on the total anodic charge density associated with anodic dissolution of Pb as a function of time. Included is an insert of the total charge density evolution at short times. The surface to volume ratio was.785 cm 2 :3 cm 3 = cm 1. An additional scale is included which conveys the concentration of lead in the total stagnant test solution during the test if the anodic charge density measured released aqueous lead species at a 1%, 1%, and 1% production. 1 Equivalent Lead Concentration (ppb) 15 FEBRUARY 219 Vol. 75 Issue 2 CORROSIONJOURNAL.ORG

5 reduces corrosion rate. Despite the phosphate-free solution s initially lower charge density, the anodic charge density is eventually suppressed in the phosphate-containing solution in less than 2 h. From a public health perspective, it is significant to note that it takes a relatively long period of time for phosphate inhibitor to become effective; undoubtedly this traces to scale formation to achieve some critical condition in terms of coverage and degree of protection. Meanwhile, even inhibited waters might not be safe to drink at least initially. After this suppression, the accumulated anodic charge due to lead oxidation in the phosphate-free solution increases at a faster rate, which is observed at several different phosphate concentrations in Figure 3. As shown, it was observed that corrosion mitigation increases monotonically with concentration, not linearly. Specifically, phosphate did decrease corrosion susceptibility between mg/l and 1 mg/l concentrations. However, at higher phosphate levels, various phosphate concentrations resulted in the same mitigation effect, as demonstrated by the 1 mg/l and mg/l trials (Figure 3). This may suggest that at these concentrations, a greater amount of phosphate does not translate into a better scale nor improve corrosion mitigation, suggesting that there is an upper limit to the concentration of phosphate that needs to be added to drinking water. Moreover, there seems to be an incubation period for effective inhibition and reduction in residual corrosioninduced release rates based on some critical properties of the scale layer, which should be determined in future work. CONCLUSIONS AND FUTURE PROSPECTS Anodic potentiostatic testing is demonstrated to be an excellent, rapid, and reproducible method for assessing the corrosion behavior of lead in the presence of varying amounts of phosphate. By sequencing an OCP test prior to the potentiostatic hold, key information related to OCP behavior, current density, and total charge density can be collected over a series of days. Additionally, comparing the results of several trials with varying concentrations of phosphate enables a direct assessment of its scaling benefit toward reduced lead oxidation in a dose-response investigation. The duration of these tests proved long enough to determine information regarding the long-term passivating effect of a phosphate scale on lead corrosion. A key finding was that phosphate-based passivation does not occur immediately, but eventually reduces lead corrosion. Furthermore, by investigating the effect of phosphate over a range of concentrations, it was evident that phosphate s benefit does not increase without bounds with respect to concentration, making an economic case for determining the lowest necessary amount of phosphate that should be used given the time and flow rate. This method of experimentation must be explored for other experimental conditions (electrolyte composition, ph effect, water hardness, temperature effects, CO 2 effects, etc.). Further experimentation should be performed to determine the partitioning of lead between the ionic species, lead oxide, and lead phosphate in solution or on the metal surface. Further analysis should be conducted regarding the nature of these lead scales formed during Pb(II) release, particularly assessing the formation of phosphate, carbonate, and chloride compounds. Coupled with a specific ion electrode and simultaneous scale characterization, the complete fate of Pb(II) can be determined in real time. Additionally, investigations determining the relative roles of flux and volume flow in a fixed volume would provide a better understanding of the dissolution of lead into drinking water sources. ACKNOWLEDGMENTS The coauthors of this report donated their time to this research voluntarily. The authors are grateful for the partial financial support and contribution of the National Science Foundation under NSF DMR # for material supplies. References 1. M. Edwards, Environ. Sci. Technol. 8, 1 (21): p R.L. Canfield, C.R.J. Henderson, D.A. Cory-Slechta, C. Cox, T.A. Jusko, B.P. Lanphear, New England J. Med. 38, 16 (23): p W. Troesken, J. Hum. Resour. 3, 3 (28): p R. Rabin, Am. J. Public Health 98, 9 (28): p B.P. Boffardi, A.M. Sherbondy, Corrosion 7, 12 (1991): p U.S. Code of Federal Regulations (CFR) Title, Protection of Environment, Part 11, Subpart I, National Primary Drinking Water Regulations, Control of Lead and Copper (Washington, D.C.: Environmental Protection Agency, 1991), p M. Edwards, S. McNeill Laurie, J. Am. Water Works Assoc. 9, 1 (22): p L.S. McNeill, M. Edwards, J. Am. Water Works Assoc. 9, 7 (22): p J. Hu, Y. Ma, L. Zhang, F. Gan, Y.-S. Ho, Sci. Total Environ. 8, 7 (21): p R. Renner, Environ. Sci. Technol. 3, 6 (29): p W. Troesken, The Great Lead Water Pipe Disaster (Cambridge, MA: MIT Press, 26). 12. L.W. Wasserstrom, S.A. Miller, S. Triantafyllidou, M.K. Desantis, M.R. Schock, J. Am. Water Works Assoc. 19, 11 (217): p. E6-E J.G. Speight, Thermodynamic Functions (Change Of State), in Lange s Handbook of Chemistry, 17th ed. (New York, NY: McGraw Hill Professional, 217). 1. J.R. Scully, The Bridge 6, 2 (216): p D.A. Lytle, M.R. Schock, J. Am. Water Works Assoc. 97, 11 (25): p J.S. Clair, C. Cartier, S. Triantafyllidou, B. Clark, M. Edwards, Environ. Eng. Sci. 33, 1 (216): p J.S. Clair, C. Stamopoulos, M. Edwards, Corrosion 68, 9 (212): p C. Cartier, E. Doré, L. Laroche, S. Nour, M. Edwards, M. Prévost, Water Res. 7, 2 (213): p J. Hu, F. Gan, S. Triantafyllidou, C.K. Nguyen, M.A. Edwards, Corrosion 68, 11 (212): p Y. Xie, D.E. Giammar, Water Res. 5, 19 (211): p P.T. Cardew, J. Water Health 7, 1 (29): p R.B. Arnold, M. Edwards, Environ. Sci. Technol. 6, 2 (212): p C.K. Nguyen, B.N. Clark, K.R. Stone, M.A. Edwards, Corros. Sci. 53, (211): p I. Frateur, C. Deslouis, M.E. Orazem, B. Tribollet, Electrochim. Acta, 2 (1999): p C.K. Nguyen, B.N. Clark, K.R. Stone, M.A. Edwards, Corrosion 67, 6 (211): p to M. Edwards, Synergistic Impacts of Corrosive Water and Interrupted Corrosion Control on Chemical/Microbiological Water Quality: The Flint, MI Water Crisis, Flint Water Study, October 215, Corrosion-Presentation-final.pdf (December 1, 218). 27. T. Simoni, J. Parks, M. Edwards, J. Am. Water Works Assoc. 99, 6 (27): p , R. Gregory, Water Environ. J., 2 (199): p CORROSIONJOURNAL.ORG FEBRUARY 219 Vol. 75 Issue 2 151